• . 
Table of co&teate 


.• , 


subject 


Chemistry 131 
General Policies and Procedures 
Collaboration Policy 
Syllabus 
Calender A .y 
laboratories and Handouts 


Inorganic nomenclature 
Scientific Method Laboratory (75 pts) 
Data Analysis Laboratory {75 pts) 
Emission Spectroscopy Laboratory (loo pts) 
Calorimotry Laboratory (75 pts) 
Energy simulator Handout 
Molacular Geometry and Electronic 


structure Laboratory (75 pts) 
solutions Laboratory {75 pts) 
Acid Rain Laboratory (75 pt*) 
Spectrovétrie Determination 


of Kc Laboratory {100 pts) 
Acid Base Titration Laboratory (75 pts) 
Photography Laboratory (75 pts) 
Qualitative Analysis Laboratory (100 pts) 
chemical Kinetics Laboratory (75 pts) 
Electrochemistry Laboratory (7 5 pts} 


Organic Laboratory (75 pts) 


... 


• y 
• 7 '. 
•• . . . 


£A3fi 


' . . ' • . ' • • 

. 1 ; • 


-.'5 


'y

3 


io 


:y "y

14 


15 


22 

30 y 

39 ._ 


51 
51 
• • .•••:'• . . ' ' . ' • • ' 


•; -64A 
75 • -•' 7 '• yy. 

86 
97 
111 


123 


132 


145 


154 


• '• • . -•••:: •-' 
167 


--. :"' : ••'•• ' • ' : •• 

.-"•"•"•• •'.. ••":-.

191 



Chemistry 131 includes both lecture arid laboratory work. To better 
understand chemistry, we believe you should become acquainted with 
fundamental principies and concepts and be able to apply them in the 
laboratory, our goals in this course are to: 


1) train you to think scientifically (develop curiosity), 


2) train you to think analytically (solve problems), 


3) provide you with scientific information of professional and 
social valuet 

4} prepare you for future science and engineering courses, 


5) teach you basic chemical principles, 


6) develop your ability to communicate scientific information, 


7) provide you with laboratory expérience* 


As in most science or engineering courses* General Chemistry is 
difficult* To be successful* ycu must commit yourselves to academic 
excellence* You must diligently prepare for each class by reacting the 
assigned portions of the text and working as isany of the homework 
problems as. you can, be attentive in class, and be willing to ask 
questions and seek extra instruction. You were selected for this 
course on your past academic experience and performance on the 
chemistry placement exam - you have the ability to do well in this 
course* The Chemistry faculty is here to assist you in doing the best 
you can. 


Keep in mind that Chemistry is a 2 course unit, 6 semester hour course. 
You will receive twice the credit for this course as you would for the 
tsott of your Other courses* You should therefore expect to spend twice 
a* «much time in preparation for Chemistry 131, 

We hope you find this course interesting and informative; we know you 
will find it challenging. 


: • ". 
Í • ... 


> _ • J 



. . -• • .... . ..... . .,J 

• •••• ... • ... 


* • *..*•'. • • * * *. 


".**"* '*••. ' ." .*' * 


GENERAL POLICIB8 AND PROCEDURES 


1. ..The. schedule for Chemistry. 131 . is significantly different than 
••• other courses at thé Academy. Check the syllabus daily. In gênerai. 
you Will have lectures.on Monday, Wednesdayf and Friday, one 2 hour lab 
on Tuesday and a 2 hour recitation on Thursday* If you have a 


.' question, be sure to consult your syllabus or ask your instructor. If 
in doubt, go to:your classroom during the assigned period and check» 
...Your, instructor will clarify the schedule during the first lesson. 


.2*. Do not schedule any appointments (dental, medical, EI, tailor shop,, 

etc,) during any class or lab period. 


3i: If you know you will miss a class, inform your instructor at least 


one lesson In advance and submit a Form 76 before your absence if 


possible. You ere responsible for all work assigned during your 


absence. If you miss a GR, you must take the make-up exam. If you 


will miss a GR due to a trip, inform your instructor at least a week in 


advance so he can coordinate with the OIC or OR for a make-up exam to 


be administered on the trip. 


..*•**** *. * . ". 


. 4. You must keep up with the course work even if you are hospitalized 


or on bedrest* If you need EI, notify the hospital liason officer to 
. arrange EI with your instructor, or just call the Chemisty Department 
(X2960). Your instructor will give you El if you are hospitalized for 


•A an extended period. 
• .-. ... •.-.-. -..-••.• . ... ... •. .• . .
.. . . 
5. Do not leave the classroom or lab until you are dismissed by your 
; instructor. 
y 7"-6. Please keep the classroom and coat closet clean and neat. Keep» the . 


doors to the coat closet closed when not in use -if you are the last 

to usé it, close them. All outer garments (jackets, parkas, overcoats, 
.overshoes, caps, etc.) must be neatly hung or stored in the coat closet . 


.••.-.. -they are not permitted in the classroom. 


•".7;. Sectior.Marcher Responsibilities} 
a. Before class. Maintain order in the classroom. Ensure the' 
classroom and closet are neat arid orderly» Collect any assignments due 


before the start of class. 


b. Start of Class» Call the class to attention, render salute 
and report the class ready for instruction, and give your instructor 


all homework due* If your instructor does not arrive within 5 minutes 


after the start of class, report to the Chemistry Department for 


further instructions. 


. ç.,. During Exams. If the instructor is absent during an exam,• 
maintain order and give the command, "cease work, " when the time has 
expired for the exam. The time should be written on the blackboard. 
Any cadet may give "Cease Work" when the time expires. 


/ y : '..'• • • • ' • ' • • • • • "••-. • '. • ' . ' . • • • • ' . • • • " •' . " . , " • • : :• 

• 


'• . • . • •• • ... • • • • •. . . '. « . • ... . • • . ' '. . y • 

.'•' *• ' t "' *-.•. . '" .• * \" % ' •' ' '. ' . • . ' ' ' ' ' * . ' * • " • . .' "'.''• . .'"'. "..' 


ñ/ classi-C'Oîft Materials. To each lecture class, you ¡mist bring youxtextbook, 
3 ring binderf paper for taking notes, a pencil or cen, and a 
calculator fyou will not be permitted to share calculators during 
quizzes or GRs). To. each laboratory period, you rcust bring JÏ pen (all 
experimental data tnust be recorded in ink) and your lab manual or 
handout for the lab. 

9* Homework. In general, homevork will not he collected or graded-
Your instructor, however, has the option of collecting and grading your 
homework. The problem is assigned indicate the knowledge levai and 
proficiency expected of you. At least 60% of all questions <sn the Ofcs 
and quiises «ill be homework problems* 


:10. Extra Instruction* understanding of the material assigned is of 
the utmost importance. If you have difficulty, you s&tould ss&fc help 
immediately from the El room CFoom 2C56) or your instructor- The El 
room is opsn from 0800 to 1600 every day unless otherwise noted. It 
will be closed during period H-S and lunch. 

13* Laboratory Experiments, Lab experiments supplement and reinforce 
the concepts covered in the classroom. They givâ you the opportunity 
to see chemistry in action. Here at the Academy, we do eKpcrircents in 
ways unfamiliar to youu our experiments are performed at the 
microábale. With these experiments, we minlie ize student exposure to 
hazardous and toxic materials and reduce chemical waste> We are also 
able to perform more experimental work. While performing ex#erlatents, 
give SAFETY primary consideration. TK1NK BEFORE YOU ACT, ASK IF IN 
DOUBT* 


A. All labs will include the following on the first page* 
_ -j . *_ * 
.. y . '• •' • ' ' •' • • 


Collaboration at sternest; state who you collaborated with on the lab. 
If no collaboration, «tare NO COLLABORATION. rftlLUKZ T O HAV E THI S 
32ÏTIOM WIIií, RESULT IV A 3 5% GÜT 0» TES LAB. 

B, For the Ernies!on Lab, Enthalpy Lab, and the Spectrometric 
Determination of an Equilibrium constant Lato the following fooiat will 
be used: 


CollaboraticJ* Statements State who you collaborated with on the 
late» If no collaboration, e**te NO COLLABORATION» 


Objectives state the purpose or cbjactive of the experiment in 
one sentence. The objective is cheflical in nature? the objective is 
not to learn how to use an instrument* 


7.7 
Theory* In at least one full page <the equivalent of 1 full, 
fcypad double spaced patje) discuss th* theory governing the results of 
tfce experiment* The section generally begins with a short introduction 
detailing why yeni are doing this experiment, followed fry a discussion 
of tha procedures u&eid to obtain the necessary data. To adequately 
write this sectiont y°11 should consult your handout and/or textbook. 
; • • . : 



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¿-o . K^.I -^ _ . •• •• •-. -j5--;-v j. _ ^ . . ,j -; •_ .• . ;. ..-._.. . .. .-•.'.•'-• *-" • " •"•"'"."'.,•'..' • •, '• : '• ; _. ;•• -,-. 

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• ••-•• .: • . jf


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Do ribt: «imply copy another source, Compose this theory inyourown 


yy words. 


J *• s l ,f'.. . '. -• 

pisou*ftion of Resultst Answer thequestions in thehandout. 


conolu*iont What can you conclude about theexperiment? Give 


:numerical resuits in theconclusion anderror analysis. A; 


.••\ 

C. Fortheremaining 75point latos, theprelab exercise will befilled7 
out BEFORE coming toclass, the labwill befill out during class and 
the conclusions: will answered. Unless directed otherwise byyour 
.-\y-:m • **;.

instructed alllab reports will bedue the Monday following the lab 
day; Only 4ofthe 10(75point labs) will begraded. However, all 
,.must becompleted and turned in. or youwill receive!an INCOMPLETE FOR 


THE .COURSE, yyyyyyyy yy.;yyy.~ yy y^yyyyyyyy yyy-yy. '-yy 

"yy -y-'*-'' >-'*' ---•:' •--' '•7 .r-;-•-:.,--•«-,'..'y "'A. y ?-'"'• „ /'•'' -"•" A:> 

12, Late Work* Allassigned work isdue atthe beginning of the 


period' Late penalties will beassessed according to theschedule 


. ..neJrOW* y.,-s v::-...-... „.,' .-.y y ¿yA.i ,y....y.y yy y y y. ; •-.••,-,-v •;:•': •;:• J.-•• '-y :':'•• 
77/77 , ;v-• • •<1lesson late y; y25% y y 
A^;77;í yy yy.y < 2Isasons late /y7S0%' .-y7; ,;7;. 

* „•'..-•:,.. .:.-7"'"A"'V;;"2 .lessone:late ^.777iil0'*;• -y• ,,,•.•,; 7y/;7 '"'(,"' 
I-, ••¡.'7 

Since Chemistry. 13Í does not follow thenormal M/T-Óay schedule, each 


class meeting (lecture orlab) is onelesson. 
..y¿ "'A., -y ""• 7.7 :.' y.yy. :"•; .:\y-"•"• .•^"*7 ,•; • -7. •' ••" "-.--1' -. ", 

W-l" 

13. computer Aided Instruction* ïou should have been issued1 three 
Computer di^Kt with chemietry computer aided instruction software. The 

instruction program:descriptions arelisted below: 7-7.77 


yfr^y 

\-y^l • • v . -f-• r

. "Ï

QlMK EZS-3XJ&Q& 

• •=-. '...•• 3 ' 


1 ThA-J^ejaantfti •Periodic Table, Nomenclature of the 
Elements, Isotopes/Atomic Weights, Properties of Some 
77;: ..y..'-.. Elemente, Mystery Element Quiz '," \y'"> 

,


> • -.-.*yy\ y''-y.-.y,v ,-y.-y,yy-' >'¡7.,A7;;Í:•'•' . yyy y it ; yy '••..•„••-... y-.y... ,?\ •y-.i, . i>; A-.y^ 

Inàl^i«aiiL_^ffiîSÊïiiiAtïiSfi! Binary salts, Variable oxidation 

S^s' ^ SíTjts, BiíMsry G»japoüi"iSíí& oí' Twc¡ isoîïïaetàis, Acids,

Sáses, Ternary Salts, Review Problems 


y:;.íí''•'J• 


Solutions; Solubility, Solubility Experiment, Freezing
Points, Molecular Weight Determination Experiment, Weight 


•i' ' J , -A.• y.> 
Percent», Molarity, Dilutions ' 7 yy.

yyyy?^'yy-y'yíi7.: ... ••'-,y.v A "A , -.:"'"7 7v '<" •••-7••"••'' 

2 Atomic g-o-rtaala ^rtd a^l«cula r Weights: Chemical Formulas, 
)yy y7 .
..;V: Atomic We i %**te, Moleculair Fonnulas/Weights, • Gram Molecular 

.y y. -..y > ":";" Weights, cram/tfole Problems ;'7:"7. * y'"y.'\.y \y 

•"•'.•:"-• A :. • ••••• ' . 7 •'• . • . A .: "•••• : y •' • ••;.' • ••..-.•.; -. . , .-; ;-• y y •/ ' ' y . • • . -: 
''. Percent Composition andEmpirical Formulasa Introduction, 

y^;i ^sr^T^^t Co^si^sitio^T E^^irical Formulas, Mg-HCl Experiment, •'•;' 
.AÍ;" .7 "A -Empirical Formula.Problems •
•• '' '
'' -7-y'y-'y :"-y


çjjej tüáses : .. Arcade style game testing your knowledge of 
Chemica
ChemicaChemical
ll reaction 
reactionreactions 
ss 


> ^ '-.^ •"•','• -.: 'i-' '-y 7, 7^" ... y^ < : ", H '-' •*-7 -1 7 7--y -• 7 7 .',-^y.. 7 7^ .A1 .yy y y'À; yy'.y

^•:'¿-v..-:.^i'!'!:'::!'i yyyyív y .;.,y:;, : y-•••••. y-*.r 

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•A A' ; y.,y ï '•vAA.'v'-. ' ' ..,•-:' ^.v--, •-•:; . v•
•• y. A * '=. ,-. .,:il ^ ...yf y •••• ;• \s -••• y-y 7 
*•}••••.. •* • ^••^••.••--•• • •••••? ^ i^r..>. . ^-.^-. • cr^ •• i '..-^ i •• ' •••• • > -, -^^-, • • >••• :•*••• • r ; 7 .-•'••• --i !•• •'.••-. •• ••••• •--:•...•,•••• . y r; ",.^i . .-o;^ -•:-•'•• • i.--1 ^ • • • 'H -. ...^^' ^ --r 

i^.^i >^; ; iî i K : Ai v-•J -Î ,i. • •. -i ;••••.•• Vi . : ••-•'; ••:••,:. :. v ~¡ ' -" -: -i.'-, • 'X-• -L. ! J -••'• •' :•-^ f •• •• Al ^-y ."••• •.-'• -'-!'•:.."•'-••: \.\ --. '•-M", '.• !,V -•/ • ' -"••'? '• -f •-'"••• ' '' n—~"V: • ' •• ; '" h • '=•; ' -y-; •,-••' 

:j::!

1. !>-"'• SÍ "-*i. ::ÍÍ?,VÍ¿':'C-f -r-':V^'--,••'", y--^ ^ --'i.'-"-'•.*.•'••• *f=f-ï' y L£¡y --¡'V-í-' 'f..--•:•••• ' ' ;> ••' * V •<;•.•. y'-Vr'-'r •: ;•!.>•••.'• .'.' *.IÏ -"-IvÎ! ¡'. ::y;--T:-'. -:-J-i • -t.>.i.^o.-cj,'i .^/v^-Ii 
V;-" . '*• : :^-^--,r-...'A-.<. <-,-, • -y-.->,-Cv t.rívíi ^ •»».•--.-V •: i.V-. .•:...•••••".".••. ~r •ï> ; r*, 'r-.'.^ l •.Í -:••• ••.• ••( -..-• >.•• f ,-... , V -..••*-.• .I.Ü. r.-.:i ••.,-• .' <, 
i '.-.ir-1 -.'i-.---.: .•• :.'/ " .••« f•,-, .• ..• -• .,1-' .-:.••••.• 


•A 3 Chemical Formulas and Equations: Chetnical Formulas, 
Chemical Equations, Writing Chemical Equations, Balancing 
yy77.7.yy/'chemical Equations, Review Problems: 7 


Acids and Bases In Water: pH, Water, Strong Acids and 
Bases, Measuring pH, Neutralization, Acid-Base Titrations, 
Titration Experiment, Strong vs Weak Acids 


Metric System: Prefixes. Temperature, volume, Weight, ' 
Concentration, Density, Special symbols and Functions 


14. Additional CAI is available on the VAX-A computer via Falconnet, 
The programs and descriptions are as follows: 
Program QCBgription 


Math Tutorial on math and scientific notation 
7'.Loge-'.y Tutorial and drill on logarithms 
Names Drill on chemical nomenclature 


.7 Namegame Arcade-*style nomenclature game (Bonus Pointe) 
Mole .:,'•'•.Mole concept and molecular weights 
Limit: Tutorial and drill on limiting reagent problems 


Thermo Thermodynamic concepts and problems 
Atom Tutorial and drill on electronic structure 
•y;-;:. .Force.'77. Tutorial and drill on intermolecular forces 7 
Gas Tutorial and drill on gas laws 
• '.•-'. pH '•;••• Tutorial and drill On pH problems 
Acid Tutorial and drill on strong and weak acids 

77 ËMF "'':".•' Tutorial and drill oh electrochemical concepts 


yQua1•'•• ' Dry run of quelitatíve analysla before the 1éb y 
yft&âox;;y'Tutorial and drill Oh redox reactions 

(jlíafli•
•••'. Tutoria1 on the periodic table, eleménts* etc. •.7 


Alchemy A chemical knowledge adventure game (Bonus Pointa) 


We will give you information detailing the steps necessary to access 
the CAI software on the VAX-A. 


15. Bonus Points from CAI. Aft: you may have noticed from above, bonus 
points are available when playing NAMEGAKE and ALCHEMY. Normally, Full 
Collaboration is permitted when using CAI* However, since points may 
be awarded for these two, INDIVIDUAL EFFORT is required. These games 
are available for use throughout the year, however, you may earn bonus 
points only during certain dates. we will give you additional 
information during the semester7 Ajj*£4 instructions. çarefully.1 .;When 
playing these games* 
. _• \ -. 

' • • 



NAMEGAME ig an arcade-style quiz where you give the correct 
name or chemical formula for elements and compounds. You are 
awarded 10 points for each correct answer and penalized 4 
points for each wrong answer. You can play the game as many 
times as you desire, only your top score counts. Bonus 
points will be awarded to the top 30 scorers. NAMEGAME will 
be stopped at the end of Lesson T-21. 


ALCHEMY is an arcade-style adventure game where you work your 
way through the land of Alchemy by overcoming adversaries 
with your knowledge of chemistry. ALCHEMY has ten levels you 
can progress through during the semester, earning up to 3 s 
points for each successful completion of a level. The number 
of points you earn depends on the number of questions missed. 
You lose 10 points for the first question missed ("fatal 
encounter1*) and 5 points for subsequent missed questions. 
After missing a question, you're removed from the game and 
must reenter. If you miss more than 5 questions but finally 
complete the level, you earn 5 bonus points. Levels will be 
open for bonus points as we enter the block of material that 


is "tested" in them. Alchemy will be stopped at the end of 
Lesson T-42.. The following is the schedule for the 10 levels 
of alchemy: 


Level Number Opened closed 


1 710 Aug 17 AtiS 

y'':7;2"- •"•.,. • .7 l o Aug '. 7,y ..24 ^ug 
..3: ; 24 Aug 7y-.'7ASept ' 
7* 7 ..' í..'Septy77y¡ 21 Sep t 


5 21 Sept 5 Oct 
è 3 Oct 19 Oct 
7 7' 19 OCt 2 Nov 
8 2 Nov 16 Nov 
9, 16 Nov 30 Nov 


:.- in' .'•:'••': ' 30 «OV A il Dec y y 


REMINDER: SINCE ALCHEMY AND NAMEGAME ARE PLAYED FOR POINTS, YOU MUST 
WORK ALONE* y yy 


,"•••• ' . -'y'." '-• •• •' 



a 


16. 
GRADING; 
7.7 
4 GRs (300 pts/each) 1200 
8 Course Quizzes (30 pts/each) 240 
2 Instructor Pop Quizzes (30 pts/each) 60 
Lab Reports: 4 G 100 pts/each 
400 


(Emission, Enthalpy, 
Determination of Kc, 
Qualitative Analysis) 


4 $ 75 pts/each 300 
(4 of the remaining 10 

labs will be graded 


randomly.) 


Final 
B00 


TOTAL 


AAA 


Points for Prog 


2 GRs 600 
2 100 pt labs 200 
2 75 pt labs ISO 
3 quizzes -jm 


Frog Points 
\

10
10104
440 (35* of points) 


Guaranteed Grade Guts 


Above 88% Out of Final Exam (Including Alchemy and 
Namegame) 


•A---"7.
Above 80% 


Below 50% <F--' 


,. .. . L. •.. 



D7C COUABORATIOK JSTATBHEHT FOR INTRODUCTORY COURSES 

FULL COLLABORATION* Collaboration is defined here to mean persons 
working together in joint intellectual effort. Collaboration «Soes not 


include copying another's work. These assignments* prepared either 


inside or outside of class,, may reflect collaboration with fellow 


students, faculty, or use of other coures». All such collaboration 
moisit be documented. 


TBB DEPARTMENT OF CHEMISTRY DO£« NOT ACCEPT COPIED WÙSK. 



• 10' 
• 
• 
Chemistry 131 SYLLABUB Fal l 1990 
Veck1. 
Pat e 
9 Aug 
10 Aug 
LennonUS APAT-l 
Mo. 
131 
12• 
Topic 
Matter and Measurement 
Atoms and Molecules 
Reading 
' ?-i.5 
*¿. '¡y¿.h 
Homework 
U.16,20,28,34 
36,42.61 
1.7.6,0.12.ia 
25,33,40,49,S8 
LAB: Scientific Method , Quiz #1 - 16 Aug 
• 13 Aug M-2 3 Chemical Formulais 3,1-3,5 3,5,13,17,22.30 
. 15' Aug M-3 4 Chcaical Foraulas 3.6-3.8 31,32,36,43,49,^1 
17 Aug H-4 5 Cas*£ 4.1-4.A 5,11,19,20,39,40, 
AFACT; 61-67 44 
3, LAB; Data Analyses 
20 Aug22 Aug 
T-4 
T-5 
6 
7 
Gase.-c 
Electronic Structure 
4,5-4,6AFACT: 65-92 
5.1-5.3 
45,48,56,62,71 
4,6,11,14,20,22 
• • 24 Aug T-6 a Electronic Structure S.4-5.6 26,30,36,30,45 
48,54 
LAB; 
27 Aug 
29 Aug 
Eml£«lan; 
M-7 9 
H-8 10 
Quiz #2 - 30 Aug 
Electronic Structure 
Periodic Table 
. 
5.7 64.66*68.78 
AFACT: 133-150 
6,3-6.4 5.6.11,14.X6.ZZ,24 
Si. Aug M-9 11 Periodic Table 6.5-6.6 26,30,36,42,60,63: 
5. LAB: HOME; Ctt #1 on Wednesday 5 Sep 95 (Chap 1-6J 
5 Sep H-10 12 Thermochemistry 7.1-7.2 4.B.lQ,14,*& 
- 6 Sep T-10 13 Thermochemistry 7,3-7.4 20.25,30,32,34,40 
7 Sep 
• • 
M-11 14 Theraochenistry 
^^^^^^^^j^UA 
7,5-7,7 49.50,55,56,60,68 
6. LAB: Enthalpy 
• 
.10 Sep 
12 Sep 
14 Sep 
T-HT-12Î-13 
15 
jjfi 
17 
Spontaneity of Reactions 
Spontaneity of Reactions 
Covalcnt Banding 
20.1-20.4 
20.5-2Û.O 
8.1-8.3 
1.3.6,10,16,24,30, 
34 
36.46.50,52,73 
2,12.14.19,20,22, 
26.26,34.42,46,60 
• 
. • • . • • • 
••' 
• • . • -• 
• : " . . • 
• • • ' 

• • ' • : • • ' . '• • : y ." -. 



-Il 
LAB: Energy Sianlator, Bring Geoaetry Klta to Letture; Quiz #3 - 20 Sep 
Ï7 Sep ÍIH14 IB Molecular Structure 9.1-9.2 Ü,6,10,14,24,26 
19 Sep7 ; M-15 19 Molecular Structure 9,3-9.5 32,37,40,44,46 
21 Sap M-%6 20 Liquids and Solids Í0.1-lp:2 2,4,6.8,15,16,20 

• * • •'. 
;LAB: Geometry, Bring Csoaetry Kits to Lab; Dr. BürUett 2S Sep 90, 1900 
24 Sep T-16 21 Liquids and Selids 10,3-10,5 22,24,27,32,34,36, 
y q a . _ ' 
26 Sep T-17 22 Liquids and Solid* 10,6-10,7 40,46,48,50,61,70 
26 Sep T-ja 23 Solutions 11,1-11.2 4,6.8,11,13,14,20 
9, LAB; Solutions; CR # 2 on Tuesday 2 Oct 90 (Chap 7-10. 20J 
< 
3 Oct M-20 24 Solutions 11.3-11.4 2B.32.36.33.42,46 
4 Oct 1-20 25 Air Po11 xition/Acid Rain 19.5 46,47,63,64 
..' 
A 5 
Oct H-21 26 Aqueous Reactions 12.1-12.32.4,6.6,13,14,16, 
16.22 
lÓ.y LAB: Acid Rain ; Quia « 4 - 11 Out 90 
: id Oct M'22 27 Aqueous Reactions 12.4-12.5 24,26,36.3&,42.4&, 
64 
it Oct M-23 28 Gaseous Equilibrium 13.1^13.3 . l,2>3,fi, 9,12,l4,u> 
-.117 
LAB: Deteruinâtlorn of Kc; Quia #5-1 8 Oct 90 
. --• . . . • . • • . . . 
IS Oct . T-23 29 Gaseous Equlllbrlua 13.4-13.520.7 
13.26,36.33,43,44, 
&2,ss.S6.se 
17 Oct T-24 30 Acids and Bases 14.1-14,2 5.6.12.16.24.28.30 
;'•'• • • •. 38,44fS0rSSi56,63, 
¿9 Oct T-25 31 Acids and Bases 14.3-14.5 69,76 
Á12.'' LAB: Acid Base Titration; GR # 3 on Friday 26 Oct 90 (Chaps 11-15, 19,5) 
22 Oct H-26 32 Acid Base/FFT Equilibria 15.1-15v? 4,6,14,20,22,34 
• • ' • • " ' • ' 24 Oct H-27 33 Acid Base/PPT Equilibria IS. 3 26,38,44.52.56,60 
26 Oct H-28 34 Chttnistry of Photography Handout 
"V:; ; '•.".-." 
"•'; . ••'' 
• • • • -. • • • • • : . . • • -• ' • 
. : • • • . . -• • • . • 
A : •* • ••:.' .••'•'•: 7 
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. . .' . ' y '••:•• • . ;•.• ' 
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1. 1 
I?. LAB: Photography 
29 Oct T-23 35 Qualitative Analysis 17.1-17.3 l,2,3,6,iZ,:l6 
31 Oct T-29 3fi Qualitative Analysis 17.4-17. 5 2Z; 24,28,-3»,45 52 
2,4,10, lii, 24,28 
2 Nov T-30 37 Rata of Reaction 13.1-18.Í 36,¿0,44 
• Vi^HW^MM^ ^ 
14. . LAB; Qualltatlye Analysis; Quiz #6-8 Hey 90 
• H-31 30 Fflté of Reaction 19.3-13.5 46,48,50 
7 Nov M-32 39 Sate.of Reaction lfi.6-18.7 52,S3,62,64 
9 Nav H-33 40 Atmosphere 19.1-13.433,34.i9,4Bt5fi, 
47 

15,. LAB: Kinetics , Quia #7-1 5 Mov 90 
14 Nev K-34 41 Electcocheáistry 21,3-21.2 2.4.1^,22,24,3/. 
-— •
16/
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• — • 
LAB; 
K-35Hone 
42 Electrouhfloiietry 2i.3-?1. 4 40,48.56 
i;' 
•:.19 
Nov 
£1 Nov 
Z-35/ 43•"• I* 34 7 44 : 
Electrochemistry 
E l s et r oc he n i s t ry 
21.5-21,621.7AFACT: 20-23 
to, 67,63, 6-5 
20.29,42 
17. LAB: ELectrecheáistry\ CI # 4 on Friday 3Ü Mov 90 (Chap Î7-19,21> 
26 Nov K-Î7 4 S Organic Chemistry 26.1-26.2 1,3,4,7,14 
28 Hdv : K-38: 46 : Organic Cheat*try 26,3-26.4 íá,45,46,50 
30 Nov H-39 47 Organic Cheais try 26.S-2Ú.& 35>36,38.40.42 
18. LAB: Organic; Quix #8-6 Dec 90 
3 Dec5 Dee7 Pec 
T-39 
T-4Ú1 
t-41 
4B 
49 
50 
Polyaers 
Pol yuan 
Applied Chemistry Option 
27.1-77.227,3-27.5-y 
2.4,6,8,10.12 
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19. LAB: Hone 
10 Dec K-42 51 Applied Chenistry Option 
U Dec T-42 Clean-up 


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15 

INORGANIC NOMENCLATURE 
• chamistry 131 
Chemistry is a lot like living in a foreign country. If you 
speak the language, life is more fun! Learning inorganic 
nomenclature is the first step toward learning the language of 
chemistry. Learning the relationship between chemical formulas and 
chemical names is essential for understanding chemical problems. 
Inorganic nomenclature is a systematic way of naming the thousands of 
inorganic compounds which we encounter. 
Your textbook discusses nomenclature of inorganic compounds on 
pages 72-78. This handout is intended to supplement and complement 
your book. Another source of nomenclature rules and practice in the 
CAI lesson NAMES. Contact your instructor if you want a copy. 
BEFORE ME GE7T STARTED . . . 
There are certain elements* cations* and anions that you MUST 
know before you get started. MEMORIZE these names and formulas; it's 
as simple as thati For any name you should he able to write the 
formula and for any formula ysu should be able to write the name 
(including spelling.) 
. 
1. The first S6 elementa and Au, U, Ag* Sn, Hg, and Ph. 
2. Compounds with common names 
HaO water 
Jî3ûa hydrogen 
peroxide 
MMg ammonia 
. 
3.
.. . 
Cations 
••• y...' 
Hanfia Symbol • 
; 
• ; •. 
.-; 
• . 
•• 
; 
.-'. 
; 
' 
. • 
. 
Lithium ion. sodium ion, etc.
Beryllium ionf magnesium ion, etc.
Hydronium ionAluminum,III)
Iron(II), Iron(III) 
Copper(I) , copper(n) 
Silver(I) 
Mercury{I), Mercury(II) 
•• 
a 
Fe3 + . 
Cu + , 
. 
Fe ' + 
cu2 + 
, Kg?* 
• 
• 
• 
•7 
• 
• • 
• 
. 
•• 
•. 
Li*, Na + , K+, Rb+, Cs+ 
Be2*, Mg*+ , Caih , Srî+,Baï+ 
H* 
Al
AlAl
a+ 
++ 
.. 
• • • 
• •• 
• • •• 
. 
. • 
-. 
• -
1 
.' 
. :. 
• • 
. 
• 
•••'.'
• 
• ' • 
• .. • 
• • . 

: ;' r 

. . •••'• • • '••' -•' y . * . " . y ;•


3 + *+• 


Lead(II),Lead<IV) 
Fb Fb 
NHJ 


Ammonium 


4. Common Monatomic Anions 
name Symbol 
Mame Symbol 


Hydride H- 
Iodide 


F-


Fluoride 
oxide 


Cl-


Chloride 
Sulfide 


Br~ 


Bromide 


5. Common Polyatomic Anions (Two or more atoms}' 
Hams Formula 
Ifame Formula 

Nitrate NOi" Sulfate so4 * '• 
Permanganate: Mn04 " Carbonate coa *Hydroxide 
OH-Chromât* cro4 
?" 
Cyanide CH-phosphate s 


-


Perchlorate 
Peroxide 


cio*- 
oa 

clo^-


Thiocyanate 
Chlorate 


ÔCH


6* Acids 


Sulfuric HjSO* 
Perchloric HCIO4 


Chloric HClb9


Phosphor ie H s F O 4 


Mitric UNO3 
HydrObromic HBr 


Hydrochloric HCl 
KydrOiodiç HI 


7* Bases 


Li thiujn hydro* ide LÍO H Magnesium hydroxide Mg(oH)3 

Sodium hydroxide. NaoH Calcium hydroxide. Ca(OK)i 
Potassium.hydroxide KoH Strontium hydroxide sr(OH)a 
Cesium hydroxide CSOH 
Rubidium hydroxide RbOH 


There are four main classes of inorganic compounds we will learn 


how to name: 


1. Compounds made from single valence metals. 
2. Compounds made from multivalent metals. 
3. Compounds containing only nonmetallie elements. 
•.4. - Acids. 

;C0MFOyHS5 X*P£ ¿"ROtt SIMGLE VALENCE METALS 

. : In general, thes e compounds are made; « s ¿«g -tt leas t cr>s of the 

m-at^l s irs sr^^ s ¿ft ¿rscï ¿là , Th^S è ¡CKS-LAIE --1 ¡jr.t^j.t.^ fclect,-?îi'is ren^ïil y 
t they have a low ionizatio n eaergv s sa d always hav e oxidatio n numbers 
f-irîrresponding t o thei r .group number . Fo r exemple-* K, K&"> es , : etc . ,• 
taavs c.x.i.ijat:u3n. ¡nuïatïiiï:» o f -i-I oiily . .M i wt? hs."*& t o ^ o i s ç^eicii y th e 
catir a (th e ¿h>ni.c Bvats.ii .füiio^ d by th e ím i©;?., vîs sauist: a 13 o ensur e 
tha t t.îiô sum of the okidatión niissfcars ' i s zt=ra (fu r we are dealing ' 
with neutra l ¡Bv±%.zuL&&$ „ KflVhi atí s same?, examples; 

Nfcfê SCitLiUES ï.vâïr.v^.e 

Cajo . cesium- oxid e 

•• BaFi ' barium f iuoricie 

.HOW hydrcuPïSïit cyani c 

£ÏrsM j strontium n it ride 


Ce(OH)Î .calcium hydroxide 


LisPQ. 4 i-ithiiii K pisas^sbasfcier 

Hg#£> 4 )&&% li s a i i ¡¡jo s w 1 îat e 

; Now, you tr y a couple ï 

e . Na ? CrQ 4 • • ..•__.. • •. -7 •'. 
b . maçne»iu % florid a 
.lii-que&tiar c (a) , th e catio n i e Na* (sodium) axid the anion i s Cr©4?~7 
( chromât.**} . ïîiUfef t.V.4* ctsfnpijïma i s fê^i^L-^rs^ëi^ • Tn questio n (b) , 
¡Eagnsaiiim e^i.at s i n compound form a s Jfry".. ruisotviiie is- the F~ anion. 

H'o aaairie th e zt-o'jop.&'jiriñ. 3riQí.ftr?tl, we ïi^sd tw a 5"*» . Therefore.T tiift 
fûs^ïMl* fa r Bti^.cjíís&iuvíí fluo r id * i s ifc^i . 

I f « e hav e 2m ios-fic: ^íiT^HQ-iariíl fía a t i,*GiïtaiïL3 ÏWi?..v<.-gep a ï o-na >/it,h 
the «jat.íftvri, r.heív v s riô^S tt , ^eííi.í y v.he n^i&fcsr of'fcyti.î^çiéfc av.o ^ i n 
th e coisjpouuvi „ Fo r i^^r tance, 


: ' y HaltCOs . sodiu m hydrogen h^arfc'i^íi'ís 
Na2HPp4 síííilwss ía&^cí*y-HÍíx»gen phosphat e 
NafíaPO* ¡s edita s dihysírcgor í phosphat e 


i n th e ?tfcíív* ïMt&sftpLaa-. v « as s Greek prefix.befor e th e word "hydro?^7' 


ví'.J'V. •:': 


: ',; t-\ 

to specify the number of hydrogens. In case you've forgotten, here 


r


•are the prefixes: .. ;.,.'y-. ry.7y" ;A;*^ ! 7"7'7yy 
mono-1 (Omitted if no other prefixes are used) 
di-2 


• w. < 
rv ' •"••J-"' 

 -.•!•• • 

tri-3


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y-y deçà- 10 

•yyy -" ''. "."":y<y 
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Notice that for HaHCD** the number of hydrogens is not specified. 


Lsis because carbonate has a -2 charger if sodium (Na + ) combines 
with C03 
a*Y there is only ^room" for one hydrogen. Thus, the mono-
is inferred- in this case* ; :'"y:7,'v;' 


COMPOUNDS MADE FROM MULTIVALENT METALS (STOCK SYSTEM) yyt. 

Nonionic metals can usually have more than one oxidation state, 
Iron, for instance, can be found as Fe3.*.-. and asFei + . Most other 
transition metals can also assume more than one oxidation state. 
way we distinguish between compounds which differ in oxidation 
numbers is by using Roman numerals, which is known as the Stock 


. system. The Roman humeral indicates the oxidation number of the 
cation and is enclosed in parentheses immediately after the metà1 
. name. Take a look at these examples: 


•¡•J i . ; 
•¡•J i .; 
• - "" ••':•• ' • • / • • 
•'-'••' • ' • • • • • ; • . ; * 
y -t • 

pyy-7 ':-,x FeSO. yy iron(ir) sulfate yyyyyy yyyyy 

•ïy try. 
• !-."^ ""• •..-! •.' 
yyyy. Fo3 <ap4)3 iron(III) sulfate 77:7 

y y -Pbûï lead (IV) oxide 
«v y ' Sn(«0i)a 77, tin(II) nitrate ^^y y y7-;77 7:77 

L '•;-'••' 'y * ' : • ï .• • • "•'. ": 

'.':•'• ' '". , 

••.':'•. j v . 

.I".'-- ..'.I---¿ -CUCH copper (I) cyanide"' ::'-7.;: yy.:7: 

• ~ m-:''J 
Hg(OH>a mercury(II) hydroxide y.yyyyyy:

• y •• •• *-! 
'" " "" ^ " 1 

•• -•:-;.•...• £ •• • 

 •'-.•; •;';'. > 
A:y :
!&£l% i :

Note the sum of the oxidation numbers(s) of the metal exactly 
cancels the sum of the negative charge(s) of the anion. Note also 
that you need not memorize the oxidation number of the metal because 
you can always back calculate its value by knowing the charge on the 
anion 7 In other words* memorise the anions, both name and charge, 


_. y. -^: .y 

•'•••. •"'•:"*: .-; • "-. • -:-y.
yyy ' ^jÇ'7 

••:•• "^ ' '

Let 's trya coupléi : 

'• • • -T 

• • : m -y •'• ' '•''•i 

h }'•' • f:

V v. •; 

C-y ' .r.i'^.'V • _

^ '.'.•. . 

: ' ;«.•"•• • :' . . 

'y '^A*"CuSO4 _ _. _ 

Î . 

:

.• } '•• -•• •-•


.y

"•>'•' ; 

' -:••:• . ; • -* -;.-•-• -F .

.:. t>-iron(ill) carbonate 

— 

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T.-W --: • • ''."• 

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19 


in (a), remember to k«y off the anioni Since you've already 
memorized i?) the anions, you know that sulfate has a -2 charge. To 
make the compound neutral; you know that copper must have a + 2 


charge. Since copper ie multivalent, we will expías its oxidation 


state when we name the .compound: cocmerfll) sulfate. For (b) , the 

(III) tells us that the iron is in the >3 state fFeï+), Vîe knotf fram 
our memorized list of anions that carbonate has a -2 charge {CO3 ~) . 
To make a neutral iron(III) carbonate formula, we Must balance the 
negative and positive charge. Two Fe3*'s and three C03 
37's would 
balance the charge at +6 and ~<5: F*ai(C0?>3. 
AS with moat topics in chemistry, there are always some 
exceptions. A number of metals exist whith are not in Groups 1 or II 
and have one oxidation state. For these compounds you do not: specify 
the oxidation state in parenthesis. The exceptions ycu should kiiow 
are Ag(+1), Kif+2), Zn{+2), «ftd Al i+3). Here's some examples of 
what we're talking abouti 


: y ••: ,. AgtïnO* silver permanganate 

• y • 
• ' • •' . -• • -; . • • . ' . 
Hi(Cl04) 3 nickel perchlorate 

• .. •. . 
AIH3 aluminum hydride 

Hamember, we used no Roman numeral* in naming these exceptions. Note 
also that most transition metals can exhibit a +2 Jon due to 
ionisation of the 2s electrons* 


COMPOUNDS CONTAINING ONLY HONWETALL3C ELEMENTS 


Let's move to the opposite end of the periodic table. Here we 
find the noracetallic elementa which are defined as having a high 
(negative) electron affinity. In general, the nonr-etal 1 ic elements 
Include those elements to the right and above the dark, solid 
stair-step line on the periodic table that forms a diagonal from 
boron (B) to astatine (At)t Compounds formed ¡solely irom these 
elements will be covalent in nature, when naming these molecules, we 
use Greek prefixes because the oxidation number or each element is 
not always obvious. The compound name begins with the most 
electropositive atom (the least electronegative) and then the other 
element is named (once again with tha "ida" ending as in the anion). 
Let's look at some examples: 


CBr4 carbon tetrabromide 


FF3 phosphorus trifluoride 
K*203 -, di nitrogen pent oxide 


. • 


XeFfi xenon hexafluoride 


H«re are a couple for you to tryl 


a . N O • • 7 - :•• "• -•••• • • -• , -• : • • " • • • 
• : ' • • 
• -•'•" ."• 


'•'•'' . . . 


• .. "• • • y ' . 

b. /sulfur hexaíluaríá* 
•"-.•rfl.*"'^-" 
Tïi* formul a in $&) 1.aÍÁÍJ;.C¿Í.Í¿ML.S£Í?U,ÍÜÍ1SÍ^.-• 17 tbei-s i s oxly one o f th e 
ÍÍÍOÍS:Ü • « i e c t r;£*¡¡i s? & i t i V*Ï ¿t t s # a , t ïies "wt^ o ~-JVÎ pr « Ê i x c a n fe* ú i: ap c & '.?.., Th a 
M35dri6-* prefi x i s alway s used; for. t."r*j» sor e electrr^^g.'jiíiivt í stem , i n 

fbi , n c 5reíi n orr th e ssaiiii r i apil e s one «.ulEuïy ÍÍSK;? r'luííiri.tí& iñeans 
. «ií t 'F"* s.- í'jtereí^re f th e £«<£ÍRÍJI1UI Í& 5F$ . 

As a fina l jrüfrbe, iií& shoul d tsiàk afccMï. naüi^ ' acida . You shoi-Iíü 

.'••.• HMO s nitric acid 

Hi$®4 sulfu r Us acid 

Al l oxyaci.<?& fí^ád » í?ís-iítaifl¿.T¡íf asygcrO ^r ^ ÏI^I^Î Î ÎÏ íta r th e ' ñama 

of tb * ¿tri.loti» Vor «fxaaipi*, 

HCl*>«i jp#!ji?uhii>ric a-sld (£rt>« psï'sïslfcsrAit:* ion ; 

ïî^CitO^ cferoai e aeii J íís:<í* cñrcisi&ts i*¿ro 
Qtlí^s? «c-iíís hav e th.« sfcsa» ffiírmíílA a s SÍ^ÜS^ o f th e compounds 
discusse d abc^ e bu t hav-* ÍS dií'fereíi t rifliES, 


HJSÎ.(ÎU5 ) fey&E'OOhlíKE'iAS flWi.d 
HEjcíasf) feytSrstorftsiis aci a 


fíCif íáí| í • îïVdK-iX3y*i>ic *ai d 

Mosit. #£ th e ti^ a v « vll i ¡sud th«; aladre*ía t 1 OÍ* fo r .a^u^u s (w.f) afte r 
th e CCJs^Oïiîiiî t o i^Hca U tha t it ; ¿ s difcSfclvie-S a;: wate r ^niï is * 
ther«fí5r* , wn a*3 itf •. Fo r instaras, , ÍITiAq) wsul d fce calle d r//^^. i odi e 
acJLtf v#h«T-!î! *r\t m WÎÎIÎI S î?c neinsetf hydroge n iíx*i£í y 

HoEiií:nclavU*r*. i * »fi âapart&rct ¡>art # r th e lançsa ^ o f chèrc1atrv , 
Zi?t v^ e t w ftassirig tïie&e íítt&^^s - You'l l b * s««li ^ much sor e o í 

• tftesi. tÉiTisy^ï^ïi'^ ch e y*as;ÍEi 
•y\ Tfca ne^ t pag a i&fcr,cjsra:r.«í.&t,¡í3r* ^rf c «fee-fit. UÍKS thi s t o hel p 
y ou p r ^s ¡zti t*£ y OÍIÍ: r- raaencí -at^ r*t ¿sk 1 i I a l í í 

y '. 

.

.. 

XOKSStCL%?2&£ WORE SHEET 

tí«üít let' s seis how ïiiuch you4" ve learned ! Use th e E d lowing '^o 
practic e your nato&jicii&ture skills!' Ï 

Same this ggsllawlncy; 

(HH4)3CO3 ;-•___ 
>*Qfi 

Al20s 
BaCla ';.'•.;; ,__"• 
CÜ»S ;. .r...'.'" ' •.'.•'.•• 
F3ÏSO* " 
CO; 
(MKO^HPO* 
— . 
F*sOs ' .'•' ;7 
cai^toy, s 

¿¿X&^ei3BUi&l^f^£L_^e_£^A£Eill|£ï 

s&monium staïfat* 

.^—^^ •••—-• 

calcium oxide 

potassium permanganate 
ailver phosphate 


oispperiïT) a^ict© 
m&gjftesiuoe iodid e ,__„__ 
titaniiamiEVÎ chloride 

..-.••• -'•.:';-• 

'• 

• • • . ' • -• • 
.. • 

7 " 

• 
; .. . •', '•: 

'... 

• 

'•••'. , 

. .". '?:-'. 

• . 
• 
ÎÏ5Ô 
ïî2SO* 
UeH 
••' . 
line 
BeSO ¿'7 
Hl&r 
PE"s 
P4O1Í 
.: 
Helo* 
• y •. 
-
.-
barium iodid e 
sod i un; sulfid e 
ironiï.ï) sulfate 
. 
sulfur trioxide 
aluminum sulfide 
carfeonic &ci& 
vanadium(V) oxide 
• . . . • 
• •. •
• 
• •. • . 
• • • 
• 
• 
• . 
' • " • •.-:• ; 
. ' .•'"•• ' y'.. . . 
. . • • . 
' • • • ' ... ... . 
• . . • . . • • • • . . • 
.. 7' • ; • ; . 
• •• ' 
• 
• 
• 


Select! í i *¡ Hatead 
Chemistry 131 


$M$£HVS: 

To become proficient in using the scientifi c method, ihi s lab 

will 1&iiá you through tft* scient i fie aethod ?if ps-p-blero solving, YC;D 

will jn*ïïa observatiens, examine your observations jfer patterns , ísafce 

and verify hypotheses, pôr fiama aH pe ri usent ¡s that «ere de signad íor 

you. aíid design your cam experiments. Finally, ye*xi will be able to 

düfcfirBiin<3! the reaction pathvay for & !7-îocfc reaction," 

Frsrïcis 3a a .in suggested centuries &go that as a result: uf 

careful obsesvation of phenomena tïxër*: vault) always emerge a logical 

explanation s ofch*rs believe that aiKl^ss e sepe rinses Ration without a 

grec anací ved objective le a da nowhere* The scientis t works* in two 

direction»; he collect s data a^di fron his observât tares î,o propose© 

explaining hypotheses through a process oï induction, then he seeks 

to verify tfcie explana tien through-a prï>tresi& «f dedüictlon, I t i» fcfte 

ra¡re scientis t who has enauçft imagination tx> «reate new valid 

hypotheses, 3fcst se¡í£ttti*t» jus t ""get mileage" out of Sfcinifttiie el ¡HE1 s 

hypotheses by doing variations of his earp^i'.ispents. I t waji iorty 

S'ears after noted theoretician Albert Einstein proposed the 

convelía n of mass to energy that the invention s-f the atomic büi^ 

validated his hypothesis, 

Ths inductive method propelled iteienc* to prominence during the 

EfilightenEa-anr. peïriewl ci' the fterwisaance. Prsvioïîaly the ancient 

ñresks «osrfced entirely through deduction and the ficxltm of the. four 

*fci£»ftftt*M {Earth, Air* Fire, ana Water?„ I^fcer, the pursuit af the 

^Materia Prim*- py Alcfeeaist» testifie d to th* futilit y «.£ A pîiraly 

deductive process tha t begins with &.n idea and eîanies the result s of 

any attíjariiaent whisis d**a not tiarn out as ^radicted. £1; 

Modern science dates from the sixteen!:, century whe?i. system ti c 

observation asMl exp arisen t sieve lcp«& tb* i Ï? d'activé part of the 

proceüü u The foratatioR of a hypothesis i s an attempt ta ^usise why 

something happe313, tiHiica i t is inductiva. JOedíícti^e p'fedictiíín <¡>í

result s leads t o verification when the hypothesis à& tested . 

¡Repeated observât ion-predict i tîjn-verification-revision of a hypothesis 

liiâda teî em acceptable «xpifimation called ¡a thscry. This approach, 

in .its sittpleet ferm, consists of (several distiiitït steps Ï 

1. HaXiag ebeervett&ae. ïha observations may be Qualitative (the 
sky i® blue, BCT is fun) or quantitative ffche pressure i s 1 
atmosphère - thfc temperatura .1» 21° C) , Many aheteaiccil reactions 
involve changes in £&ïcr, eve lu t ton of $ gas, or formation oz a 
precipitate , However* r*ot al l reactions GOCTJ*' instantly, some take 
j&iautea, hours, dayef o?- even yaarg* (2} 
3y fc#aK£»g f©2 patterns i a the observation. This prcces-s aften 
result s in the forsaatioft ef a natural or seieEtifia law, fcr 
sasunple, studies of countless cheaicëL reactions have show^ that the 
substances present ¿after a rsüction ha^* the same tsta l eaass as tha t 

'•...,. • 2 3 


of substances present before the reaction took place. These 
Observations can be generalised as the law of conservation of mass. 


(2) y . 
3. Zonulating théories. À theory consists of. a set of assumptions 
put forth to explain the observed behavior of matter. At firsts the 
set of assumptions is called an synthesis, if the tentative 
hypothesis survives the tests of many experiments , we gain confidence 
in it* value and call it a theory. At this time, it is important to 
distinguish between an observation and a theory* An observation is a 
fact that endures time, while * t !**•;:>¡a y is an interpretation and may 
change when more facts are known. (2) 
4* Designing experiment* t& i.*«tftíie- theories. Typically* science 

is self-correcting* continuously testing its theories. We must 
continué to do experiments and refine our theories in light of new 
observations if We hope tu approach a more correct understanding of 


natarai phenomena. (2) 


lahen designing an experiment that involves several variables, 
one must change only one variable at a time* If you were to change 
two variables for the same experimentr it wou.é be impossible to 
determine which of the changed parameters caused the change in the 
observation. 


The preceding discussion describes the ideal scientific method. 
However, it is important to understand that science does not always 
progress smoothly and efficiently. Scientists are human; They have 
prejudices; they misinterpret data; they become emotionally attached 
to their theories and become non-objective. The scientific method is 
only as effective as the humans using it. It does not automatidally 
lead tc ^r^gress.' 


This experiment is designed to give you some practice in the 
ways of scientisfes. The scientiSt observes resuIts of procedures, 
¡EsJkes hypotheses to explain these results* verifies the hypotheces, 
des igns: experiments to isolate one v.* r í abl e, a nd t i ne 1 ly assemble sa 
theory about the reaction. The chemical reaction that we are 
investigating is called a "clock reaction." Chemicals are mixed and 
after an intervalof several minutes the sudden appearance of a blue 
color signals the completion of the procesen-a chemical reaction has 
taken place. You will, first carry out some simp Je combinations o£ 
solutions* record the results-, and try to ©xpIaiA in chemical terms 
why changée occur. As successive steps become more complex¿ look for 
consistent patterns and exceptions, verify changes, and finally try 
to- put together a description of the y roces ses which together make up 
the clock reaction. 

ssperiÉentel Information 


1. clean the pipets before and after the lab. 
2, Keep the pipete used for the iodine solution separated from the 


rest. Iodine will staiii the pipets: and we want to minimize thé 


number of pipets we stain. 


pH&gë' 2 


.;:)^'^*™y-'-?-y'y': y:y:.y yyyyy yyy :yyy 

Record all observations andyour explanations inthe space provided. 


Be sure to observe allreactions for at least 10minutes before 


assuming nothing has happened. Donotstir the solutions after 
initial mixing. 


l. Place about 3 drops ofKI (potassium iodide), KCl (potassium 
chloride)., (NH4) jS3Op, (ammonium peroxydisulfate) , Ka2s2o3 (sodium 
thiosulfate), andIa onseparate locations Ofyour reaction surface. 
Add 1 drop ofstarch solution toeach ofthese solutions. Record 
your observations below. 
Solution Observation 


a* KI, starch .7-• -•-,• 7 y. ._ ..,.•. ..-.' 


. b. KC1, stardh -,•••••,. •' '•'.••• •.-.•• ,..•..•;• .--7 -..••: A , 

'-e* y (KM.)Xa¿a$** starch 

•MM 11 • 

• A7d^y:-7ïWiS20î.^stérch. y 
yyy'i'^yyy ;iV/:':;étáxtdr;;; y'"/. ...yyy 

Starch indicates the presence ofwhat Species?, ,__ 


2. Try various combinations ofthe solutions used instep 1,taking 
them two atatime (use about 3drops each). Then addldropof 
starch toeach. Which combinations give apositive result 
(appearance ordisappearance ofablue color)? This isa keystepin 
analyzing the reaction. 


Solution ---yy -observation 


,] ,/ a. Ki; KC1>s^¿ch\ ^.7'A-y ...'•'.'-;-
--. y 7-7" : y 

b, iaf(HH4)2fl3o,; starch 

—*"*^"™**""^ 

yic.7 ^«¿i\:Na3 S3 Oiy..^at^y7.: 

yd y • y.K¿V7Í2> :i**M?*=V.''^-S C'y 

• * KClí ^MH4^^IBiQi> staireh 

'• y-;ft,' ;•JWÍ>;7«aJs>Ô>,7 starch yy 

r-y§ïy:'.kGiy íytyst*á^y yyyyy 

!™-™™""~ w 


ti-.-., (HH4) 3 SÍO* ;-i?a¿fliOs> starch 

i* (HH4) iiSiô«, I2V stAtóh : 

^'. HéaSjO^,vIa, starch/ 7__^. 

Page;3Ú 


7:y7;'7'y.7,.yy: 
. ; 23'


3. 
onseparate areas onthe reaction surface place: 
7(a) A 3 drops: of KI 
: (b) 3 drops ofKiandl drop ofNaiS3Oa 
''7:7;(^):;7'^;;d¿Óps of ;KC1;:7 


To each ofthese, add1drop ofstarch. Then add3 drops of 
(KH4)2S;Oi* Record your results below and explain why youthink a 
positive result was orwas notobtained, what seemed to be the 
effect ofNajSïO-j? 


fl^lution 
Observation 


à. Kl,(KH¿) aSgO/j, starch y • -" •
•• • . -• y: -' • •
•• • ••• '.' ..y 

by,:jti, Haj330,, (WH*)3SjO,,. 


starch 


" '\-r. • 

e. KCl, (KTIÎV) 2Sa<Ji-starch 
—^— 

Try7to Verify your hypothesis byexperimenting with these solutions. 
State clearly thé experiment you performed and your conclusions* : 


Briefly summarize what you think are the roles ofKI,KCl, 
(MH4)ïSïOBr NaaS203, andstarch* show your conclusions toyour 
instructor tohave the validity ofyour observations andhypotheses 
checked, y.7 


solution • •-' y-y y •••"•'•:;' Bttl* 

pyy^ •' '.A*.. A*I • 7 /:;;y..'i •":••77 

':*y ' *: -KCÍ • 

yyyyyyy. 

ypy. (WK*)2S2Oa y:--.y:¡~y'-.yy 

'••:'•'. • . ! . " ' v '"

y':'Úy ïlé'îSaO.a• 

'•—• • • 


•• ¿e> ; starch --.7••••'••' •y-'' y.7 : 
M M M^^^^ M

• * 

• y -•-. '> • v • -y7;7 y.f.Eajg
• ij!i l ï.V! , ;.'--:. , ->.. . :•' ••.••••-••,: . '-;••.:•.• J"..--. •! 
; V:

y.-'-'.•: '-• : "yyyyy-'•:: .•"',.'• L .y•-••.•;-• •'•••' .-. 

• ->••'; •'•••'. isüíi V -v.-'. ' •••• • -.' ••'•• : y *•• • ••:. -y y i f :

•-••.' '* •

••¡y.v-y *„.-.•• yy.-. • "•••• •'••v.'--, A y '••-• yy-.y


':' • .:•-•;;'v:.-i -•:\ i ¿--r. 


2b 

4 y-• NOW fo r the! • ful i c lock react ionÏ. •..-.'.' Hake up two sol ut i p rïîs à s 
described belov, •'; 

Solution ïi. Measure Si drops of 0.20M Kl, 4 drops of o. OOS M 
Na2S^03, and 1 drop of starch. 

Solution II: Measure 8 drops 'a§ 6..10M (NH<) iS ;Oa. Record the 
room temperature* 

Temperature : y...7 •_.••'Í. "'"•••' 

Quickly mix solution I and II recording the time of mixing. 


When the solution suddenly turns blue record the tine. Repeat this 
procédure until you get a constant time interval. 


Time of reaction; 


_^^^^^^^^^^^^ 


5. Refer to your results in step 3 and verify that they arc 
consistent with results in Step 4; if not, repeat any part about 
which you are unsure. Summarize What you think is occurring. 
6. Concentration Effect. The so-called "Law of Mass Action11 
suggests that ths rate of a chesleal reaction is affected by changes 
in the concentration of the reactants* Hypothesize what you think 
the results would be if you were to decrease by one-half the amounts 
of KI, (HH*)iS20Sr Na^SjOs, and starch? 
flftlntioa frtdiqttfl Bete 

"• •• 7e* "• - ' V T ..-•." " ", •."' : ,.• . 

.-..S^.e .-.. _ " * . • • • -• : -, • • • 


b. (NK^SaQ i -, ..;. 7 -• •-••• •' " ; . -•• y ; 7-, 
• ; c. . NajSiQi •••••'.-. . y . • y y .-.. y : -. r • • 
7 d.7 starch ••••••., •--• "-• • • • y' : 


Now test your hypotheses by designing an experiment. Remember that 
you must ahenge only one variable at a, time. If, for example, 10 
drops of KI are used instead of 20 drops, the total volume will be 
less and hence all concentrations would be altered. To remedy this 
y OU: must replace the mis sing710 drops. ; Ifthis is dórie with 10 drops 
of deionized water, the correct volume of solution would be obtained. 


Page 5 



¿1 

Cléàtly state your experiment and record your observations, 
PolwtlOD Tims Of reaction Observation 


y a.y'77 7y77 ; 

:'' :Py y-.J. 
•• 
7c,'. . •'. .. . . -. ' • • . •V 
'.'.• ' 
yy-'aiy. y ^^^^^^^^^T^^^ 
• /. 
Step * 
. 
Explain your results,; .. . 
.. . •..• • 

7* Temperature Effects* Run the full clock reaction (Step 4) ata 
different temperature. Record and explain your results. What was 
the temperature. 


Temperatures1 Time of reaction: 


yy-yyr y-•'.•.-•• . : •• . , | • • •• 

y^yc^ci^siomy 


How that you have been guided through a systematic analysis of 
chemical reactions, from simple binary combinations to complex 
processes, you should be able to suggest the actual reactions taking 

yy place in the "clock1* reaction, in addition, you should be able to 
discuss how changes in concentration and temperature affect reaction 
rates. These principles are important and will, therefore, be used 


77-throughout this course. The key point when designing your own 
experiment, as you will do in most of our labs, is to change only one 


y variable at, a time. 


Vyy'/ly H^nioin; Allev L., J. Ch»p* Ed.:-, 5&. 434 (1961) 

y •' 2 > Ziimàçhl, S t eveíi ":g. ;,. Cheai s try . 2 «d> Jte A i ngtùn, HÂt D iC: Heath 

'y-y mîid coapany, 1939, pp 2-4 "' y 

y'r -'7 'y',. ' "A y '• r : , J ,77: . 7,y.,y7; . ":,. -y'.'-, -y y-, : y .. ." 7 7 y •• -•-•" • '• "• 

à* y%9 9 É« t jc> n, -V tl 1 l#m L,,t Chemistry. Principles and Reactions, 
y.y'-'-'Or-iah'dQ:, fL, Satthderi College Publishing, 19S9. 


• >•. • . -• •. -y • 

Page 6 

rVï:

-v



Experimental Report: 7 A 7, 


1. Tell which of the following reactions are consistent with your 
observations and which are not* 
consistent <Y/M. 
— • S 2 Q a 7 -* s j o à ? •• •!"-—r^- > S O V f ~ 


-


x* + s2o8 
a 7 -—.--— > -"so:»;*_ 


a 


-^^^^^^^^^^^^^^^^^^^^^^^^^^™ BaOe1" + i~ :—T-*-—--> so4 - •+• la 


•-. ^ 


I2 + S;Oa 
?" —r™-.-—-> Iy + SO*2' 
r 
+ s2oá?7 7-*.--—--->; j2 + soja 


r^ ' -Ï • 

3Ey + 504 2~ T~r~Tr > 12 + S2 0ty " 

» 


c i -" * s z ó *2 * '•-—--•—> ei -2 :+ . s o42 y 

'Í;7.'••';+' starch^^r-^*^-* :I2-starch (blue color) 

2. Devise a reaction scheme (combination of reactions above) which 
might explain the results in the experiment. 
3 *••• HOW did Changes in concentrátioii 3nd> changes in temperature 
;afieOt the rate of reaction? 


!*" •••' 


H :•>**: -y: 

^ -i?,.: v 


: . -.• •. i-J 


yy'yy:y:"yyy'y[ %J~éi:^h^zÀ\^y y yy-y 

...";> 


"-"' •":"•' -'."V '•'." '• !J/-y '. 1: 'JS.'.V.:-".'. • 7 ;• ;K ^j-^ ..1 .i. ¡- • : : :J;.;,'•: \. J-: '• \ ' . ^ „ . •• •-...f^-;..^ '' '• J"-^ ' •': •... • " 

•y''-F *<•:'•'•-. \ ••: ••• • • ."'"-.-r;'"~ t .:,'••• • "••.•';•••••'• :/,"'" -" : '/.•' ". -.-:"... -' . _ • ."-'• " •"•.-.•" '"'•-• J"--.\L '"r-- ' ' ' •• 

y y ti 

•:>'"r-:.:-y.i:. 
. i i .-. -••_. 

.%&• 

^^^^^^~ 

;-. : , :".iYJ 

."••• •: -• Ï 

I' A ( ^ 

/•• y 

• y •• 
Pag* -a. 

.•".••AV . .: 

yy ïii'



JO 


DATA ANALYSIS 

chemistry 131 

INTRODUCTION . :77'7 ''y'/.w" y 


A common tool used to analyze experimental data is a graph. 
Producing a quality graph (or plot) is not as easy as you might 
think. Most of you have experience plotting points on a number line. 
Unfortunately, the proper methods may not have been reinforced or 
thoroughly developed. This Lab outlines the requirements for 
creating quality graphs from experimental data, and then discusses 
some background and tools for manipulating data tq produce acceptable 
graphs. y ., 


THE BASIC REQUIREMENTS OF GRAJHINQ 


If you examine a piece of quality graph paper„ you immediately 
notice it is divided into several different sized grids. Usually 
each -of these grids is denoted by differing darkness of the grid 
lines. The purpose of these divisions is to help you graph points 
accurately, and then extract data from your graph. NEVER use square 
ruled paper (usually reserved for engineering drawings} to produce 
graphs. Figure 1 shows the difference between these two types of 


paper. .••*. A-A y '.-.>'-• ••-•• 

1::±:::: 

•* 7 '.---• 


•. •. >
yy----


9


QUALITY GARBAGE 

Figure i 

l 7 PROPERLY SCALED AXES 

The first step to drawing an acceptable graph is to decide how 


the axes should be aligned on the graph paper. You should arrange 


the graph so that it fills the majority of the paper (nothing is 
.Worse in graphing than a postage stamp sized graph on a sea of empty 


paper). But which axis should be the "horizontal" axis, and which 


the «vertical"? 7 - < " ;77 


Easy enough. Those in the know have assigned special names to 
data they wish to graph. All you need to do is follow their lead, 
Examine yóur data and decide which set is not affected by the7 
experiment. This is the independent: data, or the independent 


GRAPHING 1 


'•'.:'y..'!j! 

• y y 

. -: 


•. • 


. • •• 


il 

variable. Usually, this is a quantity that you control or set during 
the experiment, such as temperature, pressure* volume, or an 
uncontrollable quantity, like time. The remaining data is what you 
observed as a result of your experiment. This is the dependent 
variable. ALWAYS plot the dependent variable vs. the independent 
variable. Put another way* the dependent variable is plotted on the 
Y-axis, while the independent variable is plotted on the X-axis. 


Now you are ready to draw the axes on your graph paper* First, 


•

each axis should be drawn approximately one inch from the edge of the 
graph, never on the very edge itself. This allows you to write the 
scales and titles of each axis on the graph as opposed tc writing on 
the borders. The second requirement is that the X-axis should always 
be placed so that it runs parallel to the bittoa or right side of the 
page, never on the top or left side, Figure 2 shows examples of the 
correct and incorrect manner of placing the axis on the graph papâr. 

•. • 
• •• 
• • • " . -. A• 
• 
• • • 
• • . 
• 
. . • . . 
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y y T-*sT" " 
y •••
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.: B • , . • • ; . • . , 
X-aíif v • :
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• • 
• • • • • • • -. • = • • • • . . . 
. . ' • • ; . 
. ' • • • •: • 
• • • • • 
Figure 2 • .• •.. 

In additionf you should scale the axes so that your plot or line 
will fill most of the graph paper. If you can increase the magnitude 
Of your, scale (i.e.. Spread it Otit) and Still fit the axis on the 
page, DO IT. Examine your data and identify the range of values you 
need to plot. Fotf example, your values.for the independent variable 


(X) range from 400 to SCO units, and your graph paper has 35 major 
divisions along the X-axis you've just drawn. The range of data {400 
-800 + 400ï divided by the number of major divisions (25) is 16. 
Thus, you would **fillH the graph if every major division equalled 16 
units. Great, but aost graph paper has 10 minor divisions per major 
division (see Figure 3). ïJcing 16 units for a major division is 
awkward. Instead, use only 20 major divisions to represent your 
range of 400 units. Now, the maior divisions represent 20 units 
each, and you've filled your graph paper with graphI 
Two other tools for filling the page are: 1) Do not use zero 
as the origin if your data doesn't go to sero (as in our above 
example), or to rotate the page and place the X-axis along the long 
side of the paper* 


Finally, the minor divisions of your axis should indícate the 
same number of significant figures as your data contains. If your 
measured value is 1.93, you have to be able to read 1.93 on your 


:-gráp*i¿7 7" 


GRAPHING 2 


•y • 


y¿ 

LABELS 


Both axes must be labeled* The label should be a one or two 
word description of the quantity represented by the axis, and the 
units of the measured quantity. As an example, assume you have 
measured pressure as a function of time. The Y-axis would be labeled 
"Pressure (torr)"(the dependent variable} and the X-axis labeled 
"Time(seconds)"(the independent variable). 


y When choosing the labels, it is important to remember the 
scaling chosen when you drew the axes. Consider the following set of 
data: 


TJME (seconds\ • y 
PRESSUR E ftorr. 

•


0 750 


• y • • 
;io • : 870 
50 930 
100 13 30 
150 2010 


It may be difficult, and it would be inconvenient, to write the A 
number 2010 on the Y-axis. Instead, change the number to 2.010 x io3 
and label the axis with single whole numbers. This has been done in 
Figure 3, Each whole number on the axis actually represents a much 
larger number. You have now chosen a different way to represent the 
scale on the graph. 


Note that in order to plot the data as a number less than ten, 
we multiplied it by lo-3 . Thus, any number we read from the plot is 
actually Y X 10 where Y' is the value read from the graph. 
The new label would be "Pressure x 10'3 (torr)**, indicating that the 
actual value of pressure Was multiplied by 10"3 before being plotted. 


3. PLOTTING POINTS 
Now that you have correctly drawn and labeled each axis, you are 
ready to plot the data. Always use a sharp pencil or fine point 
writing instrument to place the data points on the graph. Around 
each point draw a small geometric figure such as a circle, square, 
rectangle, triangle, etc* It is best to use a stencil to do this. 
These geometric figures will help to correlate multiple sets of data 
on the same graph into groups, as well as help you locate the data 
points. Each set of data now has its own geometric figure, which you 
should list in a legend under the title block of the graph. 


4. TBS CORVE (OR LINE) 
Plotted data is of little use if some trend or relationship 
cannot be established. To see a trend in the data clearly, we draw a 
curve (which may be a straight line) through the data* There are two 
general ways for drawing this curve: the "eye-ball1* method, and the 
mathematical fit method. Let's first discuss the "eye-ball** method. 


/••: . 


: yfy . 
GRAPHING 3 
,y -


Fr£U££3 

GRAPHING 4 

BEST COPV AVAILABLE 


Eye-ball Method 


NEVE R play "connect the dots" with your data. The object is to 
draw a smooth curve which best represents your data. Two types of 
tools are available to help you do this: rulers and French curves. 
Examine your data and determine (by "eye-ball") if a straight line or 
a curved line best matches your data. 


If a straight line is required, then the ruler is the tool you 
need. 


If you believe a curved line is the answer, then a French curve 
will best suit your needs. Choose the curve that allows you to draw 
the best fit to your data (it is not uncommon to use two or three 
different French Curves to construct a single curved line). Align 
the curve you chose to three data points at a time. It is not 
necessary that the curve pass through each point, you might need to 
select an "average" position with respect to several data points. 
This requires that even though you are only matching three points, 
you must keep in mind the entire set of data while constructing your 
graph. 

If your data doesn't appear to be either a straight line or a 
curve. Of if you're just not sure, then the next method of fitting 
your data to a curve should be used. In this method we use a 
mathematical model as the tool to fit the data. 


Mathematical Fit Method 


There are many methods to mathematically manipulate experimental 
data* You may have heard of some of these manipulations as least 
squares, power curve, linear regression, etc. Any of these are very 
useful tools, though some are much more complicated than others. An 
excellent source book is Curve Fitting for Programmable Calculators 
by William H. Kolg, It contains in-depth discussions of several 
curve fitting tools, and programs for programmable calculators (or 
personal computers) to do the curve fitting analysis. I will limit 
this discussion to the most common form, linear regression analysis. 


Linear regression analysis, or curve fit, is used when and only 
when you are reasonably sure your data represents a straight line of 
the form y *» ax + b. Many calculators have a linear regression 
analysis function built in, or m*y be programmed to perform the 
analysis using three mathematical formulas. These three formula are 
used to determine the slope (m), the y-intercept (b), and the 
"goodness" of the fit (r*). The equations are listed below if you do 
not have a calculator to do this for you. 


n 

n m .. •.——— — — b -' 

s*i* : £*LL2 
n 


GRAPHING 5 



:[%.yi7y Sf^i] 

r=7l


myy : iBii2]^2 -&yu*} 

n 

7The .coefficient of fit, r*, is an indication of how well the 
data fits a linear graph* Generally, the closer r* is to 1.00, the 
better the fit of the data to a straight line. In each of the 
equations n is the number of data points. Recall that £ is the 
summation symbol. For X = 1,2, and 3; ixt means 1 + 2 + 3, or 6. Be 
careful to note the difference between IXi2 and (ïXi)2 . The first 
means l2 + 2 2 + 32 = 14, while the second would mean ( 1 + 2 + 3)2 


If your data does not represent a straight line, then you should 
resort to the other tools available. These include the least squares 
analysis, logarithmic, power, or polynomial curve fitting models in 
order to obtain a mathematical relationship for the data. Detailed 
information on these tools may be found in many textbooks on 
Statistics, as well as the source book I've already mentioned by w.M. 
Kolg. 


Computer Based Spreadsheets 


Computer based spreadsheets make the handling and manipulation 
of large amounts of data a bréese. The most useful ones include a 
graphing package that will take any data you input, and produce 
acceptable graphs. During the laboratory period you will make 
several plots and perform curve fitting with the aid of a \ 
spread sheet. ,.''.'.••' 


5. TITLE BLOCK 
The title block contains at least two parts: the graph title 
and data points legend (as seen in Figure 3) . Make your title ay 
short, clear description of the graph's contents. This makes it; 
easier to reference in a text, as well as provide the viewer a 
starting knowledge of the data portrayed. Always place the title 
block near the top of your graph (remember, the x-axis is at the 
bottom). 


-.•• .• • • <•. 

GRAPHING 6 



PROCBPTJRB 

1. Using a spreadsheet graph the following data set. If the plot 
appears linear, find the slope and intercept by curve fitting and 
report the equation of the line in the form y = mx + b. Make sure 
you obtain a printout of your data, graph, and curve fit. 
dependent Variable Independent variable 


... .

•'••-. • 

-, i •: . 7;.-. '••' :•:..-:' •.;."•' ! " A 0 

..'=i A. 7 -


•'•; 2 • 

• • • . : • • • . • • : "" " ' •

J-• ' ''* >•-* '

-y 13 

22 •:'&••• • 
: 'A . ••.:•.'.: • .•;'••:".'' • ."•. ' '•'•'. V.

31 10 

" "••• •' . 

''...y-y 34 11 

•"•If . v " • •. ."""•.-. • 

y -..-43 14 

•:. 46 15 
: 77- 52 1 7 ••;-. '.. y

•., 
;•'• -'• :.•' . •' y " -• '

55 18 

2. Now plot the following data on a separate graph. 
dependent variable independent variable 


y y :..y 

•:;:5y 7-17 
2 
3 

••:• 78.7 

•/«"•' ' "7. 
40 6 

-•-.g".'; 

85 7:7


• -3 y •••:-: y y 

125 i i 
:•.::;;;?:. 229 'yy. 1 5 

tlot appear linear? •': y. 

This data can be fit to the general form y-x n + b. in order to 
obtain a linear graph you must take the logarithm of the data points 
to make the general form log y=n log x+ log b. 


use the spreadsheet to take the logarithm of your data points 
and replot the data. Mow use the curve fitting routine to find the 
slope and intercept of lag y=n log x•+ log b. Report the equation 
for the data in the form y=x n + b. Hake sure you obtain your 
printouts.. yy


• \ '•>• . •'•. •.'• • 
.:•• .-...• -:¿-• 

. : '..;. ' y . 

y . 
•y^y^:,í^: -: .;.y 
••v. • , • f __1 .^:: 
•^-y\ -t" \.^ '• 
GRAPHING 7 ':,':•' -?7Vy 
: J"--.. 
....-.".V. ::. .j.r.v.i :.• \ , ':'.V' 


3v 7 We'll how look at some typical data for a lab you will perform 
later this semester - the determination of an equilibrium constant, 
Ke. Thermodynamic information (aH and ¿s) can be found by measuring 
Kc at a known temperature and using the equation 


lu K -V-AH Q J + _AS 


H I rJ R 


R is the gas constant, 8.31 J/mol K and T is temperature in 
Kelvin. 


Make an appropriate graph for the following data and find AH and 


ILB. 

T (*C) 


.131 21 

ixs Í 7 

laz 37 

• . • • 
• A....... 
. . .•.•. 


y 


. . • •• i 


•.. • 


: •• •• . 7 


. .. _ :.-• 

GRAPHING a 



• ' • •. 
yy M 

PRELAB GRAPHIKQ ASSIGNMENT 


For the following sets of data, produce two plots on the same 
piece of graph paper plotting volume as the dependent variable and 
temperature as the independent variable. Be sure to draw the 
appropriate "smooth curve" through each set of data. If the plot 
appears to be linear, perform a linear regression analysis by hand y-
using the equations on p.5 and provide the mathematical relationship 


for that set of data. 


• y y .• .
• . . •. 
. • .
" •••• • • . : . : • • • •

Data Set #1


Volume (ml)


Q.2
Ó-6
0.9
1*3
1.7


2.0


2.2


2.3
2.4.
2.5
. :-.••'.•'

: : • • • . 

 Temp (K)


 298


 304


 314


 320


 334 
350 
359


 370 
381 
392

 '• ••• '•

 . 
.-
• 
: 
.' 
..., 
•. . . 
'• y.• 
.. . 
. 
. • 
•..• 
• ' . . • . . ' 
' 
Data Set #2 

 Volume {ml}


 1.0 
1.2 
1.5 
1;6 
1.8 
¿;0 
2.2 
2.3 
2, 5 
2.7 
• '. • . . . . . . • • • . ' " 
• • 
Temp <K) 
• • 
• 
292 
30¿ 
312 
322 
330 
337 
345 
352 • 
361 
370 
• • . : • 

• • ,. y •--.••. . -. • . • . • • : • A . ..: •• -y :. • 

• • • . • • • • • . • : . . • : : . • • , . • : . . . . • _ 
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• 
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GRAPHING 9 



y'm. 

EMISSION SPECTROSCOPY 
Chemistry 131 


XBSOJRY. 


Spectroscopy employs radiation to probe the structure of atoms 
and molecules. There are many types of spectroscopy. Each type 
probes different limitations. However, visible light spectroscopy 
is the easiest because we can use our eyes as the detector. 


Joseph Fraunhofer, in 1814, was the first to make spectroscopic 
observations while studying the sun* Later, scientist turned the 
spectroscope to study flames* In 188S, Balmer recorded a series of 
visible lines which were emitted from atomic hydrogen. This 
spectrum was different from other spectra which had been observed. 
The other spectra contained all of the colors (similar to a rainbow) 
while the spectrum of atomic hydrogen only contained a few narrow 


(sharp) lines of only a few colors. Analysis of other elements 


showed that each type of atom had its own unique spectrum but each 


consisted only of narrow lines. 


In analyzing the visible spectrus of atomic hydrogen, J. R. 
Rydberg devised a mathematical equation which summarized the 
observed wavelengths* 

'.iW


i 

(-1) 


*Vl hitill" 

In equation (1), * is the wavelength of the observed line (in 
meters)# nnnli - 2, R is a constant which is called the Rydberg 
constant (1.09677 x 10T m"1 ) , and niïlitia] = 3,4,5.. .where each 
value of n¡mt¡,i predicted a different x. 


Other scientists found other series of lines in the IR and UV 
regions of the specti'um. y Each 'series corresponded to an electron 
"falling" from a higher energy level down to n i„,i = 3 (IR) and 
"Finni -1 ÍVTV). Thus, one simplistic equation fully described the 
complete spectrum for atomic hydrogen. But what did this say about 
what the hydrogen atom was doing? 


Niels Bohr hit upon the model for hydrogen which successfully 
accounted for the known experimental results* In analyzing the 
visible spectrum of atomic hydrogen in 1913, Bohr found that an 
energy of an electron ina given orbit is 


Ë = ™ 7 <n > 1,2.3...) ;.'•;••,"-• 


->•••'-y '-y' Emission 1 \--. ;-• ;•• 

y y y .: .-•••'•:•:•'-• '.'••:'••: 7 V'.-: 77.;-A''7-7.7--77 y-y. ;.:7ï: : .-7'v;7 ' ••••:.-7A ••••• ': y.7 . ;• : yy y yA-yyy.y: 



4U 


where B is a constant equal to -2.179 x I0~lff J/photon. To find the 
energy of a transition of an electron going from an initial higher 
energy level to a final lower energy level, equation 2 is used-
This results in 


• AE •- Enitii i y '•'. Ë f inrnt: <3J 
AE = -B (4) 


'ihlt tal 


rinn J 


rearranging, substituting in the constant B, and dropping the &, 
equation 4 becomes 

'['-. 2.179 x Ù.Î* f 1 ;• - '--i-1 -1 Í5) 

v nfln*l AiDlilil 

To find £ in the common unit of kilojoules per mole, equation 5 

becomes. 

photons [ . 1 
Photon mo ie 

£ „ z,i79 K KJ-ÜJ 6.022 x 10S3pht lkJ 1 1

103J "final "initia l 


kJ 


J (6) 

lnfi n.l "inlii.lmole 

Once we know the energy of a photon, we can readily calculate the 
wavelength of the photon using the relationship 


« he 


E =CO 


T 


where h is Planck's constant (h = 6.63 x 10'3 4 J.sec) and c is the 
speed of light (2.998 x 10s z?r) - The frequency (v) can be 
determined by making use of the relationships between the speed of 


1ight, wave1ength and frequency. ; 


(Û) 


and thus, 


E = hV Í9) 


Emission 2 


-,-\ :FZ\-• >


•T -• K -..'.r'. 

%% 

In summary, the emission levels of hydrogen are the result of an 
excited electron from a higher energy level "relaxing" to a lower 
energy level. The energy is released in the form of electromagnetic 
rad i at ion.7- The fó11owing diagram shows some of the energ y ; 1eveIs 
and transitions of the hydrogen atom. Lines in the Balmer series 
(visible light) arise from transitions from upper energy levels 


(n>2) to n=2 level. The Lyman series lines (UV) result from 
transitions from higher energy levels (n>l) to n = l. The Paschen 


aeries lines (IR) are the result of transitions from higher energy 
:1eyels'A'fn>3) •, tp7n^3;7'levéi. 7 


Note that this concept worked well for the hydrogen atom, but 
hot for' any other : atom7 at al111 

Energy Transition of Hydrogen Atom 


n-5 
y •-••-' • •'• n-4 
I Faachen 
E1ÍERCV •• " • • :• 
•yí 1 r }f * Ser±es • :r tí-3 
y •_ 
-Sal mer 
Serles ;; >• 
1 üf--> f y.yy n-2 
' • : . 

Lyman Serie* 


ty M

1ÜÜ n-1 

(Ground state ) 

• Spectroscopes 
You have already had experience with various spectroscopic 7 
devices. Your experiences probably include rainbows and prisms. 
From these you should know the names we associate with various 
regions of the visible spectrum* You should also recognize that 
certain procedures are required for proper opération* "obviously the 
proper viewing method for a rainbow is different from the viewing 
method for a prism. 


•y Thé: spectroscope requ ires a min imum of Atwji» çomp onents:';: ¿¡i 
dispersive medium to break the light apart by wavelength (or 
frequency} and a measuring device. Our spectroscope uses a 
diffraction grating which is a piece of clear plastic with thousands 
of microscopic scratches» Our measuring device7 is a glass rod. 
> t-

Emission :3 

••> : ••••. •-. , 


Both of these are mounted on a cardboard box which allows for proper 
control and reliable measurements* 


EXPERIMENTAL " 


••/-.. yy -:'•.:••" '.•'.•' •• y y ...; .-• /. .:• .; y • . • . • A ..yy . ••:: y. -

Surpose 


.l'ï Gain an understanding,'ôf; instruments by operating *• 
sp e et roscope. 


2) Understand the operating parameters of a visible light 
spectroscope. 
"". • '7 ' 7'7 " .. ' 7 . •'."'. '•'• ' 7 . /.••'. y y •.:•.••"•'A • • 'y. ••'•*••'.-.• '•. ': -' "•


3) observe* record and analyze atomic spectra. 


•;• Pr o o s dur e'••.. 
Your spectroscope box is pre-assembled. The remaining tasks 
are to attach the diffraction grating and install the glass 
measuring rod. 


> Class Rod 


A / 


Pigure 1. spectroscope 


If instructed, remove and discard the small piece of 
transparent plastic taped over the "eyepiece" opening (see Figure 1) 
of your spectroscope, mien your instructor gives you a new piece of 
grating, handle it carefully. Touch only the corners. Touching the 
middle allows oil from your fingers to cover the tiny groves::; 


Attach the grating only after performing the following 
procedure. Touching only the corners of the gratingr hold the 
grating over the "eyepiece" opening* Now point the slit opening 
toward a fluorescent light and look through the grating into the 
"eyepiece" opening. If your grating is positioned correctly you 
should see the visible spectrum spread out to the left or right side 
of the slit. If you don't see the spectrum, rotate your grating 1/4 
turn (90s) and look again. When you see the spectrum spread-out as 
shown in Figure 2, carefully tape the corners of the grating 
in-place* If you still don't see the spectrum after rotating the 
diffraction grating, get help from your instructor* 

Emission 4 



greeti 

blue 
orange/ye11 o v 

Figure 2. Fluorescent Light, Mercury spectrum 

Next obtain a glass rod from your instructor. The glass rod 
should have a narrow scratch near the middle. If you cannot find 
the scratch ask your instructor for help in creating your own 
scratch. Place the glass rod through the two holes in the side of 
the spectroscope box near the slit end. Look through the grating 
and verify the scratch on the glass rod is visible near the spectrum 
and that you can move the scratch all the way from the red-end to 
the blue-end of the spectrum* Now you're rea<iy to use your 
spectroscope. 


yAdjusting the Slit 


Look closely at the spectrum your spectroscope produces from a 
fluorescent lamp. In the fluorescent lamp, mercury atoms are 
excited by the flow of electricity through the lamp. Excited 
mercury atoms emit photons of light of various energies (or 
wavelengths) including W (ultraviolet). A mixture of powdered 
chemicals on the inside surface of the lamp absorbs this light and 
emits many different wavelengths. The many wavelengths, combined, 
produce white light - the rainbow like spectrum in your 
spectroscope. But included in the spectrum are three of the 
original mercury lines shining slightly brighter than the rest, see 
Figure 3. y


oratige/yel low blue 
58D nu 436 run 

Figure 3. Kereury Calibration Spectrum 


You will use the mercury lines to observe the effect of 
changing the width of the slit opening of your spectroscope. Start 


Emission 5 



"•'•'••..'•' -y 

by removing the two pieces of black electrical tape from the slit 
opening. What happened to the three mercury lines? 


How, replace the electrical tape but make the slit as narrow as 
possible. Describe the appearance of the three mercury lines. 


Now adjust the width of the slit for the best combination of 
sharpness and brightness of the mercury lines. Measure and record 
the width here. 


Calibration 


*• . : ;..-•:• £ . 


The calibration of your spectroscope simply means measuring how 
well your spectroscope separates wavelengths. This is done by 
measuring the distance between known emission lines. A convenient 
source of emission lines is the mercury fluorescent lamp. Look 
through the eyepiece of your spectroscope and find the three mercury 


•• • . •" • • "•••.••"• '•' • ' v 


y>:.yy y 

Emission 6 



:y&5


lines. Hove the glass rod so that the scratch is aligned with one 
of the emission lines. Now measure the length of the remaining 
glass rod on one side of the box. Align the scratch on the next 
emission line. Again, measure the length of the glass rod. (Be sure 
to measure from the same side of the box,. Repeat this procedure y 
for thé final71ine. Record the data in the following box. 


Glass 


Wavelength 


Hg BinIsSion Line Rod 


(hii) 


Distance (rem) 


blue 436 


green 


SA*. 


orange/ye11ov 580 


The calibration curve is determined by plotting glass rod 
distance (x) versus wavelength (y). This curve will be used tp 
determine the wavelengths of the hydrogen emission lines and others. 
Figure 4 shows a calibration curve with the hydrogen lines 
superimposed on the curve. 


"Í-: 


• Í. / 
Emission 7 



Mercury Calibration 
ip 
VI 
I » 
wavelength 
A(na) 
extrap lin 
Hg data 
H data 
. 1 . • . ..!• ' 
yy - 
"J* . . - y • 
Rjod Distance (ma) 
Figure 4 
¿/¿> 4 > 

• • .' "! 

• U 

?. . • •• 


Hydrogen Atom Emission 


Look at the hydrogen discharge tube with your spectroscope. 
Describe what you see. (Be sure to include the number of lines, the 
color of each line, the brightness of each line, the thickness of 
each line, the rod-position measurement of each line, and any other 
properties you notice.) A sketch may be helpful. Prepare a tabic 
in the space below to collect your data, Leave some space to 
write-in the actual wavelength which you will interpolate from the 
calibration graph* 


Helium Atom Emission 


Repeat the same procedure for the helium discharge tube. 


Neon spectrum 


Repeat for neon. You may not be able to record all of the 
Observed spectmm. 


• y • Ï . 
Emission: 9 



SrCiz Flame Spectrum 
Repeat for SrCl3 Flame Spectrum. You may not be able to record 


all of the observed spectrum. 


• . • 


ICI Flame Spectrum ••-• 


Repeat for KCl Flame Spectrum. You may not be able to record 


all of the observed spectrum. 


CSC12 Flam* spectrum A 
Repeat for CaCl7 Flame Spectrum. You may not be able to record 


all of the observed spectrum. 


' • A . 


nací Flam» Spectrum 


Repeat for NaCl Flame Spectrum. You may not be able to record 
ail of the observed spectrum. .7. 


• • . 


• . 
y y 


Emission 10 



CONCLUSIONS 


;J.> Does your calibration line have a positive slope or negative 
slope? Does it matter? Why? 


'A. . • y - ;.;• : y "•.-.•. • ••••• .-.A ; ,-. . y-.-..-.. . . ... •'.•.. • . : •'•••••-•. 

• y -. : . -y -. • : • . • • -. ' . . -. . • . . • • •* . .-. . ••• • • • . • • • • • • -• : • • : .-. -. -. . • • . . • • 
2. Does the straight calibration line go through all of the data 
points from the mercury spectrum. If not, why? 
3. Calculate the energy associated for each hydrogen emission line. 
(use equation 7)7, 


':' • • : . . • • . • • : •• ' • . 


. . -. • • . • • • • • ; • • ' 


: • • •••... 


. : . • .• •. -. • • :• 

yy-.yy -•-•.•'••. 

4. Calculate the nSllitit l value from equation 1 for each hydrogen 
spectrum line. Does the lowest n iritJt l value result from the 
longest wavelength or the shortest wavelength? What is the color of 
the line associated with the lowest niniti*i value? 
• : 
5. Calculate the percent error of your experimentally determined 
hydrogen emission lines versus the theoretical hydrogen emission 


lines. 


• •'.• • y-.-' 


Emission 11 



50 


FREÍAS QUESTIONS 


Complete the following questions before coming to the lab. 
Many of the answers will be found in this handout. Your instructor 
may require you to have this prior to beginning the actual 
experiment, 


l. What wavelength corresponds to the low energy end of the 
visible spectrum? 
• .
yy": y y. . -• -.-.. 

• • :•'.' 


• • .. . 


. • • 


• • • : • ' • . • • ' . . . 


.


: A • • • • ' • • ' •'.''' • .'• : • -. . . . . . 
• • • '.'.-' 

: A .; . • :. • • ' . • . • • . ' . ' • ' • 


... ;•• : .--. •;-.. . yyy.y, 

2. What wavelength corresponds to the high energy end of the 
visible spectrum? 
: • • • • . • • • • • . y y y' • • . ' • . • , . .•••• • 
• ' •. ' ; 

• • • • . -. • • • • • • . ' • .. y • . : • • • . • • . -• • • • . -. -. • 
-.... ' 

• . • ' . . " • • • •' . : . . . . . • • . . • . •. : . -• • • • . . -. • . 
. . . 
• • . • • • • • • • ' • • . 

yy' ^ -. -. • • •• •••:•• ,, - •:,-, ' :••.- 7 7 ,. • ,.. • • ; 7 ••••,,.••7; .;. y . 

y y •.-• . ¿ .. •• -y -y -7 . ..- .. : . 7 ••..-.-•... • ••• .-, -.-. •;• ; 

• • y y : ••...•. . A .., . ;.-• •:•.. : 
• . . . • • y y •••• • y 

• -' . • • • A . y • • • • . _ • • • , , . . . : , . 
. • • -. . -. . .• •• • • • -. • . . . . 


-. . . • ' . . . • 7-• . • • • 
'• i "-"• : . . • • 

3. what is the function of the diffraction grating on the 
spectroscope? 
• y • y 7 . • • •. -.•-• :.••• • •• • • ••• -•• ' y..-A ;,•••• y : • v 
. y 
• • . • • . . . • • . . : . ; . . . . . . . . . 
• • • • • • : . • • . • • . . : • • • • • • • : • • . -. . • • . • • • ••.. : • • -• 
• ..y.-. . :•' y y y 
7 ..... 7 .••: ,...•. . .7.7 • ...... y ••/ ...yy y .•-... y, . •. 
.•••' ' ...-•/ . ' y •'•'.; ...... . . • • • .. . 
• ...... . 
•• . 7 '"': : ' 7 • . ' • • • ' • . • • -. . • • • • ': •••• '." ' • . -• -. • • • . 7 • ' • • • '..' •• • 
• . -• • • . • • • • • • 
• • , . . : . , -. . . . . , . • • • . ... . . 
. • • 
• • . -. • • • . • • • . --• • • . . . • • . . • . . • . / 
-• • • • • • : • . •• • . • 
: • -• • • . . L . " . .' . • ' " . . . • ' ' • •• • • • " 
. '•' 
4. What is the purpose of a calibration graph? 

••• •: 

" y "-...' . ''• y ••• -'• -7-. ":• • • • : •'..'"•. . 

•,' •••• • • . • . -. • • • • • " : • . • • • : •-. \ ; . • 

:•-••••' • 

. •• • . •• ••; • •?-• •;-.. .- -.• -. •• . ...• . A-. .-.:: • •..• 

-• ' • . . • ' • • ' • • . • . . • • • . . . . . • • . • • • - . : • •.• , . ; . 

" . . . • • 

5. The Uv emission from what element causes the fluorescent 
powders in a fluorescent tube to glow? 
• • •. ' • 

..'-'-• • • • • 

y ' 7•••••• • •.. .7 • '•" •••'" ' y • ' 
"• 7 •''•'.' . : '•• : ••.".'.' • • y .-'•'. • '• 

7-' , .'y•••. 
• ... : • •' 
y 
•y. ; 

'• :•":• y .••"• ' .7 '•:-;' '••• "'• ••-' 7 - 
• ' • 

-. : • • ' • ' . ' ' • • . . . ' '.'•• 

7 
. ••: • • • . • ' • • • • •" • ' • 
y • : • 
•• .- .. •.• •.. 

7 " : ..• 
y • • ;' 

,. • • • • . 

. 
• 
' ... 

• 

'••'•'• • y • ..-. • •

. •. . .. •'-• ; ' 

Emission 12 

• • • • " . • • . -• • • . • . . . . . 

. A ••'•••••. --y ..-: •• • . •,- .-. --.-• .-•• •.•• : y • • ••: . •• . •• • 


i: 7.. ÛALORIMETRY 
SEAT OF FORMATION OF HAGNESIUH OXIDE 
chemistry 131 

•*'y INTRODUCTION 

;,7- In this laboratory, we will introduce one of the most of ten-used 
techniques in thermochemistry - calorimetry. Although we often think 
of calorimetry in terms of finding the number of calories in a: certain • 
amount of food, calorimetry is valuable to the chemist in measuring 
basic thermodynamic data. Along with learning calorimetry techniques, 
you will use the data you collect, along with heats of formation and 
Hess's law, to determine ûHt" for HgO. 


THEORY 


AS a quick rev lew*; iërteirtbei' that th* enthalpy change, AH, of a 
chemical reaction is called the heat of reaction and representsthe y 
amount of heat gained or lost as the reaction proceeds from reactants 
to products. Also, the heat of formation, ¿fïe°, is defined as the 
amount of heat absorbed when one mole of a compound is formed from its 
elements in their standard states. Another important concept is that 
the enthalpy change of a reaction is independent of its path and 
dépends only oh; the initial and final states of the reactants and; 7 
products. This principle is known as Hess's Law and is an example of a 
state function. Hess's law states that the enthalpy change of a 
reaction is the same whether it occurs in one step or in many steps, ''.y 

You will apply Hess's Law in determining AHr° for the following 
yy--. reaction: 

Mg,sj + 1/2 o 2(r). ^ Mgo(i).
• . 
(l) 


7AHf ? ''•=• ?'; 


This reaction is extremely exothermic and too hard to perform 
calorimetrically. However, we can find the heat of formation for MgO 
by combining a series of reactions which are much safer and suitable 
for a calorimetry experiment. Here is one possible scheme (perhaps you 
can think of others): -. 


•Kg.(.¿> + 2H+ ,vy, _*. .Hg.**',,-, +:H2:((} (2) 
;yAft;^ ?;y 


(,} t 2H^f M jy-> Mg;**!^) + 7Hi.0'(£i (3) 


7yy.A¿--:?:'..7 


ENTHALPY 1 



Hï(fc) + 1/2 0 Î( 0 _&. HjO(í) (4) 
AHf° - -2B5.B kJ/mole 
Using Hess' Law, reactions (2), (3), and (4) can be combined to 
give reaction (1). Work this out for yourself in the boy below: 

Thus, if the AH for reactions (2), (3), and {4) are available, you 
can find ùH for reaction (1). You'll use a calorimeter to measure the 
enthalpy changes for reactions (2> and <3). AH for reaction (4) is 
given so you can find AH for reaction (1)» the formation of MgO, 

Let's examine calorimetry in a bit more detail. First, almost any 
type of container can be used as a calorimeter, but a well insulated 
container is best. Why? Because we must account for all heat that is 
absorbed or evolved by the chemical process. Therefore, we want a 
container from which heat cannot escape to thé surroundings. A 
chemical process occurring in such a system is adiabattcj that is, no 
heat is exchanged between the container and the surroundings. You will 
use a thermos bottle to achieve adiabatic conditions. (You may ask 
yourself at this point if the thermos bottle will provide a truly 
adiabatic environment.) 

From here on out we will be talking about adiabatic processes. 
The important thing to remember when conducting thermochemistry 
experiments is that you must account for all heat gained or lost during 
a reaction. For example, if an exothermic reaction occurs in a 
calorimeter, the heat can basically go two places: (1) the reaction 
mixture, which can be measured as a temperature rise, and (2) the walls 
of the calorimeter. Since heat is not transferred between the 

ENTHALPY 2 



• 


calorimeter and the outside, the following must be true: 


heat from — heat absorbed by . heat absorbed by 
reaction calorimeter react1oh mixture 


or 
-<3I-*ÍV *• qc.ii :+. qü „ (where q - ^resents heat) (5) 
Thus the heat lost by the chemical reaction is exactly equal to 
all the heat gained, why is there a negative sign for qrxn? 


Now let's examine each of the terms in the equation:: 


Seat from Reaction (qFin) ; If the heat (q) comes from the 
enthalpy of a chemical reaction, we can replace q by nAH, where n is 
the number of moles of product and AH is the enthalpy change per mole 
due to the reaction or phase change* The number of moles of product 
will be based on the moles of the limiting reagent. 

Heat Absorbed by the Calorimeter (gfltl): For our calorimeter, the 
heat simply changes the temperature of the calorimeter* We can replace 
q by CCAT , where Ce is the heat capacity and AT is the temperature 
change. Note that the heat capacity is the amount of heat required to 
raise the temperature of the substance by one Kelvin. The units of Cc 
are usually J/K. Thus we need the heat capacity of the calorimeter, 
Cat which is specific for each calorimeter. You will find Cç for your 
calorimeter by performing reaction (7) below for which we have provided 
AHr3tn. You will then use equation (6} on the next page to solve for 
Cç:*y 

Beat Absorbed by Reaction Mixture (qn<,): This term is equal to 
mCpAT. The mass of the solution is m and is equal to the total volume 
of the solution times the density of solution. We simplify by using 
density equal to 1.00 g/mL since all the solutions are essentially 
aqueous. Although for pure water the heat capacity is 4.184 J/(g "C), 
when ions or molecules are dissolved in it the heat capacity changes. 
For the NaOH and HCl reaction the heat capacity of the solution, cr, is 
equal to 4.025 J/(g *c) while for the Mg reactions, CP equals 3.862 
J/<g °c>. 


ENTHALPY 3 



¡V.'í :'; 

:".-Z '.! '•' 

54 


.7;'.Usi.hgf-these expression s fo r q, equatio n (5) i s rewritte n as : ; 

-JlAH .=• CcAT + mCpAT V, (6 ) 

.ywheres-A y.'-' 

yyy y n = number of moles of limiting reagent. 
'yyy77 ^H = heat evolved in the reaction. ->••• 
7 Cç7:» .calorimeter constant (specific for your calorimeter);; 
A T = temperature change resulting from the reaction. 
y7- m = mass of the solution. 
cp - specific heat of the solution. 


-Note that we used q •= mC^AT for the liquid in the calorimeter 
because we know its exact composition. For the calorimeter itself, 
Thowever* a heat capacity, Ctr was used because the exact composition of 
the calorimeter is not known. 7 

Now let's apply the theory of calorimetry to your experiment. To 
find AH for a reaction, all the other values in Equation (6) must be 
known. The first objective will be to calibrate the calorimeter to 
find Cc * This calibration is accomplished by producing a known 
quantity of heat from reaction (7) below and measuring AT. Since AH, 
th£ heat capacity, and the quantities of NaOH and HCl are known 
(assuming that all volumes ara additive), the only unknown in equation7 
(6)7Ï& Cc which you can now calculate. 


, H*Ok<M ) ;.SH HClj-q) — > ^Mát^q)+._«-"(•••*•)•.+ 'H.aP(-.tï =• i7> 


"; ifirVar^-HB^V?:' kJ/raole;7 


•J---. t • i 
7-7once the value of cc is known, it can be used for reactions (2) 
and (3) to solve for the unknown A Hf. 


:To :^#e eguat ion: •f 6); you^ if!i ï: need AT, the temp e rature change for7 
the reaction mixture and thé':;cal.orÍmatér¿;-. However, you will not be 
able to simply measure the maximum temperature because the temperature 
may never actually reach the maximum. This irregular behavior is due y 
to thé calorimeter's inability to absorb heat as quickly as the 
reaction mixtu^y You will use/.a standard graphing technique to' 
extrapolate to the maximum temperature for each of your experiments. 


During this experiment you'll measure and record the temperature 
"at vg iyén t ime intervals. YOU will plot; this::: data;. on7a graph 1 ike that 
shown in Figure 1. Notice the temperature vs. time plot is irregular 
until it stabilizes on a slowly decreasing temperature line. Ideally, 
you want the maximum temperature that would be produced if the reaction 
happened instantaneously at the time of mixing (timé - o), not some 
time later when the reaction mixture and the calorimeter tend to cool 
slightly. The proper method to obtain AT is depicted in Figure i. The 
rate of cooling is used to extrapolate back to what the maximum 


: yy yy.7-';77y'.7yyyyyy-Ayy-y.- .--yy.••yy-yry.-'-'y'' ••'-777'•-'•7"'•'-•• -y"-> yy..:.-yyyy 

;y^,'y.y- •'•' . . * '" . ENTHALP Y A., y ' .""' -. y-



temperature should have been. This is the intersection of the 
extrapolation line with the y-axis (time = O) . The difference between 
the extrapolated maximum, T», and the initial temperature, Ti, is AT. 


EXPERIMENTAL: 


This experiment will be done with a PARTNER. 


A. Prelab calculationst 
The solutions of HCl and NaOH that you'll use in your reactions 
must be prepared from our stock solutions. You will make 200 mL of a 
1,00 M HCl solution from a 6.20 M stock solution. You'll also make 100 
mL of a 1.00 M NaOH solution from a 3.ID M stock solution. Prior to 
lab, calculate the volume of each stock solution and H3o that you'll 
mix together to make your solutions. 


To make 100 mL of 1.00 M HCl, start with __L.mL 
of 6,20 M HCl and dilute it in a volumetric flask to a 
final volume of ; mL, 

. . .. 

You wil l need,to repeat thi s twice since you wil l need a tota l of 200 
mL 1.00 H HCl. 

To make 100 pL of 1.00 K NaOH, start with raL 

of 3.10 M NaOH and dilut e i t In a volumetric flask to a 

final volume of uT


1. Prepare stock solutions» (One partner should do this while 
the other partner is weighing the solid chemicals.) Use two clean 100 
mL beakers to get stock solutions from the side shelf and mix them with 
the calculated amounts of water in 250 mL beakers. Record accurately 
the exact molarity of HCl and NaOH as labeled on the containers and the 
exact volumes of HCl, NaOH* and water used. You'll need this data to 
calculate exact concentrations of your solutions. 
2. Weigh solid chemicals: (The other partner should do this 
While the first partner is preparing the solutions*) The Mettler 
balances are delicate instruments. Do not spill any chemicals on the 
balance. If you do spill chemicals, clean them up immediately. To 
weigh a chemical, place a clean, small piece of paper on the top of the 
balance* Press the TAKE button to BERO the balance. Add the solid 
slowly onto the paper until the proper mass has been weighed* 
ENTHALPY 5 



• 


• • . 


a. Neigh 0*20 to 0.30 g of Mg turnings to the nearest 
milligram on the Hettler balance. Record the exact mass of Mg used. 
b. separately weigh 0.50 to 0,60 g MgO weighed to the nearest 
milligram. Again, record the exact mass. 
. . . .'. • " . • • ' . • •' • •

• 

.• . : •

B. ;': Calorimeter calibratlon (Finding Ce ). 
• ••'.• •. 

WASHING: The precision theraometers and thermos bottles 

ara: «Kpéhsivè. , Treat thee vith care. If you break a 

thermometer, tel l the lab instructor LamedLately so the 

nercury spill can be cleaned up. Also, HCl stains banch 

tops; clean up spills immediately! •
••• . :' •;•",.; ."•.'•• • •.-:-7 ' 

1. Record your calorimeter number and insure the calorimeter and 
thermometer are clean and have been rinsed with de-ionised water. 
• • . : . 


2. Using a 50 mL graduated cylinder* measure 50 mL of your 1 M 
HCl solution. Pour the HCl into the calorimeter. Insert the 
thermometer assembly and record the temperature when it stabilizes 
(This should take less than 3 minutes). Measure all temperatures to 
the nearest 0.05&C with the aid of a magnifying glass. 


3. Wow rinse the graduated cylinder successively with tap water, 
de-ionized water, and 5 ^ of 1 « HaOH solution. Measure 50 mL of your 
1 M NaOH solution in the graduated cylinder and pour into a small, 
clean, dry beaker* Record the stabilised temperature of the HaOH 
solution. The average between this temperature and the HCl temperature 
in the calorimeter previously recorded is to be used as your initial 
temperature (T¡)EMTHALPY 
6 



yy i? 


4. Noting the time, pour the WaOH into the calorimeter. 
Immediately insert the thermometer assembly and gently swirl the 
reactants. Record the temperatures to the nearest 0,050C at half-
minute intervals until a maximum is reached; then record at one minute 
intervals until enough data is obtained for an extrapolation (5 to 10 
minutes). Gently swirl the reaction mixture between each reading. 
Tim. Tempera tur* C Tim» Ifiipfr itur t 
L '.. ' 
0 3*0 

-. ' . • • : . ' .'-'. • " . •• .7 A ' 7. 

30 3fl0 

' . -: • 

60 :•";•.,' . '.'.' . .-:• : ' y 390 

;'.'.' .' •••-' 

:

... • . 

• • . . •• •

ÍO 430 

12 0 4SO 

IS O 41 0 

HG 31 0 

'••'• • ; • '. 

a*o S40 

240 570 

±70 Í0 0 

300 

• 

ENTHALPY 7 



ïy'yy:.. y;-yy, y ';'y ' A--y; 'y '¡"'y'yyyy*. •' v':V.'.,: ;•'';,.'., ' 7:777'.'..•" 58-> 

-7.177,77. -. -: -'A''": •'.•' •• -A-A.y, ; .,7 :7 : . ..-.;• •. y.y, ;.-•._•:•_ y y-. 7 

. ' . :. .': . . • • ••'.•;• • -y -• -v; : . • • .'..'.. . '••••"-' . ' • .'.'. 


c. HeatofReactionof«g andHCII 
1. Clean and dry your calorimeter and add to it 50 mL ol your 1M 
HCl solution and 50 mL of de-ionized water. Record the stabilized 
temperature (T^. 
"7-'7 '•• '. yy.'-'y-A. ....7-.• •• 

2. Noting the time (t « 0)» drop theMg into the calorimeter, 
immediately putthe thermometer assembly in place and swirl gently. 
Record temperatures at haIf-minute intervals until a maximum is 
reached; then record at oneminute intervals for at least five tninutes. 
Gently swirl the reaction mixture between each reading. 
..y 7. "••'• • "••'••'•'."'. : :'•:'•' y-' • y • • • ' : • y' ' .-. ' . ' . . . • ' : ' .. . A; 
• . • 
Tim* Temper*_ty_r_g_ 
• 
• • •: • : --; • -• -. . . . . • • • . • . 
• . • 
• : . ' • 
.;-.*! •a : ' •.-• • -
A -; 
.." 
• • . • 
• . 
• • • . ••'•' •' : . : . . . .
..
ao 
. . y .y . . • y . y . : • 7 • ••:• 

• y ••'••'. • yy . . •••:• • •.••• • 
••• ' •••'. , •• ••• 

• .. ; 
•. " •" :•'.'• . -•• 
9Ù 
• 
• • y

..." .-• y... . 

1= 0 

• 

. • 

. •• • 

'••-"•••" yyy^ i• ••;.: ::':

•':— 
. 150

A . • •"•.:••'••"• • 7. 7 .-. " • : : ' • • . ••' -• . . • ' • • • :• • 
.. "• A' ..y •• y y ..-•••• ... •.••••• • A . y . . y y y 

1.0 
.• 7 •. 

7,7 •.' y . .-. .-y., y -...-.:•:.' y 

• 


Ï10 .. . 

• • . • . . • . . . . . •

• 

. 
. 
• ' • 

. • . . , . . -• . . . . . • . '.'•• • . • " • . . . . ' • • • 

••'•••..•y ' "•• " • -• • .•••; • -• . . y .

... . • 

. _ • • 

. 240 

• . . . . • • . -. .; . . • • • . . 
• 
• 

. . • . • . • : , . . . . • • — • . • • 

;

• • ' . • • • •..•....-:" . -:'• • :•'•' • • • ' 

'•'.' .' •:.'.•.:• . : '-y--. . . • • • • . . • . :: 

• • •>

• .... -'• •
. •• • -y 

-• . • . 

•

• .... • 

.-• 

300• 

• • • ' • • • . • '. . y • . . • • 

3** 

y y . • '•*•' •: : "•'•''• •' :' .7 7 yy'-' . -: 

• • .. 
• •• • • • ' 

y*** yyyyy .


. . ' . . • • • • • . • • • • • • . • . .. . ••. . . : • : • • : • • 

• ••• . • • : • • ' : ' . • ' . •' ' -. . " •"' . '•-. • : • ' • • • . ' . • . : 

390 

. ' . • • • • • ' 
• • • • • : : • : . 

•.•••-' '.:.•'••
420 

..•••" • 

• y , • .. . : . . • • • ; ••••• • • -. . 
• • • • • • ' • -. • -. . . . . . . 
.. : ••. . • •. • . . . y . • . • . .. .• 

•

45 0 

• • • . •. • • • 

••••••.-A.-777-•. y y ••. .-• • . ,• -. • 77; 

.-.' •: -' y y '•-.:'""" !'••••'••" ' y • ' • 


• '•• • '.• • . ' • ' • ' : • • • -' : ' • • • " • . " • • • '••'•' -'. : • ' • , 
'••'.-. :.' -. •• • • •' < -; ' ' A • y :-••. •• •.. 

ENTHALPY 8 

• 

• 

• 

.'-V •i; :.: "'-''V-• '• .• 7 •..,•: '••'•" .': . '•--•• 


• ': 
D. Heat of Reaction of HgO and HCl; 
1. Using the same calorimeter, repeat the procedure used in Part C 
using HgO in place of Hg. (Note: You must swirl the calorimeter 
vigorously while reacting HgO and HCl since the MgO tends to form lumps 
at thi* bottom of the calorimeter and fails to dissolve. This could 
cause considerable error.) Hold the thermometer assembly in the 
calorimeter so it will not rattle while swirling. Record temperatures 
at half-minute intervals until a maximum is reached; then record at one 
minute intervals for five minutes. Gently swirl the reaction mixture 
between each reading, when it appears that temperature changes have 
ceased, quickly look into the calorimeter. If a white solid remains at 
the bottom, the calorimeter must be swirled more vigorously until all 
of the HgO has dissolved. 
-• 


• • " ' . • • ' • " • ' " ' " -y -' • • ' . ' -' ' • > • . ' • 

. • 


• 

• 

. 

• 

: •

' • ••

'

' •

. " . 

y <_ 

• 


' 


. 


• ' • ' : . -. ' • • • -• • -' • ' . . 
7 • " '•-'.•', '•' '• y y' 

. 


•A . 
•


ENTHALPY 9 


• : • "' 
.• • 


: '• • '• 



Timí


30 


60 


• • . • . 


SO 
i 20 
ISO 


i ta 

240 
3?0 
3 DO 
330 


n*o 

35Q 
430 
4SO 


 Tempérât trt 


• 


' 


• . 
-


. 


• •: 


EHTHALPY 10 


• 



.-•' y 
M 
Tine v*. Tempi ra ( u re EntrapnUt ion 
30Te-4. 
_ „ .• 
J ime.O. 
Jeinge rgjjare 
2 3.5 
•• 
Temperatur+^S 
-
-r-
no 
'.*o .
SO12015 0 
iso3102402703(JO_ ^ 
ïjo( 
3a/iff 
28 . 12 
28.BS 
.
..
: a8 i B4 
aa, so 
aa.?i 
28.7? 
23.6* 
2fl . 64 
^ T 
120 
{teconde ) 
leo 
• 
240 • 
' 

:• -: • . • . '•' 
• • ' • -. .-y -• 


:-.


• : • • • ••..•. -• • • 
'•'••• • •-' •• -^ " . . • • • " 7 : • ..• -• y . • • • .;•. ••'.:•• 7 y 

• y '••.. • 7 -.7. -'• : -; " y 7 . ' • / • • .-. ' : y 7 .• .... '•'.; 
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FIGURE 1 

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ENTHALPY 11 
" •• • : •'. 
• • • • . :• • 
•'• y . 


BBBQIitfi 

A. calculations: 
1. Record all volumes and weights of chemicals used and show 
calculations for all values necessary to report the required values. 
Be sure to include: 
•».; • Dilution of all stock solutions. 
b. For each reaction: 
Holes of reactants and products. 
Limiting "Reactant" calculations 
calorimeter constant. 
d. aH for reactions 2 and 3 and ¿Hf" of HgO. 
B. Orapbsi 

1. In this experiment, a good graph is critical to finding the 
correct temperature changes for each reaction and is an integral part 
of your calculations. Without it, your calculations are of no value. 
Refer to ybur "Data Analysis" handout if you've forgotten the proper 
way to draw a graph. Label Ti and T« on the temperature axis and show 
your extrapolation to find AT. You will probably need to draw each . 
graph on a separate sheet of graph paper. 
a. After plotting your data points of temperature vs. time, 
draw a sraooth curve through these points* Your initial point at time = 
o will have the initial temperature (Tj) of the solution prior to 
mixing (refer to Figure 1) . Draw a straight line from your 
"stabilized** data (data recorded over the one minute intervals) back to 
the Temperature axis. This is your extrapolation. The extrapolated7 
temperature, T*. is 'the temperature where this extrapolation line 
intersects the y-axis. 
b. Use T» from the extrapolation line* and Tt to calculate AT. 
ENTHALPY 12 



êï 

COMCLUSlOMBl 


A. Why was it necessary to use adiabatic conditions? 
B. Is the formation of HgO exothermic or andothermic? 
C. How does your value of AH*0 compare with the actual literature 
value? Calculate the percent difference between your value and the 
actual literature value. This is: 
-77-y. - (exper intentai value) - (actual value) 


* Difference — —-1 — X 100 
(actual value)/ 
Comment on possible sources of error. 


ENTHALPY 13 



h h 

DATA TROH AHALYBUI GT PRESENT 
ABD rUTDRB ENERGY BESOORCBB 

>>y yy 

: 


Hajior Steven E. Dunlap 

Department of Chemistry 
United States Air Force Academy 


. • 

• • • 


.. .•; • :


CURRENT MAJqR.BHERqY BOURCEB 


oi l 
43% 

Coal 23* .
..
• • 
... 
Hatiiral: Gas 22% : 
Nuclear 6% 

. : • 


. 


OIL MSODRCBB 


WORLD 
600 


billion barrels of known reserves


525 


billion barrels of estiaated undicovered recoverable


21.7 
billion barrels per year current consumption 
years remain of known reserves at current 
consumption 


GMXTZB BTATE9 


24.6 
billion barrels of known reserves 
(4% of world's, had 15% initially) 
5.9 billion barrels per year current consumption 
3.. 16 billion barrels per year current production 
7*8. years remain of known reserves at current production 
OWITED STATES IIIPOilTS 


2.74 billion barrels per year (46.4%) 
19% canada 12.6 years remain at current production 
10% a.k.; 9.7 years remain at current production 
io% Hex!co 27.8 years remain at current production 
20% Africa 29.8 years remain at current production 


7* Ven^zuie 39.8 years remain at current production 
13% Middle East . : 


• 



•-.:••:•• y y *>$: 

çoA^jagogRcea 

•
•• 'iwtoi-.'.'9JQ7 
7 billion tons of known reserves 
6-12 trillion tons of estimated undiscovered recoverable 
DOTTED STATES 
283 billion7tbns of known reserves (29% of world's) 


0.818 
billion tons per year current consumption 
346 
ysars remain óf known reserves at current 
consumption 


HATORAL QAS RESOURCES 


WORLD 


3*6 7 quadrillion SCF known reserves 
6-12 quadrillion SCF estimated undiscovered recoverable 
59 trillion SCF per year current consumption 
61 years remain of known reserves at current 


consumption 


UNITED STATES 


185 trillion SCF known reserves (5t of world's) 
¡18*2 trillion SCF withdrawals per year 


10.2 Ay#ars remain of known reserves/átycurrent consump. 
U.S. IMPORTS 
•jai7'5.:.'.''..trillion*. SCF; par.year.":'• 
7sr:*; 
0¿06 trillio n SCF per year 

yy. WORLD ' yy 

4.3 : .; milliph:. tons under current mining constraints 
26 million tons technically obtainable 
tWiTID STATES 


12rl^/7 years under current mining practices and planned 
coasumption 


50^60 
years using all technically obtainable Uranium 



hb-

ESTIMATION OF OIL RESERVES 


SEDIMENTARY BASINS 


600 geologically known 
160 productive 
240 non productive 
100 unexplored because of hostile conditions 
loo will probably never be explored because of 


conditions 


BREAKDOWN OF RESERVES (BILLIONS OF BARRELS) 


5)2.5 Cumulative production to date 


511.5 Proved reserves 
101.2 Inferred reserves 
513.97 Assessed undiscovered reserves 
294.0 Estimated undiscovered reserves 

fr7 

OIL RESERVES TERMINOLOGY 



•My yyy 
EFFECT OF O.S. CONSERVATION 


TRANSPORTATION 


0.5-1.2 billion barrels potentially saveable annually 
3vie bill ion barrels per1 year70urrent .çonsumpt ion 
1.5-4.5 year extension 


RESIDENTIAL DBAOE 


ENERGY USED FOR HEATING 


:.3¿ «'•"• 


trillion SCF natural gas per ysar 


0.19 
billion barrels of oil per year 
0.30 
quadr i 11 i on BTU o f e lee tr ici ty >,_•• 14',•:'powfttv1 . plant*. 
' i ENERGY USED FOR AIRGONDITIOHXNG 
0* 36 quadr i 11 i on BTU .of else tr ic i tyt IS power pi ants 
EMERGÏ USED FOR APPLIANCES 
1V53 quadrillion BTU of electricity,, 74 power plants 


COMMERCIAL USAGE 


2,2 
trillion 5CF natural gas per year 
billion barrels of oil per year 


0VQ6 


quadrillion BTU of electricity* 109 power plants 


POTENTIAL SAVINGS 
100 Power plants 


: • 7r • • . ; • 

i • -, . •• ••\ 

• •.'.• '''-.$.. 


• • • • • • ' . . . • • . • • . -y . 


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; C.\"v .-.::V 7: ' "•••• C'-;: ' " .• • • : /y -'' ••',;• -i-.; : -- '•!';;-• 

EFFECT Of QRQWTH gN WORLD OIL RESERVES 


PERCENT PATE • YEARS 

OF INCREASE ...... .' REMAINING 

-•:. 


• •'.•• 


• •• . . y • 

0.00 27 
: • •

1-43 Projected by EIA 23 

•

2,66 Current U.S. 21 

7.oo u.s. prior to 1970 .9 

-•. • : 

• . •

y y =• 

: • • • ' • . : ' . . • '•;•:: . ' • • 


. A 

• . • . • • • • . . -, • 


-

• • •' • . -. 

• • •'• •

PgQVIBBHBtiTg 


• v " • ' • . ' • • • • • . '• ' -• -. • 

Total replacement of present major energy sources. 


••'• " •-' -.... y • .':•' .' • •• • .y ... y 
• 


. • Solve the problem of the heterogeneity of current energy 


• . 
forms. 


: ...•••• 


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7 ••'•• --.'•• • . A 7, -7 : .•• '. •• •'• . ..; 7/ • • y y_. 

MQCLEAR FUEL FROM RgJEQCESJJJM OR 

.


BREEDER REACTORS 

CAPABILITIES 

: •'


;'-.•• ••.'• • •••-' - •"• -y .' -7 :' .-''. ' • ••. 7 . 

0.6 : percent of uranium ore's potential energy is 
harnessed by present once-through reactors 
.-.•'••' •": 


. . ' " . •• ' . • • ' -. • . • ' . ' . . . ' . . ' • : : . . ' . 


• • . • ; • • 


• 


70 percent of ore's potential energy obtainable 
using reprocessing or breeders 


looos of years fuel possible at current rate of use 
and with present mining practices 


.. '. • • • . . ' • ' ... . • • • . . -• • • , -. •. • 

100,000s years if lower grade ur n is considered 


'.'•••.'••':. 7 .7- '"'•• • . -, . : y • . •• : • • . ' . ; -, ... 

• 


BUftHT DfHKVfS 

• • • .• y • '• • -•••• . : •• . 

300 additional 1000 MW reactors to replace fossil 
fuel generation 


. •' ''•'• • ' •'.-. • 7 ' • • • • ' ' -. ' . A . .•• • '•.• ' . --' • •'•. 


1ÓÓ0 ; more 1000 HW reactors to meet total U.S. 

• energy demand 
• . ' . ' • ' • ' • • . . ' •'...'..••.'.•:. • •.' A • : 
7 • • : '•.• • • • • •• •.. • ' . -. . . " -" '• 

. -; • • . • • • ; . . . . -• • • -. . • • • • • •

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:•' • ••••.-.. "••••• • • • • •'•:•• ' • • * • . . • 
BESTCOPyAVAILABLE 


A . 
STATUS 


France has a 1300MW breeder on-line. 


England should have a large commercial breeder go on-line 
this year. • 


Soviets have one breeder under construction and another 


planned. 


Britain has reprocessed from the start. 


France has the world's largest reprocsssing center at 
LaHague. 


U.S. had an experimental breeder in the 1950s. 
U.S. has a 350 MW breeder reactor on order for Clinch 
River* Tn. (construction subject to resolution of national 
policy debate) 
u.s. presently has: 
100 operable once-through fission reactors 
.
..-.1 ;'in startup . • '•\ 
19 construction permits granted 
0 on order 


• 


• . • 


FHOTPTOLTAICS 

y'' CAWlBlLlTIlfl/RBOUIftEMEBTB 

Landmass of Maryland and Connecticut to meet current U.S. 
electricity demand. 


Landmass of Nevada or Utah required to meet total U.S, 
energy demand. 


STATUS 


0.1 MW facility operated by Alabama Power company 
660,000 MW current U.S. generating capacity 
$10-30 cost of cells per watt would increase electricity 
bills 100 percent. 
$*1~.30 cost per watt predicted in 1990s by DOE 
10 year current call life 
20-30 year cell life necessary to make competitive 


BEST COPT AVAtLABLE 


7;77yBj[OMABS. 
;;;NOpb.;: 
;,:.'7"-.77:7. 
Has eclipsed nuclear power in the United States 
75.
.. 
y,•.,:•
S2-: 
percent of present forest required to meet current 
. .U.&' electricity demand 
percent of U.S. landmass required to meet total U.S. 
energy demand 
10 : years required to properly establish a wood fuel 
plantation 
FUEL CROPS 
177 percent of Brazil's autoaotive fuel mads by 
fermenting sugar cane 
sugar cane doesn't grow appreciably in the U.S. 
Î1corn gives half the yield of sugar cane 
percent of U.S. landmass required to meet automotive 
fuel requirement 
206 
percent of U.S. landmass currently cultivated 
percent of landmass required by some "hopeful1* 
•future' -plants 
1200 percent of U.S. landmass required using Calvin's 
Euphorbia Lathyris 

7.-' .7iJtataiia"7y. 

FAST AND PRESENT PRODUCTION 

OIL FROM SHALS 


4400 barrels per day (max)r Scotland, 1859-1962 
3000 barrels per day (max), Australia, 1862-1952 


50 commercial plants. United States, 18*9 
50,000 barrels per day, China, currently 
25,000 barrels per day, Soviets, currently 


5,000 barrels per day, U.S., currently 


OIL FROM COAL 


7 so barrelB per day# Scotland, 1850 
100,000 barrels per day, Germany, WWII, 1/3 Wartime 
requirement 
5,ooo barrels per dayV SASOL I, 1960-present 
58,000 barrels per day, SASOL II, in startup 
300 bat-rels; :per day,.Union Cát-bide; ;19SS-^1962 

••: # . ' 

BEST COPY AVAILABLE 


•> .. . . Í .-:• 

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SHALE 


80-600 billion barrels of recoverable shale oil in high 
grade deposits of Colorado, Utah, and Wyoming 
14-102 years At current U.S. consumption 


COAL • 


a.818 billion tons p*r yea r current consumpt i on 
10 percent rate of increase in usage proposed by Ford 
2*19 billion tens per year in IB years 
meets 3.28 billion barrels of oil import requirement 


• ':••*•'=• tha t •time; • 


3.65 
billion tons per year in 29 years 
meets 5.9 billion barrels of oil total requirement 
i n: • 2.9/ y.eSr s 
42-577 
years remain at this rate:0f use 


•7:y: REQUIREMENT* 
165 synfuel plants with a capacity of 58,000 barrels per 
yyyy- yy day by 200O yy-. 
$2-4; billion per plant and 5 years to build each plant 
Immediately begin building pioneer commercial scale 
plants in this country. 
STATUS AND PREDICTIONS 
Will not meet the government goal of 0.26 billion barrels 
per year (4.5%) by 2O0O under the current administration. 
$83 billion programmed, only S13 billion ever appropriated-
With an all out effort could reach 2-3 billion barrels per 


year by 2000 and sustain it for IOO-150 years. 


QËST COPY AVAILABLE 


HYDROGEN TOR FUEL 
1979 Roger Billings marketed hydrogen powered Dodge Omnis 
724 Commercial airline aircraft assigns (twice as efficient, 
33 percent increase in payload) 
1956 USAF flew a B-57 with one engine operating on 
hydrogen. 
1500 tons of nickel to meet 1965 gap in supply and demand 
for methane 
16,000 tons of nickel produced annually by U.S. 
To meet U.S. natural gas demand, a severe strain 
would be placed on domestic nickel production. 


W«MMEjroaTIOMg 
Synfuels should meet import requirements by 2000. 


• 


Huclear fission with reprocessing or breeders must carry 
electrical load by 2O40 and entire energy demand by 2100. 


Convert to hydrogen economy between 2O40 and 2100 as fossils 
are depleted. 


Renewable^, photovoltalcs, and nuclear fussion can 
substitute for some of the above requirements but will not 
alter the major periods of transition. 


7"7 ..;...-..••...-• 


BEST COPY AVAILABLE 


• 

.-•.-. 

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Recommendati is* Time Une 

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•••: • '• .' • 

• . •• . •. . 
• '• • • 
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1W 
1W 
•' • 

•

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2000 2040 2100 

•.'• • 
. •

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T 

... 

.:• . . 

165 Srnfwl Plant* •• • 

• •
•• • y •. 
• • . 

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. •• 

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.• • • • 

1330-460 büüoo • . 

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• 
• 


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•

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•.•.•••••• •••••••• y . •• •• yy.: -
Guuiei lotionlo 
reduce OUrf BOT 


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plsntt required 

-••' 

• 
• -•. 

from. 400 to 185 . • 

• • • • : -• • . • -".. . •' . 

:•• . • 

Miiiitititmnittiiit/ifiHtuniitiUiiêntitiiitiK 

attention dependent on tha TlmWiltr of xaete oil 

Coal used tor rynfuel* and electricity 

. 

1

• ' y . : • • .. . 
: Conrtruction of 300 breeder reacton* $222 billion 
• 5.7 plantt/yeor «t $6$ billion/rear • 


1 

. • • ' • A 

. 

• y • • : 
Complete contortion io nuclear end *hrfo>Rje& sconomr 

. • • 

• 

Add 1000 aere breeder reertort 
Build hydrogen production raeUitiat 

.•• . •• 

mmrtitth*»mmimttil»i>»fci»Hinn.ra^ 

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^: 



MOLECULAR OTÓMETRY AND 
ELECTRONIC STRUCTURE LABORATORY 
Chemistry 131 


BACKGROUND 


Chemists have techniques by which molecular structures can be 
determined. Among them are X-ray and neutron diffraction for 
solids, and microwave rotation spectroscopy and electron diffraction 
for gases. Other methods used in structural analyses are infrared 
and raman spectroscopy, nuclear magnetic resonance, and mass 
spectrometry. Here at the Air Force Academy, the Department of 
Chemistry and the Prank J. Seller Research Lab have most of the 
spectroscopic instrumentation necessary to do these analyses. 


EIPERIMRMTAL 


Procedure 


Although you'll not experimentally determine the 3-D structure 
of any molécule in this course, you'll become familiar with the many 
geometric arrangements and molecular structures found in simple 
molecules, one of the easiest ways to do this is to use models. 
The kit you purchased at book issue allows you to build a model of 
each of the geometries found in your text book. Use the various 
colored balls to represent the different types of atoms in a 
molecule. There are balls available with two (red), four (blue and 
black), five (brown)* and six holes (gray) to represent the central 
atoms for all geometries. Try them out. Discover the beauty of 
three-dimensional geometries. 

Attached to this introduction is a prelab worksheet. Fill in 
each of the columns on it before y_çp cogiç to the__lab._ Be sure to 
name each molecule. Also attached is a checklist for doing Lewis 
structures. 


Equipment 


Bring: 


1) The completed prelab worksheet. 


2) Your model kit* 


3) Your textbook, 


to the lab with you. 



. 

A CHECKLIST FO» ORA*IWG LEWIS STRUCTURES 


1. extermine th* central atom. The following are guides; 
a. The least electronegative ¡element is the central atom. 
b. Cftsn thés unique atom (only r>ne of it) is the central atom. 
c. Sometimes the formula is written with the central atom in 
the middle. 
2. Arrange the other atoms around the central atom creating a 
skeleton. 
a. Oxygen rarely bonds to itself except in: 
1) O2 and o3 (ozone). 
2) Peroxides, ePg-F HfO=. 

3) Superoxides, e.g-, rcaOj, 

b . Fluor line «¿ver bonds to snore than OÏÎS atom„ 
3. connect all bonded atoms in the skeleton with one bond. 
4. Count up t.îïe tat&l number of valence electron*;. Normally t WÉ 
consider only the s and p orbita l #1 entrons as valwc e electrons. 
6orrflt forget the charge on an ifrrcic species. 
5. Subtract til* smühfcer £>£ electrrtn& already >ist^d for the single 
bonds, 
6. Distribute the remaining electrons in pairs around the atoms, 
trying to satisfy the octet rail«. Assign theta tofch*t most 
electronegative atoms first. 
7. Tf ycu run out of electrons Emigre every deserving at cm ha* an 
octet of electrons, you need to form double bonds. 
&*• If you ha*5fe extra electrons and all of the atoms have an octets 
then put the extra electrons on th& central atom arranged as pairs. 
If the central atom is in period 3* 4* 5, or 6 you are allowed to 
have uiore than eight electrons around it. 


NOTEî This is orce of many metï\Gâs useful wh&n drawing Lswis 
structures. If you learned a different taethod, use whichever is 
ea&iest for you. 


•:]yy*'y-y\C-''-<' ,•"'.'• >v:; . 
y • • . . • • • . • • • . , ' • " • . ' " • ' : '-•'• 



7,7: 


MOLECULAR GEOMETRY WORK SESSION 
Using your own paper as a work sheet, draw the Lewis structure 
each of the molecules listed below. Then tabulate your results 
in the correct categories on the attached Results sheets. For 
molecules that exhibit resonance, draw all the resohance structures. 
NH; JieF, XeF: SF4CI2 C1F3 


HF SCI; IF-im4 -£F*~ 

PF; BeCl SO, SO; O U Br 


so3 
2 -P0-3 HCK A1C13 


. • • 


:' y. 


Work out these additional, more difficult molecules. 


OH

CO F2C

H?S04 Ha2V-2. 

CO: co3 
3 H2NS 

j • -\ • 1 

.:.- .-:. .•^-•. .". 


HDLECULML GEOMETRY RESULTS 
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Chaaiatry 131 


INTRODUCTION 


A solution is a homogeneous mixture of two or more substances. By 
homogeneous, we mean that the physical and chemical properties are 
uniform throughout the entire bulX of the mixture. An example can be 
seen by placing several crystals of table salt (NaCl) in room 
temperature water. We would see that the grains of salt disappear. 
Close examination of the water/salt drop shows that we cannot 
differentiate one part of the drop from another* even if we usea 
microscope. In contrast, examination of milk (cow's, monkey's, etc), 
using the eye, would lead you to believe that this can be considered a 
solution. This, however, is nota correct assumption. If we examine 
milk under a microscope, we see small bubble-like globs of fat, 


suspended in a clear fluid. These drops vary in size, and their 


composition is quite different from the clear liquid in which they are 


found. Milk is a mixture, but it is not homogeneous, and is therefore 


not a solution, it is in fact, a heterogeneous mixture. 


THEORY 


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Solut i on Formation: : 


A solution consists of two principal parts? the solvent and the 
solute* The solvent is that part of the solution which is present in 
the greater amount. The solute i» that part which is present in the 
smaller amount. 


We will restrict our discussion of solutions to aqueous ionic 


solutions. In this case water is the solvent and some ionic solid 


(such as NaCl) is the solute. When water and a suitable ionic solid 


are mixed together, the solid is seen to dissolve (disappear). 


The salt crystal is composed of a series of sodium (Nà) and 


chlorine (CI) atoms in a repeating 3-dimensional pattern. The atoms do 


not exist as neutral atoms in the crystal, but rather as ions. An ion 


is a species which has charge, that is, positive or negative charge, it 


gets this charge by gaining or losing electrons. The loss of an -•" 

electron yields a positively charged species, while the gain of an 


electron yields a negatively charged species. Therefore, table salt 


exists as a collection of sodium ions (Na+) and chloride ions {Cl~) as 


shown in Figure 1. The reason sodium is a positive ion (cation) and 


chloride is a negative ion (anion) is due to a transfer of an electron 


from the sodium atom to the chlorine atom. These ions are held together 


in the crystal by electrostatic attraction, due to their individual 


charges. The force of attraction is a function of the magnitudes of the 


charges and the distance between the ions (the farther apart, the weaker 


the Interaction). This force of attraction which holds ions together in 
tat* cry sta 17 is. given the name lattice energy 


solutions Laboratory l 



y • " . 

- -.-••• 

8 7 

• 

• ' : •• 


• • • ".•••• '• . 

Figure 13 NaCl Structure 


". • • . . . . 


What is the electrón configuration of the sodium ions and the chloride 
ions in a crystal cf sodium chloride? 


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Are these electron configurations special (i.e., is there anything 
significant about themj? 


.• • •. 


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.• . 


When N-iCl crystals are placed in water* the electrostatic 
attraction between the ions is broken by the water molecules. This 
allows the Na+ and Cl~ ions to nave apart, and migrate into the sea of 


Solutions Laboratory 2 



-'•' sa


water molecules. Since wats r molecules are polar (i.e., the/ b?jve 
concentrations of "charge" on different parts of the molecule}-, they 
surround the individual sodi urn and chloride ions in such a way that 
they can still move through the solution, but do riot recombine (¿his 
statement pertains to an uns aturated solution, which we shall discuss 
in greater depth in a future lesson.) Figure 2 shows how this might 
look if we could observe ion s and molecules. This interaction between 
water molecules and icns is known as hydration. In this process, we 
see the "breaking" of ionic forces in the crytal, the breaking of 
hydrogen bonds in water and the formation of ion-dipole interactions in 
the solution. 

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Figure 2: Na+ and Cl- Hydration Sphere 


Solubility of Ion£s Compounds 


The driving force of the solution process is the increase in 


•'entropy" of the system. The term "entropy" is used to describe the 
amount bf randomness or disorder of a system. As an example, consider 
your dorm room as a system; does your room stay cJean all the time or 
does it require constant effort to keep it in good shape? The fact 
that it does not stay clean and neat is due to trie fact that the 
entropy of the universe wants to increase and your room is part of the 
universe, therefore, your room will always get messy unless you do 
something to prevent it. 
In the solution process, we increase the disorder of the ionic 
solid when we put that solid in the solvent. At the same time, the 
attractive forces between Jthe solute ions and the solvent molecules 
cause the solution to become more ordered. The amount of disorder 
caused by breaking up the structure of the ionic solid is greater than 
the amount of ordering of the solution, resulting in a net increase in 
entropy. Based en the above statements, all ionic solids should 
dissolve. However, we all know that this is not true. The reason all 
ionic solids do not dissolve is seen to be a function of the lattice 
energy of the solid and the energy required to hydrate that solid. 


Solutions Laboratory 3 


BEST COPY AVAILABLE 


8? 


In order for a solid to dissolve, the hydration energy must be 
larger than the lattice energy of the solid. This is the case in 
systems which yield Exothermic mixing of solutions. In the case of 
endotheritiic solution processes, additional energy is required to 
overcome the lattice energy. This energy is removed from the liquid. 
However, there comes a point where so much additional energy is needed 
to overcome the lattice interactions, that the liquid cannot provide 
enough energy and the solid will not dissolve. 


You will investigate the energy liberated or absorbed during the 
solution process with the aid of a "diode" attached to a multimeter. 
Endothermic mixing causes the meter reading to increase, while 
exothermic reactions cause the meter reading to decrease. 


•• . • ' ' . • . . . • 

Factors Whicï influence Lattice and Hydration Energies 


The smaller an ion, the more concentrated the electrical charge 
compared to a large ion. Thus, we would expect small ions to attract 
water molecules more intensely than large ions. By the same token, we 
would also expect the small ion to interact more intensely with other 
ions in a solid as well. Thus, a small ion will increase both the 
hydration and lattice energies. 


Ions with larger charges will exert a greater force of attraction on 
other ions than ones with small charges. Therefore, increasing atomic 
charge will also increase hydration and lattice energies. 


• . ; • ; ' • \ . . . : . • • ; ' ; • ' • • • ' • : • • ' • • • ' 
Conductivity 


Ions can move in solution implying that charge (determined by the 
presence or lack of electrons) can be carried through the bulk of a 
solution. If an electromotive force, in the form of two leads from a 
battery, is brought into contact with an ionic solution, we observe the 
movement of ions toward the leads of the battery. We also observe a 
current in the external circuit. In this experiment, the external 
circuit current will be used to illuminate a light emitting diode 


(LED) . 


• . ... 
LAB PROCEDURE 


•. -• •; • . • : •• 
• • ' 


A. Conductivity of Solutions 
(1) Construct your conductivity tester as shown in Figure 3. Prop the 
LED up, by placing it in a plastic straw, and inserting the free end of 
the straw in the plastic tray provided. 
• 


.: • 


•" -• ' -A •"••;: . 
• • • . • '• . • 


. . . • 
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. 
.. 

Solutions Laboratory 4 _: • 

. • 
• 

•

• .
'.' •• . . 


90 • 

9jolt 

batUry 

I f -., ^'^¿^ 

LED 

rcsisiar 

i 

drop 

Figure 3Ï Conductivity Tester. 


<2) Place the electrodes in contact with a small amount of crystalized 
Nacl;'. Kcl àhd A1C13* . ; 


Does the LED light, rib?:'••' .-.y.. :7 7:';V • ,'. 7y 

(3) Make a large drop of deionized water on your acetate sheet (the 
large drop should consist of about ± small drops*. 
(4) Insert the ends of your conductivity tester into the deionized 
watery 
7Does the LED light up? 
-V 


(5) Now* make a 2 x 5 matrix of large drops, as shown in Figure 4. 
77r7-tbpy.io^ 
bottom row t> Û o o o 


• 


each circle consists of 4 small drops 


Figure7*i Drop arrangement. 


(6) 
To the top row, sequentially add ojfâ crystal (about the size of a 
On this page) of NAG1* KCl, AgÇl> CaCla and A1C13. 
To ;t]^ó-bottom row, add 'several, crystals of each chemical. 


Solutions Laboratory 5 



• •:• •--<.', 

(ÎJ) The first drop in the top aiKi bottom rows contains NaCl in 


solution. Insert the battery leatfs into the top drop and take note of 


the brightness of the LED. 


How bright does it seem? 


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• ".. 


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-y •• •• • 

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• 

(9) Wipe off the leads with a paper towel, and place then into the 
bottom drop of NaCl solution. 
Compare the brightness of the LSD in each NaCl case; is there a 
difference? If so, why do you think there is 3 difference? 


•• y ' : 

• •• • :•. . 

'.:•.-•' • 

-•.' "ï •

•

-• =

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-., 

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: •• .. . . V \ •. •':'• 

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: ?'. "•• y

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yy 

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(10Ï Perform steps ¡s) and (9) for each of the other sets of drops* 
Write your results below. 


KCl Agci CaCla A1C13 


Top Drop 


.S;-


Bottom Drop 


son 


j: Hake sure you wipe off th-e ends of your leads before you put 
them in a new solution. 
; •£ . 

Solutions Laboratory 6 



B. Endotbermio and Exothermic Mixing 
• 


(1) Set up the meter/diode test apparatus as shown in Figure 5. A 
diode is an electronic device that allows current to flow in only one 
direction. You will use a special type of diode called a thermistor. 
A thermistor is a semiconductor material which has resistance 
proportional to VT - Thus you observe increased resistance with 
decreased temperature. 
• • . , . . • •. 


•'•••": 7 ...•'. ••' A ...y'. 
• 


' . : -:.• 


•. • 


fafe%fl^ 

• 


. • 


Figure 5* Equipment Arrangeaient 


(2) Prior to connecting the meter leads to the diode, turn the meter 
to the "kilo-ohms" setting and press the buttun in the center of the 
dial until "kilo-ohms" is displayed with only one space to jthe right of 
the decimal point* After you have set the meter, connect the diode to 
the meter leads. You should obtain a reading on the meter. If you 
obtain an overload reading on thé meter simply reconnect thé meter 
leads to the opposite leads of the diode. 
(3) Fill one of the wells of the 24-well tray half'full of deionized 
water. Place the diode into the water and allow the meter reading to 
stop fluctuating. 
(4) Add a small amount of NH4CI to th* well and observe the change in 
the meter reading* Remember that endothermic mixing causes the meter 
reading to increase, while exothermic reactions cause the meter reading 
to decrease. 
(5) Remove the diode from the solution, and thoroughly rinse it with 
deionized water. 
(6) Perform the same procedure for NaCl, KCl, MgCl2, Caci2, and AJCI3 
and record your results in chart form on the following page. Classify 
each system as either endothermic or exothermic. 
Solutions;Laboratory 7 



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HEAT OF SOLUTION DATA 


A. 
B 


. 


. 
' 


C. 
D. 
E. 
F. 
• . . 
HH4C1 
. 
• 
. 
. 
• 
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A1C13 yy:, 

Solutions Laboratory 8 



CONCLUSION 


•. 
1* All the solida you used are ionic. Identify the ions and give 
electron configurations for each ion in the five solids. 


• 


2a. Did all the solids dissolve? 


• 


-. • • 


. • • • • . 


2b, If any did not, what might explain the fact that it did not 


dissolve? 


. 


. •• 


3. What would ycu expect to take place if KBr, NaCl and Mgci2 were 
mixed together in a beaker of water? 
• 


So 1 ut i on s Labor a to ry: 9



93/ 


4. How does Cacl2 exist in solution? {Hint: Figure 2) 
: •••


... 


5. What periodic properties/concepts explain the formation 
of an ionic compound? 
" • 


. 


• • . . -: 


• ,-. .. . •. •• y . 
: 


: .... '..•'.•' • yy


6. Which salts produced endothermic mixing and which produced 
exothermic mixing? in terms of Inttio* energy and hydration energy 
gxplfritt why we see energy being released or absorbed in the mixing of 
these salts with water. 
.. 


• . 
Solutions Laboratory 10 



•.'.•• V.-. ••!".• 


t.-.1 y-


Prel&boratory Exúrciae 


1. What's the differences between a solution and mixture? 
2.Whati s the driving forcé in the dissolution process? 


• : . 
3. What is the purpose of the light emitting diode in this experiment? 
4. What is the thermistor used for in part "B"? 
5. What's the difference between an endothermic mixing process and an 
exothermic mixing process? 
y ;_ 


Solutions Laboratory 11 



•97 
. 

• . • :. •

ASI5 RAIB 

131 

••_-.: •• .

ï|^ïii2BHÇT.I<ÈÎï 

The irüíraíSíTíííl aridit y £?£ l&keï atíd natura l vatGi-wsys î:i L-ÜC^VL':: 
year s due. t o acitf ríiiii i s a laager iîonoerfi in xh*$ United S traten., Copiada 
and severa l _E-JÏftpaan natic^s , ' S.c-íd r¿íin .is generate d ïihi&n «3* s s s ^¡..h ds 
su I i' ts.r d í a it i de í £0 ? ) a M ÏÏ i trace n d i ox ici a í NOf ) a £ e hyd n* lys isd i r, w¿ ter . 
Th« í-jí-gh asi d ^oîît^nt c-f íiloudla STX! waterways jias tíAsiaijsii fíírísstu ^ÍVÍ 
th « abilit y of lake s te susta.in aquati c liíe , 

T.tt thi s ejspsrimfei-vt, ycu wil l investigate ! th e effect;,* o£ thés-" 
paîltttaîi t qa<&v~ and nûtTjraliy occurrin g c&rr<uï3i dioxid e £COj) on th s «¿oíd 
conten t of w^.ter-* This wil l b« tí-en* qualitativel y usin g ^mi ^ 
indicato r iïi °ist#r ^s a s^robs fox LVnan^^s in hydrjogetj i o si CCTÜJ.. •'-.<• .n.Ion. 
¥<m wil l ®.lso quantitativel y -ftsteríaiiiíi -the off «it s of each it a s ^ 
titratin g »ate r aoidifíetf by th e OÍ,ses t o determin e th« aci ¿ c . . ->' 
prodiiced... 

•:.•••• •:•• .. • .'. 
.... •

SISOM .... ....... 


• . • . . ••'•'.•"-..• 
y . • • ••• •• 

Ois.fis*ifeioa# 

Açiài For thi * î.ab, ye y il l dsf ía e as, ¿Í^ M e s any substanc e ti^ t 
dissociate s irt water to ^iv e tiydr^sgen ions , H+ < 


Acid Coni^4fi^¿ Acia conten t i a a aeasur a of r.he attac h i-^i-v" tri e 
rAïîBba:r of hyd^o-aô-* i ans i^MÜtóil* frirai s~ íitiá* tí i tri e aüid. ;ir^ ! Í.V CíSíl 


•üvav i*.te vnly on« b.y$r $9es, ioít p*sr mcl«CJI.1« c f acid , w'íi.ii e su¡ I f ¡.ÍT í ¿ " sc i ÍÍ . 
fiíSC^, produces £w& ^ydr^ges. ie¿vs pas? aiïliitiiile a£ a.sid. 
Acid CMï.teiit m&y be measured i n s-evsral wayss* Indicator s Ar<ï cftyrainais 
vSíttft resp^ïîd t o th e ïiydra^ n ÏOSÎÏS ira ablutio n by tírcasvcing color , 
previdinç? VÍS-ÍIS! (Sfss#sssu<sTrt; ©f r¿y£r«gen ion concentration, . A ixuzft r?ora 
accurat e m^asuftsmârn. can to© utaàs fey •*titrâ t i?*q" th e solution . This i s 
dor.® by raa&sttrintf th a amtKsnt a? a feas* (of icnbwjtï ca^Cí&ntration) r&>:iu.ir&a 
t o completely r®aot. with. th e acid . Tf»u «ti l uc^ both Ktîtho-.iiî in thi: ; 

..,-• • y . -7 • .:"•'• .':• •'' .7 A • •' •.-.,• .y : ... -A . ... . ••....•• :; : 

: . .... 

.. . 

. : • • 

• •• . ' • ' • : 

• ..•'• •.•• •.': 
, ' • . • • • • • 

: •. 

•. • • • 
•:,. /.. . • . ••'. • .. y . 
feci» ïiAltr E*3g* i 


98 


Generation of Gases 


Carbon dioxide, sulfur dioxide* and nitrogen monoxide, may be formed by 
reacting certain aqueous salts with an acid. Generation of each gas 
occurs via a complex chemical mechanism. The purpose of this experiment 
is not to investigate these mechanisms but to investigate the effect 


A each of these gases has on water. 


Carbon dioxide, C0Ï ( vill be generated by reacting a solution of sodium 
carbonate, Ha2C03, with nitric acid, HN03. In solution, sodium 
carbonate exists as sodium ions and carbonate ions: 


Na2COa(i, > 2Nay* o + c °i2"(^) Í1) 


in an acidic solution, the carbonate ion «ill combine with hydrogen ions 
to form the more stable carbonic acidi 


• ' . -• •
. • 

• • . • 

«i00!!.,!" .

«V",.,)+ 2HV«,) ¿ ^ »z°°>i«ï-(2) 

: • 

When a solution of carbonic acid is made more acidic, it decomposes to 
form carbon dioxide and water. 


• 


iWm«><=^OT î[* ) * H3°<i): <3> 


To increase thé acidity of the NajCO? solution, a strong acid (KNO3) is 
added, and the gas, C03, is released. Sulfur dioxide and nitrogen 
monoxide «ill be produced in a similar fashion. In the natural 
environment, these gases are generated by a variety of sources. 
However, the gases are identical to those generated in this experiment. 


Concentration of oases 


The balanced equations for the generation of C0J(f) from sodiuiz 
carbonate and SOj(,> from sodium sulfite show that different numbers of 
moles of gas are produced than the number of moles of NO(,j produced 
from sodium nitrite- For every mole of NazC03r one mole of CO; is 
produced. For every mole of Na2S03, one mole of S02 is produced* 
However, for every 3 moles of NaN03, only 2 moles of NO is produced. 


To compara the acid content produced in water from each of these gases, 
you need to insure that equal moles of gas are produced in your 
experiment. You vill do this by diluting the aqueous solutions of the 
salts based on the balanced chemical equations. 


•' • 


• •••:.-. • . -• . . • • •••••: . -. . . -, . .• •• • • • • 

A y : * : . • ..• . •-....• • . • . . . . • .-.. 

y • • ••• • ••-• •••• • A . • . • . y.y • . • •:,-• 

.. • 

• • .. 

7 .'.•••.:' •'•"•' • 

77 A ;,y.-y ••"•. :•••••••-• -. y y ...-• • • •. .'y• • '•: • . 7-
. y •••••• . • •" 

•; .
ACID RAIN Pag e 2 

• . • 

' • -' • '

/ •• .,:'•-•• •' A •:•-.-. ;..y; 


• 


• : 
• 


•" .
Effect of Gases on water 


The gases that you generate will react with water in a drop of Yamada 
indicator to produce acids as followsî 


™ 2u) +H2°<u; ±H 
2 
ra3<aq) (4) 

. :• • 
• 


. 


(Note: This is the reverse of equation 3.) 


Yamada indicator will be violet to blue colored if the solution is 
basicr green for neutral solutions, and yellow to red for acidic 
solutions. 

Sulfur dioxide reacts in a similar fashion to form sulfurous acid: 


so + H,0 ^H^SO, (5) 


i(t> "a"tU<-, M.";W 3(J<] ) 


• 


we'll not consider the effect of nitrogen monoxide on water because it 
quickly reacts with oxygen in the air to form nitrogen dioxide. 
Nitrogen dioxide then reacts with water forming nitric acid and nitrous 


• ••' 


2ND (.)


 +°*U> *> 2NÜ*U ) 
(6Ï 
-


• .' • :
2HD1(<> +H 1 0 m ; 

±HNOM.<0 +mQ2(^) 
• 


!•.-.

Generation of Sulfuric Acid • y. 

r 


A current area of scientific study is how sulfuric acid is formed in the 
atmosphere. Two hypotheses have been identified. The first states that 
sulfuric acid forms within water droplets when sulfurous acid is 
oxidised by nitrous acid: 


2H? S03(a q, + 2HN03( .q } > 2HjS0Maq , + N;-0(aqj + H2 0( i , £7) 

The second hypothesis is that sulfurcue acid is simply oxidised to 


sulfuric acid by atmospheric oxygen. The difference in these two 


hypotheses is the presence of nitrous acid. You will design an 


experiment to test each hypothesis, and draw your own conclusions. 


• 


-.:. • 
'"' • • ','..:' A 


ACID EtAIH Page 3 
-7 •• '•• y ' 
• •'• : 
.. y. : yy



JJtPERlHEHMi. 


A. COKCËHTRATIOH OP GA6EÛ OBHERATED 7 : 
1. Compare the balanced equations for generating CO3, SO=, and H O 
.from-"'their respective aqueous salts* 
The balanced equation, for C03 hà'ë7': 
been provided (equations 1-3). You must write and balance equations for 
SO; and HO (see paragraphs C.l.a and Û.l.a belov)* Based on these 
equations, determine the ratio of moles of aqueous salt to moles of gas 
produced from each salt. You must adjust the concentrations of the salt 
solutions to insure the same number of moles of gas will be produced in 
all three cases. Write the ratios below: 

NaH02(a<1)/HO

^^¿t-MÏ^Ïf = • ^3<*q>'S° 2U) 
t*> 

2. Your ratios should indicate that sodium sulfite and sodium 
carbonate produce the same amount of gas* while sodium nitrite produces 
comparatively less. Both the sodium carbonate and sodium sulfite must 
be diluted to produce the same amount of gas as sodium nitrite. Based 
on the ratios from paragraph A.I.-, calculate the new concentration (in 
molarity) of sodium carbonate and sodium sulfite solutions needed (the 
stock solutions are all 1.0 H): 
NewUa2C03 Rev BA2 S03 
Concentretion Concentration 


ACID RAIN Page 4 


: LUI 


a. Now use the relationship, M ^ = M; VJ , to calculate thé number: 
of drops of aqueous salt solution and deionized water that must be mixed 
to obtain the required concentrations. Recall that M is molarity of thé 
solution, set Vi=i drop and solve for V2„ V2 is the new total volume 
(in drops) required to obtain the appropriate concentration. By using 
V3-I drop, you will obtain the number of drops of water to add to dilute 
the solution (the -1 accounts for your initial drop of solution). For 
both Na2C03 and Na2S03 ( show V2, the number of drops of 1*0 M stock 
solution, and the number of drops of deionized water mixed to obtain the 
correct concentration: 


' 7 • 


Dilution Dilution 
Dat a W<t2C03 Data Na=S03 


• . . 
• .• • • y 

. 

. -• •• 

. •• -• • 


• ' . 
. • 

• 

• . 
• 


• • . 

^^^^^^^^^— « 


': •'." y •':•.;• .'.'• • 


8* EFFECT OF CO* OH WATS* 


1. Dilute your stock solution of sodium carbonate as determined by 
the above calculations to obtain the necessary concentration. Do this 
by placing several drops of your 1.0 M stock solution on a clean area of 
the reaction surface and add the correct number of drops of deionized 
water. :.y.y 
2. Place two drops of Yamada indicator in the petri dish. 
3. Rinse an unlabeled microburet several times with deionized 
water and then use it to transfer your diluted Na£C03. Place two drops 
of the diluted Ha3C03 near but not touching the drop of indicator and 
cover the dish. It will be important to have the same drop size in each 
section so the same unlabeled microburet must be used each time. 
4. Carefully lift an edge of the lid and add one drop of 3 M HNO
3 


directly on top of the drop of Na2CO¿. Describe what happens to both 
drops. Continue to make observations for at least two minutés. 


a. Evolution of carbon dioxide will be indicated by the 
appearance of bubbles (equations 1-3). The color change of the 
indicator is due to the formation of carbonic acid as the carbon dioxide 
dissolves into the drop and reacts with the water. 
b* Carbonic acid is a weak acid. This means that it doesn't 


dissociate completely to give 100% hydronium (H30*j ions and bicarbonate 
(HC03~) ions. Instead, it partially dissociates, and an equilibrium is 
é stab l i shed.:-,'y • •-; . 

ACID RAIN Page 5 


• 


+ BCD
"Ac,) + Mΰ(1) ^Z _ ¥+ (.q )3 ' U(i ÎB) 


5. Repeat steps 2-4 using 15-20 drops of freshly drawn deionised 
water instead of the indicator. Rinse out an unlabeled microburet with 
deionized water* After two minutes, insert your cleaned, unlabeled 
microburet through the entry port to remove some of the acidified water. 
Place three drops of this solution on the reaction surface and add one 
drop of indicator to it. 
6. Titrate this solution of aqueous carbonic acid to determine haw 
much acid was produced by the carbon dioxide gas. The procedure to 
titrate is: 
a. First empty an unlabeled microburet and rinse it several 
times with deionized water. 
b* Place 10-15 drops of freshly drawn O.O01 M sodium 
hydroxide (NaOH) onto the reaction surface and siphon it up with the 
clean, unlabeled microburet. Always transfer the acidified water and 
HaOH with the same microburet to ensure the drop sizes are always the 
same. He will assume that one drop baa a volume of 2.0acl0"s L. 


c. Carefully add the base (HaOH) one drop at a time to the 
acidified water on your reaction surface. Continue to carefully add 
dropwise until the Indicator turns green (indicating the solution is no 
longer acidic). Count the number of drops added. At this "equivalence 
point11 , the number of moles of base added equals the number of moles of 
hydrogen ions present in solution* 
d. If the indicator turned blue, you have passed the 
equivalence point. Assume that 1/2 of a drop of base would have yielded 
the green color, then* total the number of drops of base added. 
e. Record the colors of your acidified water with indicator 
before and after the titration. Also record the number of drops NaOH 
added. 
Color of acidified water from generation of CO2 


1&&Z* ht%W 

. 

.Number of drops 0.001 H NaÔH added: 


7.7 Clean and dry the petri dish. 
ACID RAIN Page 6 



C* EFFECT OF S02 OH HATER 


1. Use sodium sulfite (Na2SOa) to perform an experiment showing 
the effect of 50 ¡ on water. Follow the same procedure as you used in 
Showing the effect Of CO* On water (Section B). (Remember to change the 
concentration of your stock solution to allow for genarat ion of equal 
mal» s o f SOi , .ÇO-ii.'-. and HO. •). 
CAÜTIOHI ALWAYS NEUTRALIZB THE POISONOUS BD2 BY ADDING 3.0 H 
AMMONIUM HYDROXIDE (NH4OH) TO THE PETRI DISH PRIOR TO REMOVING THB LID 
AND AFTER YOU HAVE REMOVED YOUR ACIDIFIED SAMPLE» 


a* The reaction of Na¿so3 and acid is similar to the reaction 
Of acid and Ha2CO^ (equations 1-3). Hrite the chemical equations for 
the generation of SO;(I> from Na^S03(iq): 


-'.-. 7 •;• 

';; 

_ •£ 7 

. • •-,.

yy .yy"-- yy. 

••'y-'y. ' 

b. What acid is formed when SO2 dissolves in water? 
ACID RAIN Page 7 



.•t- •'. . _ J


ro4 


c. Thé formation of this acid is analogous to the formation 
of carbonic acid. Write the chemical reaction that describes the 
formation of açid when water reacts with S02. 
d. After diluting your stock solution of sodium sulfite, 
transfer two drops of the diluted solution with an unlabeled raicroburet 
(don't forget to rinse it first with déionizad water) to the petri dish. 
Add nitric acid to generats S02 and show its effect on a drop of Yamada 
indicator. REMEMBER T O NEUTRALIZE THE S03 BEFORE LIFTING THE LID OFFÍ 
; e* Compare the'effect ofCpj versus sù2 on the indicator. 

' • ;•' -•si 1 ; •' 


yy..y. 2. Repeat the generation of SO; using freshly drawn deionized 
water and your diluted Ha2so3. Use the same amount of deionized water 
as you used in B*5*b. Collect a sample after two minutes and titrate 
this sample using 0.001 M NáOH. The titration procedure is outlined in 
paragraphs B.e.a - B*6.e* REMEMBER TO NEUTRALISE THB S02! Record your 

^results belowi 


•77:7: Color change. 
Number of drops base added: 


a., Clean and dry the petri di»h. 


EFFECT OF NOî OH HATER 


l. Perform an experiment to show the effect of nitrogen dioxide 
on water using 1.0 H NaNOr. YOD MUST NEUTRALISE THE HO/HO, 
IH THE PETRI DISH WITH HH4OH| 


ACID RAIN Page 6 



á. The following two equations have been fille d in for you. 

Fil l in the las t equation for the formation of NO in an acidified drop 

of water: 

. 

Ha SO. 

• . 
•*••*.M> +NtViM> 

->• HWO 

H03"lM) + H+(MK-2l*t) 

ME> + H„0 + UNO 

*= *t,! ?t»<i> 

nu 

-

b. write the equation for the reaction of N02 and water. 
• . . •• 
• 


Generate HO to show its effect on water. Describe what 
you see: 


. .•. 


d* The brown colored gas produced is Ho2. Nitrogen monoxide 
(NO) is an unstable gas, and is immediately oxidized to H02 by 
atmospheric oxygen in the petri dish. Hitrogeii dioxide is a poisonous 
gas. IT TOO MUST BE NEUTRALIZED BY ADDING NH4QH TO THE PETRI DISH* 
followed by a wait of approximately thirty seconds* 


2. After neutralizing any excess HO2, clean and dry the petri 
dish. 
3. Titrate a sample of water acidified with N02. Record your 
results belOW. REMEMBER TO NEUTRALIZE THE POISONOUS NO2 GAS USING NH4OH 
AFTER REMOVING YOUR ACIDIFIED WATER SAMPLE, BUT PRIOR TO REMOVING THE 
PETRI DISH LID» 
• . • • • . 
4. Clean and dry the petri dish. 
• . 
ACID RAIH Page 9 


\í)b 

. 

E. FORMATION OF SULFURIC ACID 
Design an experiment to test the two hypotheses about the formation 
of sulfuric acid in a rain drop. The experiment should have at least 
three parts: 


a. Procedure (experimental)» Outline your procedure to illustrate 
what you will do. Drawings may be helpful. 
b. Observations (results). Collect your observations in a way 
that is easily related to your procedure. 
c. Conclusions. Evaluate each hypothesis in light of the 
observations made during your experiment. 
To help you get started* you have been provided with O.l M 
Ba(N03)2 * This will be used as a probe to detect the presence of so4 
2" 
by giving a white precipitate, indicating the formation of sulfuric acid 
as opposed to sulfurons acid. 


• . 
• . 
• •; 


7 . 
•• 
. • . • 
• . 
. 


• 


• • .. . . . . • • . . -. . . • . 


... 


. > • ... 
• 


• •••• . 


•


After completing your experiment and recording all necessary 
information, clean and dry the petri dish and reaction surface. Rinse 
out the unlabeled microburet several times with deionized water, and 
perform any other necessary lab clean up. 


... . -••:: 

ACID RAIN Page 10 



RESULTS 


1. Calculate the number of moles of acid produced from the 
generation of each gas. Use the data from each titration that you 
performed, and recall that at the equivalence point, the number of moles 
of acid present equals the number of moles base added. 
a. From COj: 
• 


• 


• •••• . 


b. From soaï 
. • • • • . . . . . . -• . . ' . -' : 


c. From H02r 
• 


2. Make a table listing each gas, the acid content produced, and the 
indicator colors at the beginning and end of each titration. What 
comparisons can be made concerning the capability of each gas to affect 
the environment*1 
•


y . 
• 


. . • 


..... 


• • y . • ' A 


ACID RAIN Page 11 



10(3-7 


COHCLUaiOMS 
yl:.; 7What7ÇMivbeh concluded about the capability of each gas to 


increase thi acid content of water?; 


2* Write balanced equations to explain how each gas forms an acid 


with water. 


* ;.3.* What: cisn be :¿oncl uded about : the ac idit y of ai I ra in va te r g i ve n 
that carbon dioxide is a naturally occurring gas? 


y' y 

'.-:. t 

4. Explain why acid rain is primarily a probiera in the Eastern 
United States and Canada as opposed to the Western United States. 


'•->' ' yy, 

•5k-y- How.do you think We can soIVe the acid rain problem for the 
j- ¡^ i--m ^ y- •'•• :•-• y. -'' .' ."•• : .•"••••.••' . •». ••* -\ '•.'•.'•• • "'"•. . •-••.•" • 

,


t<nií-:.í; -'•..:,:-.i::-.::\'-.-. .-.-... -yy y.:.-, . \I yyy y '-- '.riV • • A --'y." ^y I. ' 

••^-•-•••• 

ACID RAIN Page 12 


~ •• m'.~A 


PRE LAB EXERCISE 
ay7 List those gases that react with atmospheric water to causé it to be 


acidic, .. .. 


2. Write the equation for the formation of acid when SOf reacts with 
water (hydrolyzes): 


3. What is "acid content"? 
4. Do MCI and H2SO* have the some acid content? 
ACI D RAI N Page 13. 



y^& 

... .•••.• .H .

: '.O' 

.:«• • 

--•m. • 

•. .-a • 

y ••'•:: • S •• ' 

' ' ' •• .:• '•' y . y 

-.:•'• •.: -l :••<• • 

•i ••: ••••yy 

y.y •. • 7 yy 
ui 


-. . . . . . . . . • . . : ... 
• • ' . ' . • ' : • •" ' • • ' • " • . • •' • . . . 
SPECTROMETRY C DETERMINATION 
OF AN EQUILIBRIUM CONSTANT . • ' ... • • .. •• 
chemistry 131 
• • • 
. • 
••s :'• •'•-• •-••• ••:: ••," •.. 
A • 
• ; ' • y ••' . ' -• , . • • • -• • . • , • • . • . 
INTRODUCTA 
-"-• • .. • • 
'.'•'...•••' •••'.y 
•-'-. • '.•,.'• ' .• ' ' • '•• " ' 
Do all chemical reactions go to completion? Do they progress 

until all of one reactant is completely used up or do they proceed 
to only a limited extent, or for that matter not at all? These are 


important questions concerning chemical equilibria which will be 
discussed in this laboratory. 


A condition of equilibrium exists in a chemical system when two 
opposing reactions occur simultaneously at the same rate. As a 
result, products are being formed from reactants as quickly as 
reactants are being produced from products. It is important to note 
that after equilibrium has been established, there will be some 
products formed and some reactants remaining. 


Not all reactions go to completion. An example of one that 
does go to completion is the dissociation of a strong acid, such as 
HCl, in water: 


••'••• ••:•". ' y . . ' . . -. . • . y -• • •. ..:. .. ' y •'•... • .-.'•:' ' • • • • " , : • ' • -. -. • • . 

The single arrow indicates practically complete product formation 
{H30* and Cl"); we would expect no significant amount of reactants 
tb re**ih if one mole of HCl is reacted with one mole of water. 
On the other hand, a system at equilibrium such as the reaction 
of iron(III) ion with the thiocyanate ion: 
. • . 
. • 
' 

is expressed using two arrows pointing in opposite directions to 


indicate a state of equilibrium. 


"• • . •=•.-.••• ' ' • * ' • " 7 ; 


Equation (2) is a net ionic equation where Fe3 + and SCN" 


represent the reactants and FeSCN2 + represents the product. You 

cannot simply go to the corner store and pick up some Fe3 4 or SCN" 


ions. In this lab the Fe3 + ion comes from an Fe(N03)a (iron (III) 


nitrate) solution and the SCN' ion comes from a KSCN (potassium 


thiocyanate) solution. When these species are mixed in an aqueous 
.^m.piilt ibht the fol 1 owing ions formiy. 

C-yí^ y%i*iyyyy ' ..yy yyy '^-:& 

EQUILIBRIUM CONSTANT Page 1 



. -• . • • • • ' " ' ' ' ' ' • . . . . . . y '• ' ' . • ; ' . " ' . -• ' ' • • • • • ; ' '.' 


. ' •'•' '. yy •' . . ' . • . . ' : • '••-•• ' • " '• • ":•' • 7 • •


i i a 


"•••••••: • ' • • 


y • ..... -. . 


• • • " • • : • ' • • • . ' . ' . • • ' -. . ' • . . . . . . . '. 


The K*{aq) and NOi"<aq) ions remain in solution unchanged from 
reactants to products. Ions that remain unchanged in the overall 
reaction are known as spectator ions and are not shown in the net 
ionic equation. 


You can derive the correct equilibrium constant expression, Kc, 
for equation (2): 


{products ] [FeSCÍí**] 

K - ~~——"" -. (3) ; 


* [reactants] [Fe3*][SCin 
• A : '''.' •••'••': ' . • " . • • '-.-' • ' • 7 '. ' 7 • • ' •' '. '•'•' ' •. y •''."' 
Ke is the ratio of equilibrium product concentrations (the brackets 


denote concentrations in molarity} to equilibrium react«ne 


concentrationsy Each concentration term i« raised to a power given 


by the number of moles of that substance appearing in the balanced 


chemical equation. In equation (3), all molar coefficients are 


equal to 1* 


DRiFPTiurn 


• A • • ' • . , ' : . -. ' . ' . ' . . • • • • ': •: :•' . ' • : • ' ' 
• . • • :. . , . . • • '•..••• : • ' . • : . . 
There are three things you should accomplish during this 
laboratory.. 


1. Learn how a standard laboratory instrument* in this 
case á Spectrometer* is used to collect data on chemical reactions. 
• 


2. Determine the equilibrium constant for a chemical 
-•


reaction* 


3. Determine AH.AS , and aG for a chemical reaction. 
BACKGROUND 


• '• -• • • • •'. ' , " • • • • • • . • • . . • • . • -. ' 
• . ' . • -• • • y l • . . ' ' • -• • ' • ' . • . . . . • • . . . : • ' . " ' ' • . ' . . ' 
In this experiment you will determine the value of the 
equilibrium constant (Kc) for the reaction: 


• " . -• , _ • ' . ' . • ' A : . • • • • ' --. " ' ' . ' . ' • ' • • • . • -' ' : . ' • . • • ' -. . . -. 
• ' • • . ' _ . . • 


: • • • y • • • • • . • -. • . • • . • . • , • . , • • • 

. • • . • ' : • • ' . . • • " ' . . " . . " : 


Un fortunate 1 y for us,: we do hot know the equ i 1 ibr ium 
concentrations of Fe3*, SCN", or FeSCN2*, but only the initial 
concentrations of the two reactants, Fea + and SCN". The equilibrium 
concentrations of FeSCN2+ is determined colorimetrically using the 
spectrometer, iron(III) thiocyanate ion is a red-orange color and 
the intensity of color increases with increasing FeSCN*+ 
concentration in solution. TO determine the equilibrium 
concentrations of Fe3* and SCH", we use the fact that Fe3 + and SCN" 
can exist either as uncomplexed (free) ions or as complexed FeSCN1*. 


-


The equilibrium concentrations of Fe3 + and SCN are found by the 
difference between the initial concentration of ion added to the 
solution and the concentration of complexed ion (FeSCN2+). 


• . • 


A. ; 
EQUILIBRIUM CONSTANT Page 2 

• • •. 



7.This reaction must be conducted ina moderately strong: acid 
solution to prevent the formation of th<& FeOHî + ion: 


; F ^ ^ 
•-"•'?:•')••• •'•^' : 7 


.\.' .• -i/^l' 


Thé FeOH3 + Ion absorbs light at the same wavelength as FeS^if?+ .' To 
have FeOH2+ pre s en t w ou 1 d eau se cons i d e r ab3. s exp 3 r imen tal e r r o r. 


:Acüí:,s;^Í7fg to Le Chatelier's principle, addition of H*, shifts this 
equilibrium to the left* minimizing the formation of FeOH2*. 


You will usea Spectronic ¿¡o spectrcri&ter to measure the aw.ovtfit 

o :f 1 i gftt ates o ¿-te s d i iî.bai-o zhayn ¡r>G ;?fcy t'ft s F s S Ctî*: st c;.u t..! t>tt. TÏI.s 
ft b Í¿ o rbs fiCftf A, o í i: hô so 1 u¡ t. i w Ï. i-y j^roi;c Î* í i o«A 1 t o thef^£OTí* 
concentration ([FeSCN2*]) according to the following equation: 
J¿y¿^[FeSCH*+] 
.(4) 

where -f and * ar e ¡satiricall y determined consstwit's* This . 
s-ftl&tiimahip i s fciiEww as Ese*1-'s Law; * i e v^h^ si\Glar .absorptivity , ¿.n 
inter s i v© pîfop&rty cf th e liffhi* ^bssovbi?^ ^p^ci^iy .sncí t i s '¿he 
efrocxiví-i p*t¿>. .IstfiCjth- or dí^tíüStü& t)i'-; li^i t ¿FU •:;;-*• t.rsve l thras h the. 
s o 1 ut. J. isï-. 1 ri t he. -ep * ;: t r cnitó cer . For s q^ieo^ £ a o i y t i oïï& O Z Fe 3 Cïi * * , th e 
aialar .absorptivit y i s ¡* - 4,70 x i£3 M"-i,c¡a",i &£ tïj& vav^iür^t h of 
maximum absc-t'bancé (447 nm) . The pat h lengt h i s th e interna l 
diameter of th e specia l tub e tcuve-ttaj Hs&d t o hold the solution . 
You wil l us e a cuvett e with 1-- 1.17 cm. Thus, for aqueous 
solutions of FeSCN2+ the concentration of FeSCNa>i s given by: 

y-'.;7-.>yA:-;v; ;; ••;.;.•.-y .77^. 7 777y¿.7.<y .: 
.'7" 77' 

*y 
[Fe5C»?+l - -? * -, -———— > y-.- (5) 
,;':st'-' y'y'S^SO :x:iO*'Hy ; 7.-y: . :'y'y 
íilthotish absórbanos ca« be r^ad directly , tíis s^slfe 5s 
noríí.i^ípftr. This s:ak#& i t dií't^c^I t t o re? ¡sa, and ver y fc&.^y t o r&card 

•; ÏI iso irrfte i; dat a H l" Ï>S t SÎÏÎHJ , r^^s thí¡ pt^ ¿j «#Î fît t ï s nsB' J. tt a n c <? „-( ï-T} , sc ^ 1 e 
since i t i s linear and easie r t o read. 
You may then d«t<É:roei.r.«! the absórbanos: 


7"i^'yio§::T:; 
;'.; -;V( 6 | •;. 


(Hots; Us** T i n eq\.&tio*ï 6 üíjdl r¡wt % T.) A sample <yü very ïiigiï 
cclo v lntîSWK:î.tv aíMSf'íífhí virtuall y al l fctec li^h t f^-iu th « 
sp.ftiftrciiftfttftr, Thi s çlw<ss o% trA^aiittSii^ï ! c r ir¿f¿iiits absorba n^e. 


cci«vafrsfe:L^,

 as

:

-;™ïJ*t sawjuls

 &xlutlctv. i*aci£H?.e:ii3 îscriï dilute , KUV*Í .líqht 

i s t.K'ffiíisíñi.tted • through

 tÁ« s^lwtÍGri.

 At 2«r« iOJiceHt'iratlsK the; 

trsinso® itt^iiscè -of. • ligh t i s '100%-(T?41) • end• ' th& absin^'isa «ca is-' ser6," • 

EQUIiUBgí.Wi ft^SST^ÍÍT Ç^qe 3 


ÏAÏI4 


You now know how tc calculate the absorbance (equation $} and 
thus, you use this number to determine the equilibrium concentration 
of FeSCN?+ (equation 5) . From equation (a), we know that the 
equilibrium concentration of Fe5 + f[Fe * + ],<,) may be found by: 

7 [Fe^]fC 5 = [Fea*li»iti.i -îFeSCNi + ] B q (7) 


TEHPE»ftTU»E! errecTs 


The dependence of the equilibrium constant on temperature 
indicates whether the r~nation is endothermic or exothermic. If the 
equilibrium constant decreases a3 the temperature is increased, we 
know that formation of the products compared to the reactants is 
less favored at this temperature than at the lower temperature? 
thus* according to Le Chatelieras principle¡ the reaction is 
exothermic. By a similar process of deduction, we could conclude 
that if the equilibrium constant inoresses as the temperature 
increases, the reaction must be endothermic„ 

•yyy. . ••''•• y y •. • • " ...-.' 
y Recall from your study of thermochemistry that: 

¿G^AH-TÀ S (3) 


• . _ " • • • • .;•• ,. 7 -. " 
• .• •
and 


. 
AG = -RT in Kc (9) 


.-•'• • . . '. 


From data you collect in this experiment, you can solve for all 


these thermochemical parameters: 


1. By measuring transmittance and calculating equilibrium 
concentrations, you can find the value for Kc and from this you can 
solve for aiS using equation 9. 
2. By studying the reaction at different temperatures and 
calculating the equilibrium constant, it is possible to calculate 
the heat of reaction, AH, and the change in entropy aS. 
By substituting the value of &G from equation 9 into equation 8 you 

obtain: 


-KT In Ke - AH - TaS y (Í&) 


¡yyyyy'yy 'y.-'.7.7y-' • • • • .."••.'-: y "'•,' ••'•.•.••••••..••• •..'•.,• • ••'•.'•. yy': 


Dividing both sides by -RT you obtain: 

•

In Kc -J2Lx|-L| + .±l (liy 

R i I TJ..-7Rv. • 

• . • • . . • • . . ... . 
• -.


••-.. ; y ... . . -. • 

• :•': • • .: . 
... • .• • • • : •• •• 

• . 
. .. • 

--•. 
. -• 


• • . -• • • • 

EQUILIBRIUM CONSTANT p^ge 4 



. • ' 

• • • • • 7 •: •' 

'•••.= ' 

• 

. 

This equation i s i n the form of y = mx + b i n which 

y » lnKc ; X - 1/T ; b - .££ and m-" £ü 

• •••«.• t 


The reciprocal of the temperature (1/T) in Kelvin plotted versus the 
natural logarithm of the equilibrium constant (In Kt ) , should be 
straight line with a slope of -AH/R where R is the gas constant, a 


B.314 J/(mole K) and AH is the heat of reaction. What does the 
slope of this line tell you about the exothermicity or 
endothermicity of the reaction? 
• . 
•.. • 


• 


7 •' 


. 


. 7 y • . 


3* Once AG and AH are calculated, it's a simple matter to find 
A S from Equation (8)• You may also determine *s from the 
y-intercept of the graph mentioned in paragraph 2 (y-intercept 
equals A£/R) . 

• • • . ' ' ' ' ' 


. •• • . • • • . . . • • • • . -• -. . . .

EXPERIMENTAL • . 

7 


• ' 
NOTE: Fe(H03)3 IS ACIDIC. BOTH Fo(K03), and KSCN . : 
VILL STAIN THÜ LAB BENCHES. BE CAREFUL VITH THESE 

. ' • • . 

SOLUTIONS AND WIP E U P YOU R SPILLS IMMEDIATELY. 


• 

SPONGES, CLEANSER AND ELBOW CREASE ARE THE ONLY 
KETHOD YOU VILL USE TO CLEAN THE tAB BENCHES. 


A . PART 1 : TRftNSMITTANCE HEASUREHENT 
1. Thoroughly clean, rinse and dry eight 250 mL or 100 mL 
beakers. Number the beakers one through eight with a grease pencil. 
Remember, all of your solutions will have the same volume of 
40.00 mL. 
2. Using a clean beaker, obtain about 100 mL of 0*00200 M 
KSCN stock solution from the side shelf. In another beaker obtain 
about 100 raL of a mixed solution which is 0.0030O H with respect to 
re(NOa)g and 2.00 H with respect to HN03, Record the exact 
concentration of the stock solutions used. 
3. Fill your clean and marked beakers with these reagents In 
accordance with Table 1. Use a graduated cylinder to measure the 
reagents and de-ionized water. 
EQUILIBRIUM CONSTANT Page 5 


.--•y • 



• í 
• . • 


IIS 


. • 


TABLE 1 


."-' •.'•• • • ; 


Ë&&KËE Fe<N03),/HW0, ¡Ç5ÇM De-ionized Total 
VOLUME 


y •


:.ly:--10,00 ML 4.00 M. 26.00 ni. 40,00 mL 


A2À 10.00 aL fi.00aL 24.00 aL 40.00 i«L 

:-3-10.00 mL 8,00 SL 22.00 uL 40,00 aL 


4 10.00 aL 10,00 mL 20.00 SL 40.00 aL 
v'A y 10.00 aL': 12,00 aL 18.00 SL 40,00 nL 


v'6-.' 10.00 aL 14.00 ML 16,00 SL 40.00 -L 


• 7 . ;'. . 10.00 et";. 16.00 HL 16.00 mL 40,00 aL . 
... 
. 


S 10.00 et7 18.00 SL 12.00 mL 40,00 aL 


4. Mix each solution with a clean, dry stirring rod; rinse and 
dry your stirring rod after each solution. Record the actual 
vo1umes used. 


5. Be sure that the wavelength is set to 447 nra on the SPEC 
20* Calibrate the Spec 20 using the instructions on the last page 
of this handout. Now you are ready to measure the percent 
transmittance of the solutions* 
6. Measure the temperature of one of the solutions. Assume 
that the temperature is the same for all beakers. Record the 
temperature below. 
7. Obtain a cuvette from your instructor* 
a. Be sure the cuvette is clean. Fill the cuvette up to the 
bottom of the frosted circle with solution from beaker 1. 
9. Place the cuvette in the SPEC 20 with the frosted circle 
aligned with the mark on the SPEC 20. Close the cover over the 
cuvette. : 
s


10. Record the percent transmittance. Use TABLE 2 to record 
the percent transmittance* 
NOTE: Don't allow the cuvette to remain in the SPEC 20 for an 
extended period for it will warm the solution and change the 
reading. 


EQUILIBRIUM CONSTANT Page;6 



.Wx 


11. Remove the cuvette form the SPEC 20. Discard the solution 
127 Rinse the cuvette with a SKIA31 (less than 1 mL) amount of 
solution from the next beaker. Discard the rinse solutionr Fill 
the ouvette with fresh solution from th* same beaker and measure the 
percent transmittance as in steps9 through li* 

NOTE: Be sure to rihse out the cuvette with the solution you 
are going to measure next* This will avoid cseintaminstLesi of the new 
solution to be measured by the previous solution, Also( don't rinse 
the cuvette with water between measurements; lingering wat^r will 
dilute the solution to be measured resulting in an incorrect 

13. Repeat the procedure in step 12 for the remaining beakers 
numbered^ 3^ through 78yyyyy ' y.7; 

;íi77|^t;;Í!.;7:.TENPERftTÜ«e PEP EM PENCE Q$.&¿;h :'/• 

1. Set up ft cold and a hot water bath. 
•2V: Fill the blue pneumatic trough with an>ice water solution* 
(Add approximately one 400 mL beaker of Ice*) 
3. Fill another pneumatic trough with, hot tap water. 
•. •••..' * • • * : 
4* Flaca beaker» 1 through 4 in the cold bath for at least ten 
minutes* 


-yiy^y:--place beakers 5 through 8 In the hot water batli for at. 
least ten minutes. . 


NOTE* The next step© must be carried out in a TinELY FASHIO N 
to prevent large temperature changes from occurring. 


yyy 6. Remove beaker 5 from the hot water bath* At this time 
measure the temperature of the solution» Record this temperature 
below.- "yy -.y- .-.yy.vy ;;: -..y. 

• y.-. h--r 
7. Fill the cuvette and measure the percent transmittance as 
you did in Part 1. Record the percent transmittanae. 
,.'•••'•''-'••' ' ' ,. y y . y " .-•• 7. •yy 7. ' yyy 

B, Repeat steps 6 and 7 with beakers 6 through Ô. Be sure to 
rinse the cuvette with the solution before measuring the percent 
tr an em itt anee of the solution-^yfySy 


9. Hbw you cen measure the Percent transmittance of beakers 1 
through 4, in the cold water bath, following the procedure in this 
part steps 6 through S. 7y ':::': 
! . ..;•
y? 
y y. y 
ECU ILjBRI UN CONSTANT Page 7 y . .i--,..-. ._• 
77 : 
f '.y-y¡ i yf" y A ••• " yy y i:..•'*>•• 
7 .7 7' *"'''•'. "'"j ••' 


1 L8 


ROOM TEMPERATURE:
BËÂ ~T"
KER *T A i[Fe3+] 
TEMPERATURE"
C 
INITIAL fEQUILIBRIUM 
[SCN~] i[Fe3*] [SCIT] [P&BCN34 ,
~ 
| ^ c 
—i —;—) 1x| I I 'y-y\y."H 
i:; :\ 7, i41 4^4- 
r^ 1 1 •.'•• i 
• ¡y yI r . f 
1 ;• 7.}':, J:;. ;,|.;.-
6I 
7 t 
B| 
I . I: 
-4-4 
+-+ 
.171 A:,-I: 
i i 
i ' •• 
^^á^U U • .
• 
• 
y •
. • 
•••.'.'..•• • ... 
COLD TEMPERATURE: TEKPERATURE11 
T T 
2» 
3 I Í 
+•^-4 
4-
TI : : I T 
• 
• • 
• •
.y-. A... 
" . 
• • • • 
HOT TEHPEBATURE: TEMPERATURE-
T—r 
81 L 
X T T 
4H—I + mm>: 
T#-€-4 
fe#+ 
4—-4 
T 
• : 
.: 
: • ' • • 
• •
. . 
• .
. • 
• • . 
• • 

. . • ; • 


... 


EQUILIBRIUM CONSTANT Pag e 8 



A<A'y" 7:yA' m:,:: ;r'.'.•:'•' y:": . 
y";-y'-'^f 

LABORATORY PEPOHT CEQUI REMENTS 


'The¿ab^portviiiUBt include the following: 


1. Statement of objective 
7 2* Theory (Ï-2 pages). 


3. Results - since there are many calculations that must be 
made on this lab* the use of a spreadsheet program is recommended. 
Using the data collected from Part l and Part 2 calculate the 
following and include it in Table 2. 
a. Absorbance from the transmittance data (show a sample 
calculation below): 
b. Initial concentrations of Fe3*, and SCH" [show a 
sample calculation below): 
••' Ó *7 Egu i 1 ibr ium concentrât ions qf Fa3+ ,. se w-, a nd Fes CN * * 
(show a sample calculation)t 

d. Equilibrium constant at all three temperatures for 
equation (2) (show a sample calculation): 
e. 
The average Kc at each temperature. 
K« at room temperature: 
Ks at cold temperature: 
Kc at hot temperature: 
y±^ 
*G using equation: (9).77 

:g-,.: AH us i ng ; a graph of 1/T versus in Kc and u s ing 
eg>iatiott^iil.;7y7 •'>•:..' ..;. 


h. A3 from the y-intercept of the same graph and from 
equation (11) i'-l 
Place this data into Table 2* If you use a computer based 
spreadsheet, you may substitute a printout of your data for this 
table* Be sure to indicate that a spreadsheet is attached 
containing your data. 


:'y .' '-y- . • 

EQUILIBRIUM CONSTANT Page 9 



.. 
12 0 
ROOM TEMPERATURE: 
BEA 
KER 4 T 
TT 
T 
Aj[Fe?*3 
TABLE 2 
TEMPERATURE- *C 
INITIAL 
[SCN"] 
4-4 
[Faa+]
I 
•f. ^^U^B^^^^^^h 
EQUILIBRIUM 
[SCtT] ÍFeSCN3+J 
T 1 
"f-
3I 
4I 
TI6I 
}: ' 
F 
i 
I 4 . :••: 
. 
• 
7 I 
a i J— t . 
COLD TEMPERATURE:
1Ii r r 
T 
2I 
~~i—r 
3 I I 
i i; 
4| 
I I 
HOT TEHPERATUEE;
T^TT—T 51 
Hi i 
TEMPERATURET 
T 
± 
TEHPERATURET 
°C 
h 
T" 
l 
I 
. • • • • . 
J± 
7 I I 
J L 
T 
-L 
•
'• • 
y 
' • . 
: 
• • 
• : " — 
• y. 
". . 
• • • 
• • •
, •
• 
. 
• • 
. • 
• • . 
: • "• ' 
• •'. ; ; • •
• 
•••• : 
• • • 
.• • 
• 
• ' • y m . 

EQUILIBRIUM CONSTANT Pag e 10 



• .. 
' • •.-. 


• 


-, • . y ;• 

4. CONCLUSION: Compare your ÙH, ÙG and *S with the following data 
from the CRC handbook (| 296 K), 
• , .. •• 
•"•• . . : • •
7 vW»]y¿.,r £]••'., 

- aol KJ 
i, mol * \-aol •* 
SCM'iaq) 76,44 92,68 .144 y • 

. • 

•

Fe^{aq} -4a, 53 --4:60, -.315 

•'••• .
FeSCN2+(aq) 23,43 71,13- -.130 

. . . . . . .• • .• .. • 

'• " ' '.'•'•.'"•• • •• ' '-. • V • • • ' • ' . . 

• 7 • y -.-... •=••' • • • 
7 

ùH AG

ncnrx« A*rx» 

• -' : •.'-.'y. .7 * • :. "• . '•' " • ':'•••'•'. ' • y y -y 

. . . - •. 

y •' ' • . •• . • ' . • • • 7 • ' . • " • 

Experimental ....., • -..-.. 

ror 

• 

• • ' : •-. 


..... • • .'• • . . . ' 


\ Error • • 

Note: These theoretical values will probably not match your 
experimental values very closely. 


•_ . 
. y 
y ..•••• 
"• . • • • " • • . • ; 
• • " : . ' ... : . 
• .• 
. • • ' • • 
.
• 
• .... . .
y . . 
• • 
y 
• • . • . 
...• . . 
. . 
' • ' • . . . . • . . ' • ' y y • 
• . . • • • . . . . 
• • • • ' .'• • • • • " • • • • • 
... 
i •. 
•• . ' . 
•y 
• 
• 
• • 
• ' •' 
'.. • • • . ; • • • ' . 
• . 
• . ' : • . : • •' • '• •..• • •• • • • ' • 
• " 
• . -. -. . . . . • . . . • • •< , . -. -• 
yyyy. : .. yy..... y. y y. : y ,. 'yyy'yy'y' .:'•;••>..-." 
-yy 
• . : 

... ' • ' .' y' 

. 
. • .".

... .•-... ' 

• •. 

' -• : • • * • 

EQUILIBRIUM CONSTANT Page 11 



122 


INSTRUCTIONS FOR THE OPERATION OF THE SPEC 20 


1. Be sure that the instrument is plugged in and turned on, 
¿. If the SPEC 20 is not on, turn it on by rotating the power knob 


on the left front face of thé Spec 20, clockwise. It will then take 


10 minutes to warm up. 


3. Set the correct wavelength using the wavelength adjustment knob 
{on the upper:face, right hand side of the instrument). The dial is 
calibrated in nanometers. Remember to interpolate between two scale 
markings to set the right wavelength. Ask your instructor if your 
not sure how to set it correctly. 
•


4. Make sure that the sample compartment is empty and the cover 
closed. Adjust the power knob, still keeping the pover on, so that 
the meter reads 0% transmittance. 


NOTE: To avoid parallax error in your reading of the needle, 
move your head so that the needle and its reflection in the mirror, 
are superimposed* 


•. 
: ' . : , : • " • . " ' ' ' : 7 • • : • • • y •" : .'.. .. '. : ' •: 

5. Now fill the cuvette to the bottom of the frosted circle with 
• ••.. 
de-ionised water. Wipe the outside of the cuvette with a tissue and 
make sure that it is free of fingerprints. Place the cuvette in the 
compartment making sure that the line on the cuvette is aligned with 


• 


the line on the front of the sample compartment. Close the lid over 
the sample compartment. 


6* Use the knob on the right front hand face of the spec 20 to: 
adjust the needle so that it reads 100* transmittance* 


* •


7. Repeat the above steps 4-6, until the Spec 20 is calibrated. 
E. Hake sure that the outside of the cuvette is clean and dry. The 
cuvette itself should be rinsed with a portion of the sample you 
wish to measure the percent transmittance of. 
9. Fill the rinsed cuvette to the bottom of the frosted circle with 
the sample you wish to measure the percent transmittance of. Dry 
the exterior of the cuvette. 
10. Place the cuvette in the sample compartment as before with the 
line on the cuvette aligned with the line on the front of the sample 
compartment, 
11. Close the lid on the sample compartment and read the percent 
transmittance. 


' ' -' . • : • . ' • • . • • " • • • , • . : • • • • • . • • ' • • ' • • ' 

EQUILIBRIUM CONSTANT Page 12 


• . " • -• • • 

ACID BASE TITRATION 
chemistry 131 


INTRODUCTION 


Analysis of an unknown acid concentration is often determined 
by "titrating11 the unknown acid with a known amount of base- In 
this experiment you will perform a titration of a strong acid (HCl) 
with a strong base {NaOH) and you will also perform a titration of a 
weak acid (acetic acid) with a strong base (NaOH)* 

OBJECTIVE 


i. Understand the concept of the titration of weak and strong acids 
with a strong base. 
.-...• •• -;•"• y . . • . y . . ,:•••-. • 


:2V Introduce the laboratory technique cf titration* 


THEORY 


Refer to textbook for the complete theory of acid base 
titrations (Section 15.2; Chemistry. Hasterton and Hurley). The 
following is a brief review of the theory. 


The net ionic equation for the reaction between a strong acid 
and a strong base is as follows; 


B+(aa.) + OK-(aq) ->H20 (1> 


The equilibrium constant for this reaction is approximately 10L4 ¿ so 
that for all practical purposes this reaction goes to completion, 


A typical strong acid/strong base titration curve is shown 
below* The volume of titrant is plotted versus the pH of the 
solution. The equivalence point on the titration curve represents 
the point at which the the moles of base (titrant) added is 
equivalent to the number of moles of acid in your unknown solution* 
Notice that the pH is 7 at the equivalence point for a strong 
acid/strong base titration. 


Acid Base Titration Page 1 



y 124: 

Tlti-*tlAH *t Su, 00 mL 
of 1,000 M MCI VIïta 


1.000 K HaOH. 
i* 

, 

y . >y.* 
.^F . 

12 

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.


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.: 

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10 L 

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a 
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Polo t-...: 
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i •'•'. ' ^ p 

-;•" > ".-'•.. : •• ' • ' : • -' 

70..-•; . $&-.- .•>(*"•• :.../s/-: . y 10? 

Since we will not know the pH of the solution in our strong 
acid/strong base titration, how can we determine the equivalence 
point? The answer is quite simple. We vill use an indicator. 


An indicator is just a very weak acid that changes color at 
different pH ranges. The indicator that we will use for the strong 
acid/strong base titration is phenolphthalein. Phenolphthalein 
changes from clear to red around pH 9. You may ask why are we using 
an indicator that changes color at pH 9 when the equivalence point 
occurs at pH 7. Notice that the titration curve is very steep both 
before and after the equivalence point. In fact, the line is almost 
vertical between pH 4 and pH 10, Because the line is vertical 
between pH 4 and pH 10, there is very little change in the volume of 
NaOH added in this range. Thus, any indicator that changes coîbr;iii 
the vertical portion of the strong acid/strong base titration curve 
is acceptable. ' y-~'y. 

Base Titration Page 2 



125 
. 


• • . 


. 
.. 

A weak acid/strong base titration is quite different than a 
strong acid/strong base titration curve. Let's consider the 
titration of acetic acid with NaOH* 


CHjC03H(ag) + OH"(aq) ^CH3COa"(aq) + H20 (2) 


The equilibrium constant for this reaction is approximately 109 , so 
it too goes to completion. Let's consider how pH changes during the 
titration as shown in the following titration curve. 

• 

••••• y. 

... 

•. • 

Titration of 50.00 nL 
of the vctk acid-HCJH^OJ [\ 

: ' -• • • 

•

(1.OO0 H, With 1.000 H »*>0H. ... . • 

• 

• 

-

14 
: ' -..' " 
• ' • ' 
•


.... . 12 


• "."-.•• 
•


; f 

1 tn«|

lO 


Pain t 

• palC • 7. . ... 
• ••. 


. • .. • . 


. . • .


'-' 1 
pH -y : •: .-J .' •
•• •7-7 

• : .. . 
y-• -.' • 

•


f •7 y. '. 

2* 50 73 100 

il HaOH M44vA 

7J 

-. ••.


yyyyyyy 

Noticë that the pH-, starts at about 2.4 and rises rapidly,: 
About halfway to the equivalence point the pH changes very slowly. 
In this region you have a buffer system—acetic acid and the acetate 
ions produced by that addition of OH" (as shown in equation 2), At 
the equivalence point the pH is greater than 7 because in this 
region there is no acetic acid left in solution* but only acetate 
ion, which is a weak base* Notice also that the curve is not nearly 
as vertical around the equivalence point as it was with a strong 
acid/stropgbase titration. Therefore, we must carefully choose an 
indicator which changes color very close to the equivalence point. 

Acid Base Titration Page 3 



.'• ' • -. :. • •• ' -•• . 

1?6 


Before we start any titrations, however, a chemist needs to 
make sure that the concentration of the titrant is known exactly. 
The concentration of the sodium hydroxide solution may change with 
time. If.CO; Is absorbed by the NaOH solution, this would result in 
formation of carbonic acid which would neutralize some of the NaOH. 
To determine the exact concentration of the NaOH solution, 
therefore, we must "standardize" It against a known concentration of 
acid. The acid that we will use to standardize the NaOH is 
potassium hydrogen phthalate (KHP)* It reacts with the NaOH 
according to the following reaction. 

(3) 
The procedure for standardizing the NaOB is to measure out a certain 
amount of KHP* dissolve it in water and titrate with NaOH to the 
equivalence point* Using the following relationship. 


Molarity^*r„ * Volume^E„ = moles of KHP {4} 

you can determine the exact molarity of the NaOH solution. 


EROCBDPKE 


(Notai Your instructor will dirsct you on th* proper titration 


technique prior to performing the lab. If you bava any questions 


pisase askI) 


1. This lab will be conducted in pairs and written up independently. 
All collaboration must be documented2- 
Weigh 2 samples of approximately 0.4 grams of solid potassium 
hydrogen phthalate (KHP) and dissolve each in 50 mL of deionized water 
in a 200 or 250 mL Erlenmeyer flask. Add 5-6 drops of a 0.1 % 
phenophthalein indicator solution to each flask and then titrate the 
solution to a faint pink endpoint. Use these data to standardize the 
NaOH solution. 

3. Once the NaOH solution is standardized, determine the molarity 
of an unknown hydrochloric acid solution by titrating approximately 
10 mL of the unknown acid with the standardized NaOH solution, 
perform two replicates using a phenolphthalein endpoint. 
4. Determine the molarity of an unknown acetic acid solution toy 
titrating approximately 10 mL of the unknown with NaÓH to a 
Acid Base Titration Page 4 



• •' • • .'•' .' • ••• :.-. • -y 
..•:•• 

.. • 
• 


phenolphthalein endpoint. Perform two replicates with the 
phenolphthalein indicator. 


. y 7 • •••.'.'.-J . '.. ' •• •. -.:•-'.••'.'•.. • 


5. Next titrate the same unknown acetic acid solution using methyl 
orange as an indicator. Perform the titration only once. 
• • • • . ' • • • ' • ' " : ' ;• • y 


6. Your instructor will direct the entire class to perform a pK 
titration of the unknown acetic acid solution using NaOH and no 
indicator. The data from this titration will be supplied to the 
whole class. The procedure for the pH titration is as follows; 
• 


a. Measure the pH of the solution before adding titrant. 
b- Add a small volume of titrant. Stop, record volume added, 
measure and record the pH» 


c* Repeat step "b" until within about 2.0 mL of the equivalence 
point or use volume increments that give approximately 0.2-0.3 pH 
unit changes. 


d. In the vicinity of the equivalence point, many data points are 
needed, so take pH readings after every 0.10 mL titrant is added. 
e. After the equivalence point, continue taking readings (at 1-2 mL 
increments) until at least 5 mL beyond the equivalence point. 
• . . •;• 

• 

•• y 

-• • 

: -. •
• • 
• . • • , • 
" , ' • 
' . •
:•••.:
' • " 
• 
•_. •= _ :. • y . 
.y 

Y • • •• 


. • •


Acid Base Titration Page 5 



12fí 


tan, 

y . 
• 


• 


• 


• 


1 \ 


\ • • ' . '* 

Acid Base Titration Page 6 



OONCLOBIOWa jghag all calculation» 


1. What is the average molarity of the NaOH solution? 
2. What is the average molarity of the unknown HCl solution? 
. • 


• " . * -. 
3. What is the average molarity of the unknown acetic acid solution 
using phenolphthalein as an indicator? 
. •• 


Acid Base Titration Page 7 



4. What is the molarity of the unknown acetic acid solution using 
the methyl orange indicator? Is it different than the titration 
using the phenolphthalein endpoint? If so, why? 
• .. • . 
• 


• . • 


5. Plot the pH titration data of the unknown acetic acid with NaOH 
using Quattro. (Attach the graph to this lab alsoi> 
a. What is the equivalence point in mL NaOH? (Hark it on the graph 
also.) • 
• • 7 


. . .


rjjju--:; What is the pH half way to the equivalence point? 


.•• . 


. . 


ç. What is the pH of the solution at the equivalence point? 


d. What is the effective pH range of phenolphthalein as an 
indicator? (Highlight this range on the graph.) 
e. What is the effective pH range of methyl orange as an indicator? 
(Highlight this range on the graph also.) 
Acid Base Titration Page 8 



.' . : y 

•• • •:• 

• 


.


PRELAB EXERClflg 

1. Draw a typical weak acid strong base titration curve. Label the 
axes and equivalence point. 
7 '7. 7'.•••.•.' • 7y. y .". •••. 7 . •. .'.A • • y • " -:7 . 


-"•ï .A " , 


7 


: •'• ':' 


• 

• 

• . • 

•

2. Define equivalence point. 
• 

7 •'••:•.'••• •'••• y' yy. 

7 

.

• ••-'•••• "•'••:"" 

• • y --• y -•• •• 

3a. What i s the pH of a 1.00L 0.1 H aceti c acid solution? 

• • • '• : • 


" --: • ' • ' . . • • . ' • • • ' • • • ' . -. 


" • • • " " • • • • • . . • : . • • • •-. " • • : . • 

. • • 


•'.••' • • . y .••7 • 


•.• . •••..-•• 
•'.•.-:• v. :.'.'• 


. . '. •. "'.... 


• •.: '•.' 


b. What is the pH of the above solution after the addition of 500 
mL of o.l K NaOH solution? 
y .y y ' . •' . • 
' -' • . . " 


y .. . y-"yy-y


•• • -••-• 

y •"• .-•••. . yy -y 

•: .: 
' - • • 


• . y ' 
• ':• • 
C. What is the pH of the above solution after it has been titrated 
with o.l H NaOH to its equivalence point? 
1 


. 7. v 


• • . . • • 


... 


. • 


d * : what i s the pH e f the above solu t ioh a f te r ^ddlt i wti o f LOI L 
of . 1 H HaOH? 

Acid Base Titration Page 9 



• ••'.. 

•


y • • • • • . .-. •• 

PHOTOGRAPHY 

•. 
Chemistr y 13 1 

• y • • 
:;':

EfrtSpduqtioft •yy•••' : " . • • 

•: • : . 
The following appeared in the Gazette de France on 6 January 
1839: "An important discovery by our famous painter of the Diorama, 


H. Daguerre. The discovery partakes of the prodigious. it upsets 
all scientific theories of light and optics, and will revolutionize 
the art of drawing." The author was describing a paper presented by 
Daguerre at the National Academy of science on how light was used to 
"make" pictures. This was the start of photography. 
Photography is a true blend of art and science. In its 
beginning a photographer was more scientist and experimenter than 
artist. He had to prepare his film and development procedures from 
scratch. There were no off-the-shelf film or 2 4 hour development 
labs. Every part of photography was done by the photographer. 
Today modern science and technology provide us with very 
sophisticated cameras and film capable of recording images and 
detail as never before. But high technology is only a beginning for 
good photographs. Juxtaposition of subject elements, perspective, 
light and shadows—these are the things the photographer must apply 
artistic talent in order to create superior photographs. Today we 
will not be concerned about photo quality- Instead we will concern 
ourselves primarily about the chemistry of photography. 


. 7. 


Objective 


1. To understand the chemical principles involved in photographic 
paper development. 
2, Reinforce the principles of acid base reactions, precipitation 
reactions, oxidation reduction reactions and equilibrium. 

Theory1 


The basic principle of photography is that light is focused on 
photographic paper containing silver halides. This light forms an 
invisible image called a "latent image". The latent image process 
is just a very simple oxidation reduction reaction. First the 
halide (CI", Br' or I ] absorbs a photon of light (h*) and releases 
an electron in an oxidation reaction (1). The electron released 
from the halide reduces the silver ions to form metallic silver (2J. 
This metallic silver is the basis of the latent image. 


17 Bunting, Soger i The Chemistry af Photography'• 1987 í (Portions of 
the theory copied with permission of Professor Bunting.) 

Photography 1 



•'::. • : 
• 


light (hv> +-Br" > Br +•e <D 


Ag + e' -> Ag(s) (2) 

This latent image is magnified during the development process 
via a number of chemical reactions and the result is a negative 
photograph. The magnified image on the photo paper appears opposite 
that of the real object. That is the light parts of the object 
appear dark and the dark parts of the object appear light. The 
reason for this is that the reduced silver on the photo paper is 
formed in very small, evenly dispersad particles. When light hits 
the reduced silver, the small particles diffract the light and make 
it appear dark. The unexposed portion of the photo paper 
corresponds to the dark parts of the object. However, since there 
is no unreduced silver particles on the unexposed photo paper this 
region appears light. The chemistry of the negative development 
process is described below. 


". 


•


Photographic Paper 


The essence of photographic paper is a layer of silver halide 


(AgCl, AgBr, and Agi) on a clear plastic support. The silver 


halide, principally bromide, is a fine powder and must be somehow 


fixed to the surface of the plastic base. The material used to bind 


the silver halide to this support must be transparent to allow light 


to reach the silver halide grains; it must be fairly rigid to 


prevent particle movement which would blur the image; yet it must 


not be brittle so as to crack wheh the film is flexed. Finallyf and 


of utmost importance, it must allow water and solutions to penetrate 


it so that the reacting chemicals may reach the silver halide in the 
"processing steps. y'"". . 
.. 

The material which meets all of these requirements is the same 
as that used in fruit flavored "jello" desserts gelatin, Gelatin 
is a very complex and indefinite molecular structure. it is a 
protein material and is made up of amino acids. These molecules are 
typically made up of long chains of 3Q0 to 900 atoms (where n = 
100-300) as shown in the following figure. 


0

\y 


1 / í 

'OH


CH


K -•r.Çii <3> 

/ yCIim I

1 i 

H A-R-i-H" 

• '.'•• i 
R 


•A . 
. 


: 


• . 
. • 


• 


Photography 2 


The symbol K is shown to represent some group of carbon atoms 
Of unspecified length or structure. In proteins these R groups may 
contain occasional atoms of nitrogen, oxygen, sulfur or phosphorous. 
The composition of these R groups have a profound effect on the 
properties of the film. For this lab we will not discuss any of 
those effects. 


Now let's consider the preparation of the silver halide, which 
must be suspended in the gelatin. One way to form it would be to 
react silver with bromine to form silver bromide according to the 
following. 


2Ag(s) + Brï(t) > 2AgBr(s) (4) 


If we drop a chunk of silver into bromine liquid we will only 
form AgBr on the "outside** of the chunk. For the purpose of film, 
this is not an acceptable way to incorporate AgBr in the gelatin. 
Recalling your solubility rules, a better procedure would be to take 
absolution of AgNO? and mix some KBr solution with it. The reaction 
is as follows 


AgN03(aq) + KBr(aq) -> AgBr (a) + KNo3(aq) (5) 


Commercially the procedure for making film is to take AgNOa and 
mix it with various halide salts (KBr, KI, KCl) in a liquid gelatin 
at a temperature of 50-80" C for 1-2 hours. (The KI and KCl is used 
to vary film sensitivity and grain size.) The solution is quick 
chilled, solidified, shredded and washed to remove the KBr, KI, and 
KCl. Finally it is reme1ted and spread on a film or photographic 
paper in a very thin and precisely uniform layer. 


optics 


•A, . y A A 
In a typical 35 mm camera, the image if focused through a lens 
shown in the following diagram. 


Source 


ObJ eet 


35am Camera Box 

Photography 3 



Cantera lens normally contain a diaphragm—a device to próvido 
circular hole of variable diameter. The diaphragm is known as *-tv_i 
aperture and its principle function is to control the intensity of 
light which passes through the lens to the film. Adjusting the 
"f-stop" or Hf-number** on a camera lens is just making a variation 
in the aperture diameter. The f-number is equal to 


f-number = focal length/aperture diameter (G) 


As you can seer for a camera with a fixed focal length, thr 
smaller the f-number the larger the aperture diameter. The other 
adjustable settings on a camera is the shutter speed. The faster 
the shutter speed, the less light reaches the film and the slower 
the shutter speed, the more light reaches the film. 


With our pin-hole cameras, the optics are very_simple. The 
camera has no lens, but has a pin hole which allows light to enter 
irt and expose the filni. 


Photograph!c 
papar *o 
Source 
Pin hole Object 
t Cañera box 
(being 
photographed) 

To vary the amount of light that reaches the photographic paper 
there are only two controls: 


1. Size of the hole. 
2. Length of time film is exposed. 
The specifics of hole size and film exposure time is discussed 
in procedure section. 


Ow e loping the Image 


In order to produce a visible image on an exposed film, 
additional silver must be deposited in the vicinity of each of the 
small silver specks that make up the latent image. This is brought 
about by the development process, 


Photographic developers contain chemicals that are reducing 
agents. These reducing agents readily give up electrons to reduce 
the silver ions in the silver halide to metallic silver. 


V , y •" 
Photography 4 
• 


•*. What sort of materials can serve as developing agents? A 
logical guess would be some other metal more active than silver, A 
metal that is "more active" is one that can give up its electrons 
more readily than silver (i.e., has a higher oxidation potential)-
Hercury is an example. It is slightly more active than silver, and 
so can react with a silver halide as the following equation shows: 


Mg(l) 4- AgBr(s) > HgBr(s) +- Ag(s) (7) 


Bromide is just a spectator ion and it remains unchanged 
through the process. 


Mercury metal once was used as a developing agent in just this 
manner. Its effects were discovered quite accidentally by Daguerre 
in 1835. Daguerre prepared silver iodide emulsions, exposed them in 
a camera and then stored the exposed plates near some spilled 
mercury from a broken thermometer. The result was that the mercury 
developed his exposed film (containing a latent image) into a 
visible imagé. These photographs later became known as 


* dague r retypes. *? 
The reducing agents in developers in use today are all organic 
compounds soluble in water. In solution the molecules have the 
necessary mobility to get in contact with the insoluble silver 
halide in order to reduce the silver in the film emulsion. The most 
widely used reducing agent in photography today is hydroquinone 
CeH4(OH)z. Hydroquinone is a weak acid that dissociates according 
to the following reaction: 


OH 


+ 
U2Q £ H30+ 
(-o + (8) 
(aq) 
Since this equilibrium lies far to the left we must add a 
chemical in order to "activate1* hydroquinone. The chemical which 
activates hydraquinone is hydroxide. When a base is added the 
hydroxide reacts with the two protons on hydraquinone to form the 


. -. . • : • • • * -. 

••. • • '.•' 


.••• • J»



dianion of hydroquinone according to the following reaction: 
0;


OH


+ 
20H-> + 2H30 
<9) 
OH 
(hydroquinone) (hydroquinone dianicm) 


In the presence of Ag+ , |l acts as a reducing agent and 


becomes oxidized according to the following half reactions; 


'+•'••la (Oxidation half reaction) (10) 

(hydroquinone dianion) (quinone) 


Ag+(aq) + e~ > Ag(s) (reduction half 
reaction) (11) 


.y-.y-


i-''.y 

L-'•• V 

• y. . v . \ ¡
7 ."" ^ 


7v-.'i!'. . : ••'. 


Pho 6 



13fi 

Thus the overall reaction is; 


2Ag*(aq> > + 2Ag[s> (12 ) 

Uq ) 

The pH range of all developer solutions is always basic. 
Normally the pH range is 10-12 using Na3C03 as the base. 


Up to this point we have ignored an obvious question regarding 
the development reaction. Why is it that the reduction of silver in 
the emulsion occurs only in the vicinity of the silver particles of 
the latent image? Why aren't all the silver ions reduced? The 
answer to this question is that the rate of reaction of the silver 
near the latent image is much greater than the rate of reaction of 
silver not near the latent image. The developer can and will reduce 
all the silver ions in the fil» and if development is extended for 
too long a time the entire emulsion will turn black. The reason we 
can use the development reaction to produce an image is that 
reduction occurs faster near the silver particles. So we can 
develop a film or print until the image sufficiently darkens, but 
stop the development before the slower reacting silver halide is 
reduced* The silver metal of the latent image acts as a catalyst 
for reduction of the silver ions with which it is in contact. 
Chemical development would not be possible if it were not for this 
catalyst. 


Stop Beth 


When the development process is completed - that is when 
sufficient silver has been reduced to give the desired image density 


- the film is placed into a stop bath. The purpose of the stop bath 
is to prevent any further reduction of silver ions. Since the 
developer solution is only activated at pH ranges above 7, then one 
way to stop the development process is to "wash" the film in an 
acidic bath and in so doing shift the equilibrium of the developer 
solution from quinone to hydroquinone. 
Photography 7 



By adding acid, the H+ will react with the excess OH" in the 
developer to form water. Since there is no hydroxide to "activate" 
the hydroquinone, development can not occur. 


Fixer 


After the reaction is stopped, we are left with a silver image 
superimposed on a background of pale yellow silver halide. This 
silver halide which was not reduced in development must be removed 
by the fixing process. If not, a print would ultimately darken due 
to gradual reduction of more silver from exposure to light. Fixing 
is a process by which the remaihing insoluble silver halide is 
converted to a soluble material which can be washed out of the 
emulsion. A great many substances, both negative ions and neutral 
molecules, have since been found which will complex silver ions. 
Ammonia, for example, is a molecule that can dissolve silver 
chloride by bonding to it to produce a complex positive ion, 


AgCl(a) + 2NH3{aq) > Ag(NHS)3 
+(aq) + Cl(aq) (13) 


The materials commonly used in photographic fixing solutions 
today are salts containing the thiosulfate ion S20,2 -( The fixer 
used in this lab is ammonium thiosulfate or (NHOjSîOa. The fixing 
action of thiosulfate on silver bromide is as follows: 


AgBr(a) + 2s203 
3^(aq) > Ag(S203)Î*~(aq) + Br" (14) 


Thiosulfate dissolves the silver bromide and the ammonium ion can 


dissolve the silver chloride. 


AgCl(s) + 2HHJ« + (aq) > Ag(NH3) 2 * (aq) + 2H+(aq) + Cl-(aq) (IS) 

Developing, stopping, and fixing are the three sequential steps 
that must be performed in the standard processing of all black and 
white photographic materials, rollowing these three steps it is 
necessary to thoroughly rinse a film or print before drying. If any 
thiosulfate is left in the emulsion the image will not be permanent. 
Excess thiosulfate in the emulsion will turn the photo yellow and 
eventually cause the image to fade, 


reversal Procesalng 


In order to obtain a positive image, the reversal process *s 
used. Recall that the latent image is made of reduced silver and 
appears dark, even though that it represents the "light" part of the 
object. The unreduced silver on the photo paper appears light and 
represents the "dark" part of the object. in order to reverse this, 
the latent image is bleached and washed out, and the unreduced 
silver is exposed to light. This produces a photo in which light 
parts appear light and dark parts appear dark. 


The procedure for this is quite simple. The photo paper is placed 
into the developer and left there until the latent image is formed. 
The paper is then placed into a bleach bath of potassium dichromate 
to oxidize the silver from the latent image. The photographic paper 
is washed to remove the bleach and the dissolved (oxidized) silver, 
and then exposed to light so the remaining silver halide forms the 


positive latent Image.
rinsed and dried. 
The photo paper is then developed, washed, 
y7 
Photography 8 


1*0 


Procedure 


1. Build a pin hole camera. If you have any questions about the 
construction, see your instructor. Cut a 1 cm x 1 cm hole in the 
side of the box. On the inside of the box, tape a piece of aluminum 
foil over the hole. With a paper clip punch a hole (the smallest 
hole possible) in the foil. On the outside of the box, cover the 
hole with electrical tape, under safe light conditions, insert the 
photographic paper into your box and tape it against the inside of 
the box directly opposite the pin hole. Tape the box shut so that 
it is light tight. 
2. Expose the paper to your subject for approximately 1 second (on 
a sunny day) and for 3-5 seconds (on a cloudy day) . The sun ïtmst be 
to your back to avoid overexposing the paper. 
3. Under safelight conditions, develop your photographic paper. For 
negative processing, place the paper into the developer. Gently 
agitate the paper while it is submerged in the developer. When your 
image begins to appear remove the paper* Some development will occur 
after the paper has been removed from the developer. When the image 
is developed, place the paper in the acid stop bath for about 1 
minute. Next, place the paper in the ammonium thiosulfate (NH*)jS303 
fixer for about a minute. Rinse the paper completely with tap water 
and allow to dry. 
4* For reversal processing, place the paper in the developer. 
Gently agitate the paper while it is submerged in the developer and 
slightly overdevelop your image. Place the paper in the KjCr207 
bleaching solution. Your image will disappear as all of the 
metallic Ag is dissolved. Rinse your paper and turn on the lights. 
Place the paper back into the developer and the positive image will 
appear* Rinse the paper completely and allow to dry. 


Photography 9 



••"y • lAl 

concluilom 


1. What is the pH of a O.l H hydroquinone solution (Ka = 4.S x 
lO"11)? 
2. what "activates" hydroquinone to be a developing solution? 
3. Typically hydroquinone is placed in Na2C03 buffer solution. 
What is the pH of a buffer solution containing 0.15 H Ha2C05 and 
0.10 M H2C03? If the effective pH range for hydroquinone 
development solution is 11.0 or greater, would the above buffer be 
adequate? 
4. For Br" to be oxidized to Br requires a photon with an energy of 
at least 2.58 x 101* J. What wavelength of electromagnetic 
radiation does this correspond? If the safelight emits in the red 
region at 750 nm, would the safelight effect the photo paper? 
5* During negative processing, why must photo paper be placed in 
the "fixer"? 


Photography 10 

j_^^^^^^^^^^^^^^b^ 


,y 142 


6a. Given that [Ag+] = 0.0025 H in a 4 liters of spent fixing 
solution, how many grams of silver is in the solution? 


fib. How many grams of NaCl should be added to the solution to 
reduce th* [Ag+] to 1*0 x 10 * H? 


6c. wow much silver was recovered in this process? 


•"••(••• :"••• . : •" . 
7. Explain why the developer reduces silver around the latent image 
and not In the unexposed areas. 


y. -A-' 
• "•••. 
Photography 11 


•. -.-^ ' 

yy--.y-y::'y .7:7 yyyy. y y. ^ 

Sa. Potassium dichromate is used to bleach out (oxidise) the silver 


from the latent image in reversal processing. The following 


represents the overall unbalanced redox equation. Balance the 


equation. 
Ag(s) + K2Cr?07(aq) + H35D4(aq) > Ag?S04(aq) + Cr?(504)3 (aq) + K2SOil(aq) 


8b. When the bleaching solution becomes "weak" a small amount of 
sulfuric acid is added, why? 


what is the result of overexposed photo paper? 


10. Attach your photo to the lab* What could you have done during 
: n-:•:••:


this lab to improve your photo? 


••. • . .• " i 

y'-.y 

y .-v 


Photography 12 



••: Prelabératory Exercise 


I. How does light expose black and white photo paper? (Include 
.chemical equations.) 
2. : Why does reduced s i 1 ver appear bláól/t on photo papet ? 
Z - If Br' was; 6X id i zed to Br 2, i nstead of just..; Br y what effect might: 
that have on the formation bf latent imagé. 

4* ; Why7Is geiatin used as: the matrix for si1ver ha1ide? • 


S, • •' • Why i s., ace tic ac id used in the stop bath? 


phó 13 


y:y y y y ••..?;.-'•:'.:.: • 

14$ 


QUALITATIVE ANALYSIS 
IDENTIFICATION OF CHEMICAL COMPOUNDS 
Chemistry 131 


INTRODUCTION 


Qualitative analysis is the process of identifying the content 
of a sample with regard to the chemical species present. No 
assessment of the amount of the chemicals present is required. 


This lab is the culmination of your chemistry experience for the 
semester. You have learned about chemical reactions and the behavior 
of chemical species, the properties of solutions, complex ions, and 
solids. Now you must integrate all these concepts, selectively using 
each tool of knowledge to analyze qualitatively unknown chemical 
samples* 


OBJECTIVES 


Your goal in this lab is to successfully identify the cation and 
anion in five separate unknowns. You will do this in two parts. The 
first part of the lab consists of creating a "reaction matrix" by 
mixing known chemicals and observing and recording the product of any 
reaction. In the second part, you will identify the unknowns. You 
will be required to perform the second part without any collaboration 
from any source except your written notes from your reaction matrix, 
your knowledge of chemistry, and this laboratory handout. 


TffEPRY 


Ions in solution (whether cation or anion) have distinctive 
properties which allow them to be identified by controlling the 
experimental conditions. Certain ions form precipitates based on the 
solubility product constant, Ksp, of any resulting species formed. 
Others may undergo oxidation-reduction reactions that produce gases 


(visible as bubbles in solution) or changes in color of the solution. 
Some ions form complex ions with distinctive colors. Finally, some 
ions react with the water to result in either basic or acidic 
solutions. You have observed all these chemical characteristics in 
previous labs and as demonstrations during this semester. 


Salts are compounds that are made from cations and anions. In 
solution, these ions dissociate into individual species that possess 
unique characteristics* First* let's practice with the ions produced 
by salts. Below is a table of six salts and the resulting ions 
produced. Complete the table by filling in the blanks: 


• • . 


• 


• 


• y 
7 . 

— 


Salt Cations 


Anions 


AgCl 


NHjHO*, 


Ba2* 


•


FeS04 


B*a* 


.»«•.


.. .. •-


ZnS 7 • . 


ftemember that an ion may have various oxidation states, but the 
overall charge on the neutral salt must equal zero* 

A proven way of approaching this problem is to first identify 
the cation in an unknown. Four "stock" reagents will be provided 
that give distinctive information about the cation* They are 
ammonium hydroxide, hydrochloric acid, sodium hydroxide, and 
potassium permanganate. Once the cation is positively identified, 
the anion is identified by performing similar tests* You will also 
use twelve different salts that contain all the cations you will see 
in the unknowns. Reacting the salts together provides a data base of 
observation to correctly identify the cation* It is crucial that you 
récord detailed observations of the salt reactions. 


An obvious fact that is often overlooked is that the sample 
compound must be water soluble. If not, you would observe a solid in 
your unknown. Using the solubility rules narrows your investigation 
by reducing the range of possible compounds. 


staple Analysis 


Let's work through a sample qualitative analysis. Our unknown 
is a clear, colorless liquid. Mixing with a drop of hydrochloric 
acid gives a white precipitate* as does mixing with a drop of sodium 
hydroxide and ammonium hydroxide. Potassium permanganate produces no 
observable change. 


A positive reaction with hydrochloric acid may indicate the 
presence of a base; but our unknown also reacts with the bases sodium 
hydroxide and ammonium hydroxide. Since a white precipitate forms, 
let's refer to the solubility rules for help* 


Qualitative Analysis Page 2 


"• . 



14? 


RULE li All nitrates (W03 ) are soluble. 


Our unknown is soluble - it may be a nitrate. However, none of 
the precipitates formed are nitrates. 


RUL E 2: All chloride»* bromides* and iodides are soluble EXCEPT 
those of silver, Mercury(I), and lead, copper iodide is also 
insoluble. 


Our unknown could be one of the soluble halides. But, the 
addition of chloride (from the HCl) resulted in formation of a 
precipitate. Thus, our cation is either silver, mercury, or lead. 


RUL E 31 All sulfates (so4*) are soluble except lead, strontium, 
barium, calcium, and mercury(II). Silver sulfate is sparingly 
soluble. 


Our unknown could again be a soluble sulfate. However, since we 
have already limited our choice of cations to silver, mercury(I) and 
lead, we can rule out lead sulfate and probably silver sulfate. 


RULE 4t All carbonates (0O3 
2"ï, phosphates (PO4*~) J ana chromâtes 
(Cr04 
3~) are insoluble except tbe alXali natals and ammonium. 

Little here except that the cation could be an alkali metal or 
ammonium but we already know it is not. 


RULE 51 All hydroxides (OB') and sulfides (S2~) are insoluble except 
the alkali metals and ammonium. Calcium and barium sulfides 
hydrolyse in water to fo n hydroxides. 


Our unknown formed a solid with hydroxide (NH4OH and NaOH). 
Once again, the cation cannot be an alkali metal or ammonium. 


Based on these tests, we have narrowed our cation to three 
possible ions: lead, silver, and mercury(I). We need more data. 


By reacting this unknown with our four other unknowns, we can 
build a smaller reaction matrix similar to the one you will build in 
Part A* Now the detailed observations of that reaction matrix are 
used* we notice that one of the reactions is similar to that between 
sodium iodide and lead(II) nitrate, forming the yellow precipitate, 
Lead (II) iodide* Since mercury is not included in the reaction 
matrix, we eliminate it as a possibility. Silver does not give a 
yellow precipitate with iodide. We now can assign pur first cation 
as lead(II). 


Having identified onr;cation as Pb2+, we now pursue the anion. 


Qualitative Analysis Page 3 



U 8 


using a similar method of deductive reasoning and elimination, the 
anion is identified as nitrate, NO}-. (It is the only anLon in the 
reaction matrix with which lead is soluble.) 


EXPERIMENTAL 


A* REACTION MATRIX 

In the first part of this lab you'll develop a reaction matrix 
by reacting several salt solutions with four standard reagents (the 
stock solutions): HCl* NaOH, NH4OH, and KMn04. You'll also react 
each salt solution with all the other salt solutions. 


1. Place 12 drops of each stock solution on ths acetate 
reaction surface. You should lay out these solutions to mimic the 
Reaction Matrix on page 7, making the collection of data easier* 
2. Add one drop of each salt solution to a separate drop of 
stock solution* Carefully record your observations on the Reaction 
Matrix. Pay careful attention to color, general appearance, reaction 
times, bubbles, etc* You will use only observations for 
determination of the unknowns in Part B. 
3. Now complete the Reaction Matrix by reacting one drop of 
each salt with a drop of all the other salts* Again, carefully 
observe and record any changes that occur* 
4. For each mixture which reacts* you should be able to 
identify what occurred. For example, mixing AgNOj and NaCl produces 
a white precipitate; you should be able to identify the precipitate 
as AgCl. Silver chloride should form when any soluble silver 
compound and soluble chloride compound are mixed. Identify reaction 
products on the Reaction Matrix wherever possible. 
5. Using a similar methodology, you should be able to identify 
any precipitation reactions that are not predicted, or expected, 
using the solubility rules* Carefully look over your Reaction 
Matrix. Note below any precipitation reaction that the solubility 
rules would not predict: 
V." .-. . 


•" '.-• •• 
Qualitative Analysis Page 4 



6. Now Check the table for any precipitation reaction that you 
would have predicted based on the solubility rules that did not 
occur. Record these below {show the expected precipitate as well as 
reactants): 
.. 


7* After making all observations, clean your reaction surface 
and return the salt solutions to their appropriate trays. Keep the 
four stock solutions. 


B. IDENTIFICATION OF UNKNOWNS 
1. Ask your instructor for your five unknown solutions. 
a. From this moment forward, you are not allowed any 
collaboration with any source. Do not talk to your classmates. Do 
not consult any textbooks or notes except those ycu have written in 
this lab handout. Your instructor will discuss the chemistry of 
reactions of compounds, but will not be able to answer any questions 
concerning identifying the unknowns. 
b. Your unknowns are not necessarily the same salts as you 
used in Part A. However, only the ions in Part A are contained in 
these unknowns, making 9 possible cations (Al5+ , 8a*4", Cu?+ , Fe .3 + 
x^2+ T*V,Z+ Mn= + , Ag % Na+ ) and 5 possible aniens (NO*-, SO*a_
Fe Pb , 


, Cl~, and I~) 


CÓ! 


c Hake initial observation concerning these five unknowns* 
Record the numbers on each microburet which identifies the unknown on 
the Unknown Reaction Matrix (pg 8). Carefully record below the 
physical characteristics of your unknowns (color, presence of 
precipitate, etc.): 

UNKNOWN #!•?•" 


UNKNOWN #2: 


UNKNOWN #3: 


UNKNOWN #4: 


' . ' • • . ':'.•. •"•"••••; ••. ; • • • . . . • • . 

UNKNOWN 15: 


.... 


Qualitative Analysis Page S 


• 



. . • • • • \ • -: . • • • . • • . . -. . • . . • ; 


' . • : • . • • • • . ' • . • • 


• -• • . • • • . • . • • • • . . . . -. . ' 


• . • • . : 
• • . . •"-.' • • . . • 


...'• . '.'.•• .. • 


•-.. • • • . ' • : • -. -• . . : • . • • • . . . • • . • • 

2. React each unknown with the four stock solutions and record 
any observations on the Unknown Reaction Matrix. 
3. React the five unknowns together, again recording your 
observations on the unknown Reaction Matrix. 
4. Using the data you gathered in the first part, your 
knowledge of chemistry, solubility rules, etc., identify the ions 
present in your unknowns. Each unknown contains only one cation and 
one anion. 
a. THINK!I Ask yourself if the combination of ions is 
reasonable. For example, could AgCl be one of your unknown 
solutions? A quick review of the solubility rules will answer this 
one. 


• ' '• • • .•• •. •


b* RELAXi1 Use a process of elimination to narrow your 
possible ions. 


Your analysis and identification of the five unknowns (ten total 
ions) must be completed prior to the end of the laboratory period. 
Your grade for this lab will be based on correctly identifying the 
ten unknown ions, each worth ten points (100 points total for the 
lab). NO OTHER WORK IS REQUIRED. 


7 ' ' • : • . .J 

-. 

'•':•' - •••• • ' 

• 

• 7.Í --..'. 
•. .. • ' ; • • • ' • 

• •. 

Á : •"•. • '•••'• y •• ' '•'•'• ' . 

yy y-.,• y7 ".••'.., ' "; • ..'"• • 

":•..' • 

. •. 

•. . •• ,..-..-. 
• • • • . • • .• .••*••••: 

-_ . -. . y •• -. • : . . . . • • . • : . . . 

• •- •• ' '. ' . 'y : • '••' : -' '•' • • * .-• 

. • • . • • • • . • . . -; ; • ' • •. • • • • • ' . " • • • • • . • • • • . " 

• y 
y. • • '• • y 
. • y •.. 

• . 
y • • •... ••• . 

T Eft 


' ' ". 


. • 

• . 
• •• • • .

Qualitative Analysis Page 6 


= y • ' . . . . ' . ' " • ••'.• ' • ' • 


.. . •' . "'.y -'.'••y'-' •• A ; -: '•'"."•""' •.:'••••' • . •' -y '•"'..'..'' . 


. • :* 


•:-. . 7 ••' ••: ' • ' 
REACTION MATRIX (FART A) 


B *<^1 Ft(Ntt ) PbfMÛ^ j NaC H;Ufl 
*MMO j cufHa r*so 
i. 
Mn(H O^ 
J 
. 
3 it* CO 
, 2 3_ 
Ha I 
HCJ 
î**OH 
tlH4OÜ 
KMnO, 
ALiHÛjîj, 

i 
Fb( N03 )2 . • • • 
. ' ' • 
AfN03 
• 
H *î" 3 
• 
Mic: 
.'.'•• • • •• 
• • ' • . • •' N*I • 
• 
yyyyy •.'•. • 
.--'••.• 
7. , 
" • ' • . • 
NaNO, . . • _ -. • • . 
• . . . 
• 
• • • • 
• . • • 
• • 
•• . . . 
• 

• 

Qualitative Analysis Page 7 


•-••-. ..•'. • '.-• y '. 

. y. 


^ • 



152 


UNKNOWU RÉACTION MATRIX (PART B) 
1 
71 
• •••?•'• • 3 ' ' 7 
.... 4 , 5 
'•2.' 
:.3':-.-y A A' 
• 

4 
S • 
• .A .. 
HCl 
••
: . 
y 
HaOH 
NHflOH 
. 
KMl»Ó4 
UNKNOWN LEÎTEÈ. 
UNKNOWN 
NUMBER 
CATIO N 
ANION,; 

'•' ' 

• '• • • • 
. . ..." 

RESULTS 

• /'. ••-.'"'• ' 
:• ' 

Qua l Ana Pag e 8 

r. • ..• . r 

..." <" " ••. 

•: • • • 
• •. 

<^ 7 


Qualitative Analysis Ísaá--^'. 


IS* 


. • 


CHEMICAL KINETICS 
Chasistry 131 


IHTRODPCTIOH 


We may ask four basic questions ^bout a chemical process. 
First* is the reaction spontaneous? TUernKktynamics allows us to 
determine the conditions under which spontaneity occurs. Second, 
how fast will the reaction proceed? Third, what is the "sequence" 
of forming the products? Chemical kinetics provides answers to how 
fast and gives insight to the reaction mechanism. lastly, how far 
will the reaction go before reaching a state of equilibrium? Study 
of the equilibrium constant, K M , may answer tnis question. YO U 
have studied thermodynamics and equilibria in previous labs, Nov/ 
let'3 take a look, at the kinetics of a reaction. 

Kinetics is important because it allows us to determine the 
rate of the reaction. Experiments have shown that the rate of a 
chemical reaction is dependent on several factors. The four most 
important are* (l) nature of the reactants, (2) concentration of the 
reactants, (3) temperature, and (4) catalysts. 


The purpose of this experiment is to examine the effects of 
these factors on a chemical reaction. In the first experiment, you 
will observe the effects of concentration by reacting various 
concentrations of hydrochloric acid with zinc and aluminum metal. 
To observe the effect of temperature, you will repeat the 
experiments on ice cubes. 


In the second part of the lab, you will study the reaction of 
hydrogen peroxide and iodide in acidic solution: 


Ha02 + 2H+ + 31 2HaO 


m 

you will observe the effects of the change of reactant 
concentrations to determine the rate law of the reaction. From the 
rate law, we will make predictions concerning the mechanism by which 
the products are formed* 


.


REACTION SAT8B 


We can quantitatively express the rate of a reaction in terms 
of rates of change in concentration of the chemical species present 
in the reaction. This change is either written in terms of the 
disappearance of a r«*ctj)nt or the appearance of a product. For 
example, the famous reaction of A and B to form c could have the 
rate expressed three ways. Two of these ways are: 

. . 


. 


Kinetics Page 1 


• 



• 


A + B *.C 

rate ofdisappearance -A[Aj -d{A] 


of A At dt 


rete ofappearance +A[C] +4[C] 

of C * At *
** ~^ï~ 


TUB eoi&iaicm THMBT 

For most chemical reactions, the individual chemical steps 
that make up the mechanism of the reaction cannot be observed. The 
mechanism is really atheory about what occurs step-by-step as the 
reactants are converted to products* Tbe slowest step ina reaction 
ntchanln determines the overall rate of reaction. 


Factors which affect the rate ofreaction are explained by the 
collision theory. The collision theory simply states that fora 
reaction tooccur, the species must collide with enough kinetic 
energy and inthe correct orientation* Byincreasing the 
concentration ofthe reactants, you increase the number of 
collisions. You can increase the number of collisions byincreasingthe surface area ofa solid reactant and increasing the temperature. 
Increasing the temperature ofthe reaction also increases the 
kinetic energy ofthe collision. The reactants must have the proper 
orientation for an effective collision* This is necessary tobreak 
existing bonds so new ones form. 

Let's look atthe reaction of hydrogen peroxide and iodide in 
aqueous solution to produce 13". From our discussion above, the 
rate may be expressed as: 


^'•yt9téy-± iüj _-i *m . ^¡Ï£Û 

3 dt 2 dt dt 


The rate law for the appearance ofproducts is: 


rate -k[H2023K {I~1r [H*]f 

where x, y,and z are the reaction orders and k is the rate 
constant. 
" • • • • "• " • ' • • . 
: • • ... . • • • • _ -. -• • • 
• .. . . 
• 7 • • : • . : • ' ." . • .-_ • " • • • • ' . 

• " ' 
Kinetics Page 2 



• 


• " 
156 


Expressing the rate by measuring the formation of products results 
in: 
rate -d[I3-] -1 à[HtO\ 
dt dt 

Combining these two expressions gives us the general rate law for 
reaction (1): 


k^O,]* [H+]T [I^]a (2) 


King(li describes two possible reaction mechanisms. The first 
mechanism is dependent on H* and is described by 


Mechanism m 

H+ + K-0 , #• HOOH, {%} fas t 

KOOH,* •*• I " ** H-02 + HOI (4) slaw 

2 fe-1

HOI + H* -? H*OT (5) fas t 

+ "3 
(6) fas t 
H2OE + I" -* K30 + I2 

.

Í7 ) fas t 

I ) + I" -W 

H30; + 2H+ t 31 -- 2ÍL0 + I3 ' 

Recall that the rate determining step is the slow step in the 
mechanism. The reacting species are obtained from the slowest step 
but the rate law may be expressed only in terms of the overall 
reactants. Based on this mechanistic path, what is the rate lav for 
the reaction? 


{1) King, E. L., HO W Chemical Reactions Occur, W. A. Benjamin Inc., 
pp 80-83, 1963. 


Kinetics Page 3 



Another possible mechanism for this reaction does not involve H+ 
This mechanism can be described by: 


Mechanism 2: 


H a 0 2 •'•*. r • OÍT + HOI (1Î) *low 


OH" + H+< Ha0 (13) fast 
HOI + H*. H^O I {14) fast 


H!OI + I* Ia + H30 (15) fast 


(16) fast 
12 +I


H303 + 2H+ + 31' ZK20 + I," 


Based on this reaction path, what is the rate law expression for the 
reaction? 


Notice that the first reaction mechanism results in a rate law that 
is dependent on [H+]* A straightforward way to determine which 
mechanism is valid would be to do several reactions while changing 
the acidity of the solution! If no change in the rate occurs, we 
know that the rate is independent of [H*] (y = 0 for equation 2). 


In order to visualize the reaction, you must think about the 
molecular interactions. For example, when the iodide ion reacts 
with the peroxide in acidic solution, equations 3 and 4, the 
sequence of reactions might appear as follows: 


Kinetics Page 4 



H 
.. " 


¡O—O--H 


fast 


O—0: + H* 


I " 
/ H 
H 


H 


[4] .. ! 
-H +
./
0. slow 
»T-0-I" \ 


H H 


'U 


H 


i 


fast 


H—0: + H* 


H ~ \ 

ÍT 


n

: 

* • ai


fast 


H—Q: + :_I Î _I_î 

W v


:i:i: -(- :j: —* fast 


:i: :fc 


What is the rate law for the reaction? 


Remember, the rate of a reaction is affected not only by the 
concentration of the reactants but also by the orientation of the 
molecules and their kinetic energy. The orientation and spatial 
relationship of the molecules are easier to visualize if you use 
VSËPR theory. y'• 


Kinetics Page 5 



PftBUJ TEMPERATURE EFFECT? 


1 For


'•'•y- this experiment you will need aluminum foil and zinc. 
First, place four squares of zinc and four squares of aluminum on 
the reaction surface* Place five to ten drops of 0*01 M HCl on one 
square of zinc, five to ten drops of o.l H HCl on the second square, 
five to ten drops of 1.0 H HCl on the third, and five to ten drops 
of 6.0 M HCl on the last square of zinc. Repeat for the aluminum. 
Record your observations below. Include a description of the 
reaction and times from addition of reactants to first observable 
reaction: 

Zinc Aluminum 


6.0 H HCl 
1*0 H HCl 


0,1 K HCl 


0.01 H HCl 
2. Repeat the above experiment but add one drop of Cu(N03)2 
to the HCl (place the HCl on the metal first) . Record your 


observations. 


Zinc Aluminum 


OJ<H03)7/ 


6.0 H HCl 
Cu<N0s)j/ 
l.Û H HClV 


Cu(H03)i/ 
0*1 H HCl 


CuíN03)í/ 
0*01 M HCl 


Does Cu{NO3}a act as a catalyst? (Remember a catalyst must increase 


the rate of the same reaction) 


Kinetics Page 6 



3. How place the metals on an ice cube and repeat the first series 
of experiments (without the Cu(No3)2) . Record your observations below. 
Zinc Aluminum 


6.0 H HCl 
1.0 H HCl 
Ó.1 M HCl 


0*01 M HCl 


F*«T H « COBCEHTRATIOH «FFBCTS 


A* introductory Procédures 


1, You will use the digital volt-ohm meter (d-VOM) for this 
part of the experiment. If you have any questions regarding the use 
of the d-VOH, ask your instructor. The procedure to set up the 
d-VOM for use is: 


a. Turn the main switch to ohms Ü. 
b. Press the oval button until "kû" is displayed. 
c. Plug the BLACK lead into the hole marked COM. Plug 
the RED lead into the hole marked V. 
Connect the leads to the cadmium sulfide (CdS) photoelectric cell. 
The CdS Cell is encased in a plastic support, and the support 
inserted into one of the wells of a 96 cell tray. The top of the 
cell will be used as your reaction surface, and should be as level 
as possible. 


2. The CdS cell responds to changes in light intensity with a 
change in resistance. It is important that the amount of light 
reaching the cell remain constant throughout the experiment. 
a. Place your hand over the CdS cell. Did the resistance 
decrease or increase? 
This is similar to what occurs as the Ia ~ concentration 
increases. The change in resistance is directly related 
to the change in concentration (rate of formation) of l3~. 
Ij~ is a yellow species which absorbs light, reducing the 
amount of light reaching the CdS cell. You must be very 
caraful in controlling the drop six* of your reactants. 
While the rate law is dependent on the concentration of 
the reactants, the photoelectric cell is reacting to the 
amount of light which reaches the cell. The greater the 

Kinetics Page 7 



depth of the reaction solution, the lower the 
transmittance* This phenomenon is described by Beer's 
Law. This law states that the light transmitted through a 
solution decreases as the concentration of light absorbing 
material in the solution increases. We will express this 
as: 


A - «Í[X] 


where A is the absorbance of the solution, e is a constant 


known as the molar absorptivity, t is the effective path 

length that the light travels through the solution, and 


[X] is the concentration of the light absorbing material. 
You can see how the drop size will effect t. 
B. Determination of Effect of Acid concentration 
1. Place one drop of 1 H H3Oa, two drops of buffer pH 1, and 
one drop 2 H Hal on top of the photoelectric cell. Add the Hal last 
and record the initltal resistance reading* Continue to record the 
resistance every 15 seconds for 5 minutes. 
2. Repeat this experiment but use the buffer pH 3 with the 1 M 
HjQs: and 2 K Nal. Record your results. 
Kinetics Page & 



I 62 


3. Again* repeat this experiment but use the buffer pH 6 with 
the i M H202 and 2 M Nal. Record your results. 
C. Initial Rate Determination 
1, In class, we have discussed using the initial rate method 
to determine the rate law from experimentally measured data. Devise 
an experiment in which you can determine the rate law by varying the 
concentrations of Hal and H20r. Use the buffer pH 6 for each 
experiment. The table below Is probably a good starting point. You 
must determine the rate from the data you collect. 


Trial Run Drops Initial Drops Initial Drops Rate of 
x2o2 [HjO,] Nal [Hal] pH6 Formation 
i y 

Kinetics Page 9 



coNCLPaioire 


1. Graph time vs. resistance for the l H H3o2, 2 M Hal and buffer 
with pHs 1, 3, and 6. Place these graphs on the same graph paper. 
Use of a spreadsheet will be very helpful here. 
a. What do these graphs tell you about the effect of acid 
concentration on the rate of reaction of Iodide with hydrogen 


peroxide? 


b* Write the rate law for this reaction based on whether path 


1 or path 2 is supported by your data: 


2. Graph time vs* resistance for the three experiments you 
conducted in the Initial Rats section of the lab. You may also 
place these graphs on the same sheet of graph paper. 
3. What is the initial rate of each trial in question 2* You may 
determine this from your graphs by taking the slope of your curve at 
time = Ó (use only the first few data points). Recall that the rate 
of change of your measured resistance is equal to the rate of change 
in concentration of I g"* 
Kinetics Page 10 



4. Using the inital rate data, determine the order of each reactant 
in your rate law for the reaction of sodium iodide and hydrogen 
peroxide. Write the rate law: 
5. Determine the rate constant for the rate law. Use the rate 
expression developed above. 
Kinetics Page 11 


• r -y-. . v-, 

1(5 3 


PRELAB PROBLEMA 


1. Using Lewis structures, draw a possible reaction for mechanism 
2, eq. 12. 
2. What color is the l3~ ion? 
3. Write a balanced redox equation for the reaction of Al metal 
with HCl: 
4. According to the Collision Theory, what two conditions must be 
met for a reaction to occur? 
5. What is the general rate law for the following reaction? 
H 2 O 2 + 2H* +y^i-y%^t^-:y.2H2o 

Kinetics Page 12 


•: 
... i . ' 



•.-.y. f.v y 
16 6 


"•;' ^ y 

:3?y 

•••'•j. • y y -: 

.Î ••. . . i 

Kinetics 

Page 13 

.-T-. ' . _ •. tw 


6UFCTROCHBKX BTRY 
Chemistry 131 


INTRODVCTIOH 


Electrochemistry, which deals with the conversion between 
electrical energy and chemical energy, is an important area of 
chemistry that touches many aspects of our lives. Were it not for 
an application of electrochemistry in the form of batteries, we . 
would still hand crank our cars to start them and wear windup 
watches. On the other hand, we would use slide rules rather than 
Calculators. Other applications affecting our daily lives include 
electroplating and corrosion of metals. Electrochemical measurement 
techniques are used routinely to determine extremely low 
concentrations of ions in solutions, identify unknown compounds, 
monitor reaction rates, and study a whole host of other phenomena. 


voltale Cells 


To convert the energy from a chemical reaction to electricity, 
a voltaic or galvanic cell (commonly known as a battery) is used. A 
voltaic cell operates on spontaneous oxidation-reduction reactions. 
The transfer of electrons between substances produces the 
electricity. 


Voltaic cells are constructed to separate the oxidation half 
reaction from the reduction haIf-reaction. This separation forces 
electrons to flow through an external circuit where they can do 
work. We can construct a galvanic cell using the spontaneous 
reaction of copper metal and Ag* ions. In the reaction, electrons 
are transferred from the copper metal to the silver ions* 


The balanced net ionic equation for this reaction Is: 


2A + 2A*<0 


« +<M> + Cu(o <* (Ml(1) 

Which can be divided into oxidation and reduction half-reactions: 


v 


•".• • 
" ••• .-" •"•-.••. 


• .-?•-. 
:rf"-j-: 

(=•'£?• 



im 

Equations 2 and 3 show what happened. The electrons lost from 
copper are used by Ag+ ions to form silver metal while Cuï+ goes 
into solution. Figure 1 shows a voltaic cell using these reactions. 


We use a standard notation to describe galvanic cells. For the 
copper/silver cell, this notation is: 


C<i(a, / Cu'*(aq) || Ag+(*q) / Ag(s) 


The slashed lines represent phase boundaries, the two solid lines 
represent half cell boundaries* By convention, the anode (oxidation 
reaction) is written first. 


(anode) (+) A* 7v 


^-{cathode) 


AgKO_ 


Cu(N03)2 


Cu(*>/ Cu2+í»q)U Ag*(*qî/ MU ) 


Voltaic Cell 


Figure 1 


According to equation 2, the Cu metal loses electrons 
(oxidation] forming Cuî+ ions. These "lost" electrons travel 
through the wire to the Ag metal, and combine with the Ag* ions 
(reduction) in solution, and Ag is plated out on the Ag electrode* 
The sale bridge completes the circuit by allowing inert ions such as 
Na* and N03" to travel through the salt bridge and maintain 777 
ionic-neutrality of the solution* Without a salt bridge, no current 
can flow because as A3* ions leave solution the remaining solution 
will have too many negatively charged nitrate ions, N03", left. By 
"intercepting11 the flow of electrons outside of the cell, we can use 
them to drive a motor and produce work or light a lamp or power a 
calculator. This will cause a positive deflection on a voltmeter, 
verifying current flow as shown above. Oxidation occurs at the 
a nod a, reduction at the cofcJiode, An easy way to remember this: both 
oxidation and anode begin with vowels while reduction and cathode 
begin with consonants. 

ELBCTROCBEIUBTRY Page 2 



Electrolytic cells 


An electrolytic cell is similar to a galvanic cell, but instead 
of obtaining electrical energy from spontaneous reactions, an 
electric source is attached to the electrodes "forcing11 electrons to 
move in the opposite direction. Thus a non-spontaneous reaction can 
be forced to occur by supplying electrical current, some uses of 
electrolytic cells include electroplating metals onto other metals, 
refining or purifying metals, charging "rechargeable" batteries (car 
batteries, nickel/cadmium batteries, etc.), decomposing water, and 
obtaining pure metals from their ores. 


This last application Is extremely important in the production 
of aluminum from its ore, AlsO?. Hall discovered this electrolytic 
process while still a college student in 1886. Before his 
discovery, aluminum was so expensive that the rich would flaunt 
their wealth by throwing dinner parties using aluminum utensils. 
Today, production of one mole of aluminum from its ore requires 
30,000 joules of energy. Recycling aluminum requires only 26.1 
joules per mole (the amount of energy required to melt aluminum). 
Thus, the emphasis on recycling aluminum rather than producing it 
from its ore is quite obvious. 


Standard Potentials and •* Calculations 


By using half-cell potentials, you can determine the voltage 
(potential) that any two haIf-reactions will produce* The 
potentials are not an absolute figure, but are measured relative to 
a reference half-reaction. The standard reduction potentials we see 
in tables in cur book and on the classroom wall are measured against 
the Normal Hydrogen Electrode {NHE) which has a reduction 
half-reaction: 


2H*(4 0 +• 2e > H i(i ) <4) 


The NHE has been assigned an £° or voltage of 0.00 V. Actually 
any reaction could be used as the reference but since hydrogen has 
been chosen,w© still continue with it. The symbol E* means 
standard half-reaction or "half-cell1' potential in which the 
reactant and product solutions are 1.00 If (1 atm for gases) and the 
metals are pure crystals. For convenience, we normally tabulate the 
E° values at 25*C, but it's not necessary. This method enables us 
to determine the potential for numerous half-reactions. By listing 
all half-reactions as reduction half-reactions, we generate a table 
of standard reduction potentials. 


A table of standard reduction potentials enables us to 
determine the potential of a redox reaction, determine the order of 
reactivity of some species towards others, and to determine which of 
two species is more susceptible towards reduction or oxidation in a 


ELECTROCRBHXBTftY Page 3 



spontaneous redox reaction* A positive E* value indicates that the 
substance will spontaneously reduce relative to equation 4, while 
negative values mean a spontaneous oxidation will occur relative to 
equation 4. By simply adding the E05t and Er(<i values, you can 
obtain E°, The E° value for reaction 1 is simply the sum of the E° 
values for reactions 2 and 3. Remember E°O H --E°f*d for a 
particular half-reaction* 

E° -E'oldltUl l +--B0 
r,„llt,i,ll = -0.337 V + 0.7994 V - .462V (5) 


Note that the E 4 K and Er#<J values are not multiplied by the 
stoichiometrix coefficients used to balance the redox equation, 


Nerast Equation 


In the late 18Q0's# a German chemist named Walter Nernst was 
studying the thermodynamics of electrolytic solutions. Through his 
research he developed a formula to calculate the potential of a cell 
that was not at standard temperature, pressure and concentration. 
It is called the Nernst equation. Using equations we covered in 
this and earlier chapters, we can arrive at this equation. Gibbs 
standard free energies (¿G) are given by equations 6 and 7. 


ÛGÙ « -nTE° (6) 


aG - -nFE (7) 


Here n equals the moles of electrons (a"} transferred in the 


reaction and F is Faraday's Constant (96,485 Coulombs/mole e~), 


Gibbs free energy is also given by equation 8. 


ÛG - AG* + HT In Q (8) 


where R is the gas constant (a.314 J K~Lmole"*)r T the temperature 
in Kelvin* and Q the reaction quotient <[Products]/[ReactantsJ)* 
Substituting equations 6 and 7 into equation 8 we obtain equation 9. 


-nFE » -nFEc + RT In Q (9) 


Dividing both sides by -nF we obtain equation 10, which is the 


Nernst equation. 


E - E° + (HT/-nF) In Q (10) 

Since we usually conduct redox reactions at room temperature f29@K), 
the quantity (RT/nF) conveniently becomes 0.0257/n and the Nernst 
equation at these conditions becomes 


E - E* - (Q.D357/n)-In Q (11) 

mUtCTBOCHBllIBTRY Pag e 4 


1/1 


corrosion/Corroslea Protection 


Corrosion is metal deterioration. It's mainly an 
electrochemical process, and under the right conditions, it occurs 
spontaneously in nature. The cost to the U.S. economy in 1984 due 
to metallic corrosion alone was over $80 billion.1 Thus, corrosion 
is Asignificant problem. With so many aircraft constantly exposed 
to the elements, corrosion is also a major Air Force concern. 


Before we can examine ways to protect against corrosion, we 
must first understand what happens during corrosion. Corrosion 
results from the operation of spontaneous electrochemical cells on '. 
the surface of the metal being considered. We'll limit our 
discussion here to steel (iron) samples. You might wish to imagine 
the body of a new automobile or a Soviet Hig 25 during the following 
discussion. 


As in all electrochemical reactions, oxidation happens at the 
anode. This occurs because the anode is more easily oxidised than 
the other metal in the cell. For solid metal, oxidation forms metal 
cations which are soluble in water. £o how does a steel car body 
rust if it consists only of steel? 


i ' r "-.'•'• . : 7 ."• • :; ' V •••• ' y.,-•"•'•""' ••.:• • '"•••• • "••' . '.' "•'• -• ' '' ""'.•. 

Well, in order for corrosion to occur, water must be present. 
Dissolved in the water is oxygen, which undergoes the following 
reduction haIf-reaction: 


. • •


0, +•2H*0 + 4eJ > 40H" 
• 


The other half-reaction is the oxidation of iron 


• • ' • . • 


. . • • • • • • . ' • . • • . . ' • • 

Fe(i, > **2*f*> + 2«" 
• 

Write the overall oxidation/reduction reaction that occurs during 
the formation of rust: 


Thus cars "rust-out" because of the air. Of course the process 
happens rapidly if certain other conditions are present. You must 
have some contact between the solutions of reactants* Water works 
best for this. The water works better if some ions present are 
already in solution (what does this provide?). Some states sprinkle 
rock salt on snowy roads. This provides all of the requirements, A 
small scratch in the paint, moisture, and an ionic solution, and the 
result is rust on your new car. 


i. Chan j„ H, , Cha m l s t ry, 3rd ;.«d.
.. , Hind am Htuit, Inc. Je v York,1 Hi 
pp.764-785. 
• .ELECTROCHEMISTRY 
Page 5 



Àis youvrecall,; Es values apply only to pure crystalline metal 
samples. Most steel samples do not have a uniform crystalline 
structure; instead, they have many small areas of differing crystal 
structure due to stresses that occur during processing of the 
metals. These different regions have different Ee values and this 
results in the formation of a voltaic cell. Thus, in steel samples, 
some of these regions serve as cathodes while others serve as 
anodes. This explains why rust occurs at separate, distinct points 
all over the metal surface instead of covering the entire metal 
surface uniformly. 


One obvious solution to this problem is to produce steel with 
one uniform crystalline structure. Technologically, this would be 
extremely expensive, if not impossible at this time. Therefore, 
some other means of protecting steel (iron) structures must be 
devised, some of the simpler and less effective of these means 
include surface coverings (paint, etc.) and passivation. 
Passivation is the process of coating the surface of the steel 
sample with another metal (usually less reactive or not as easily 
oxidised as steel.) Both of these methods protect the steel by 
covering it* Unfortunately, once the surface coating is removed or 
scratched: off, rust develops* 


A better way to protect the steel would be to electrically 
connect the steel to be protected to a more active (more easily 
oxidizable) metal, such as magnesium. In this way, the steel 
becomes the cathode and the magnesium becomes the "sacrificial" 
anode. This process is known as cathcdic protection* In order to 
incorporate cathodic protection, the sacrificial metal must have a 
lower reduction potential than the metal to be protected* 


A '•,y- ELECTROCHEMISTRY Page• S 



•yyyy •••••'yyy.---yy-yyyyyyíji. 
EZPBMHBNTAL 


$.s Voltaic eel is (Batteries> Î 


In this experiment, you will set up voltaic cells, as in Figure 
1> but on a micro-scale. Refer to Figure l often to identify the 
parts of the cell with your setup. 


1. Cut a piece of rectangular filter paper in half so you have 
two squares. Fold it in fourths and cut it to form a "+", as 
demonstrated by your instructor. Save the leftover pieces for Pair t s; 
,C:,at$7p,7 


2: : Take s orne of the copper, zinc, 1e ad, and Iron and 1ightly 
sand off the oxide layer on them. 
v HOTEi Read Steps 3-*5 hefore performing them, because you'll 
want to take quick measurements for more accurate results. 


v 3* 7 Placé the7four metals at different corners of your newly cut 
filter paper. Place 5-10 drops of 0.10 M NaN03 solution in the 
middle of the filter paper. What function does the NaN03 perform? 


4. Place a few drops of 0.10 H Cu(NO3)3 to the side of or 
underneath the copper. Do the same with the 0.10 M FeS04, 0.10 M 
Pb(NO' )2/ and 0.X0 H Zn(NO 3)2 solutions and their respectiye metaIs, 
The drops of solution must touch their metals and these solutions 
must just touch the NaKOs solution* 
S,... 'Take your mu1timeter and se1act DC: volts (VDC) and measure 
the voltage between the following: 

a. Hake sure your leads touch bare metal and are not 
themselves immersed into the solution* 
b. Switch leads from your multimeter on the metals if you 
get a negative voltage (you want a positive voltage indicating a 
spent aneous react ion). :* , 
V.;*ÍV)::: :';^fv>A; 
cu andAZnII 7. MI^Mfff 7 
<• en and Fe V.'.-. 
'••''• . 
c u and Pb 

ELEcrnocHfiMrsTsY page 7 

• \ . • • .i . 

6* Why are these values recorded as E and not E°? 


7. Calculate the theoretical E" values for the three voltaic 
cells you constructed (using a Table of Standard Reduction 
Potentials) and write the two haIf-reactions for each cell as they 
occurred. 
a. Cu and Zn 
aim 

Reduction Half-reaction; 


oxidation HaIf-reaction: 


Measured E =.• •y.._l7_-l.-i.-_..'•' Theoretical E° = 


b. Cu and Pb 
¿itsa 

Reduction Half-reactioni 


'oxidation Half-reaction: 


Measured E -• •
••• •: y .•-.-;. Theoretical E* 
c*-7:;'CU.an4:-Fís • ElXYi 
=Reduction Half-réaction: 

oxidation Halt -reaction: 

Ee

Measured E Theoretical

ELCCTftOCHEHIflTRY Page 8 


JL75 


8. Why does the measured E not equal the theoretical E°? 
B. Reduction Potentials: 
As you recall, the standard reduction potentials of all elements 
are measured relative to the reduction of the following reaction: 


2H+ <-A <, ) + 2e_ • ----*-^> H2(r¡ É* - 0:00 Volts (12) 

In this part of the lab you will construct a table of reduction 
potentials for some metals* but you will use the following as the 
reference: 


Pb?*(i,j +2* ~ -----> Pb (i ) E° = 0.00 Volts (13) 


Note. Redox processes could be measured against any reference 
as long as they're all measured against that same reference. 


1. Using the same setup as before, measure the potential of the 
following metals and their 0,10 H solutions against the Pb/Pb3* 
system: 
Al, Cu, Fe, Pb, Mg* Zn. 


á. Based on your knowledge of the table of standard 
reduction potentials, predict the order of reduction potentials vs. 
Pb before you measure them. 


b. Make sure you first sand any oxide layers off the 
metals. To ibqkat the reduction half~reactions for the metals,, put 
the negative (black) lead on the Pb/Pb2+ electrode. The negative 
lead is the source of electrons; therefore with the black lead on 
Pb, ypii ;knpw Pb is being oxidized. 
6 i. In making ..ye ur: t a Jbi e, report all va lues ( even thé 
negative ones), rank order all potentials, and write half-reactions 
as reduction reactions. 

ELECTROCHEMISTRY Page 9 



...... . . .. 


' • • • . • • • • '"••' : ' • " •'• . • • . " -• • • . 
. . . • -, . . • . -. . -. : • ' • • • • 

•••••' • .••'••• 

• 

176 

.... 

; ..' .V 

ï + : •••••.:•:•. .

Metal jteduct í on Pot entla1 vs. Pb/Pb 

• . • . 
•••:.:...• -y ., • • 

Al 

• .•. • 

: " • • 

• 

c« 

• y • 

•••••• . 

. 

Fe 
Fe 
. ••• 

-

• .-. 
• 

"y .••• '

•

• . • 

• '• ' ; 
•yy 
•"'..'" 
¥b 

-• 

-, • 

: 

-. • • • 

"__••-• 

H* 

'• • : 

.:••". 

'•"••. 

• . 
Zo 

Using this information, write your Table of Standard Reduction 
Potentials below. THIS IS DATA FROM YOUR EXPERIMENTS. DO NOTUSEA 
TABLE OF STANDARD REDUCTION POTENTIALS FOR THIS TABLE. 


TABLE OF STANDARD REDUCTION POTENTIALS 


Reduction Half-Reaction B« fvolts\ 

-'•'•.'. •• yû:yy • ' '.'•.''.• -. 7 • ,' ' /••" "• • • • îh y 7..' ''•&•& 

yyyyyy .y. 'y .-.. ••-.-yy. •••''•• y-y y-y ,. yy yy '• y', y y -7 yy;. yy y .. y .7 . yyy: 

: • : • . •• 

' . • . . . • • ' • ' : • • " . ' " • : . • • • • . . • • • . ' • . ' . ' . • • " . ' . • • '• " ' . ' : . 

'•":•• •• •,' ! -" ' . ' • . .y ; '. : •:• y.: . •' • -. • ••.'••• .7 -•> • .' -: ' •. " : y . '• • y .'.. • • 

>.•• :-.-.. ;:• A ••.'••:••, • •.'... : ••-•... '.'.'•• ,'•• . • • 7 '... • .: : .-7 ' " ' ' '.••.': 

... . 
.... 
" y • • ••'• . .. . • • • . ' • " -• " • • . • • -•.. • .... . 

. . . . ; . . • • • • . . ' "..', • . -. * . . . • . • • : ' ' . -. • • • • . • • • 

. • • . . ' ' ' • ' • . . . . . • . • ' • • . ; • . 

.. •• . •:'•• • • . ; : • • ' : • . ' . . . y ' :• '•• '. • 

:••. • 

C* Nernst Equation t 

'•'• 7 yyyyy 

yy y . yy --. y y , .•-.: • ,••.-• • • -' y • • . y :••-• . 7-y-y •••. : y ••• -y-.y 

l. Using thé same procedure as in Part A, set up the following 
copper and lead voltaic cell: 
• 7 7 ' •;. • •••;. ....••. y . ' •• • yyyyy-.y'1-"7 -yy--•.;• -y v . -.-. • •
•• •'•...-. .• _ : yy. y . ' .: • 

Pb (i) / Pbï+(aq. 0.10 M) \\ Cú**íáq. O.ION)/ Çiï'¿^ 

•: 
••.. • 

;*... Measu r e -E for this eel I•• 

• . . . • . . . . -. • . . . _ • • • • • • . . . • . • • . • , • . . . . . . • . ! • . , -. . 
•.'•'" , 
•'...: • . • . . • • • • • • • • " • ' ' ' . 
'. . • " -. . ' : • . • ; • • . • • . -• • . " ' ' . • . • •• • ' • ' . ' . • • • . . ' • y • ••. • 

.• • ••.. y .....•' 7. " • ''•:'•' , 'y . • : •• • : . . : 

,:; .•;-.• 7 -.•••• ' .7: . ••' ; . 7 ;. •_• ;.'•• -• '• • ' 7 • .•• •. -; 7 . • • •• >••'' 

•• '• . •' ' . ' "' -: • ' ' • ' . " • '. -• , : • •: : ' • • • ': ' • ' ' '• • . :• . ' • . . . . ' ,-. . : .-' : 

ELECTROCHEMISTRY page 10 ^ 


' '•' ••"•'• ••: •. 



:' • '• •'•:'••.'• 


• . •• ' 
:••;•• -•-• 


• 


y.y y'myy 
' 

b. Calculate E' for this cell using a table of Standard 
Reduction Potentials. 
.••. • 


1. ._ .... 
• 

• • ' • • " . . * -' 
. .. -, • • • •• 

-... 

• y ••'• 

...;. . [. y--• . i: . . -. ; • .

'• A ..... .. .: .; 

•... y 
•••••• '•: •• ••: • y • • • 

• • :,' ^ .' ' 

... 

..':'• 7 7 

' • • • .• 

' • . ' . . • '•'.''• •: 

. •-• 


':.-•• yyy y • 
- . 7. ''•

•••" ,'.•• yy •'.• • '..••• y 

'.;-.'• 

'.' • 

. •"- •• • ...• 

-. . 
. . ' • ..'•'" 

' ' • • • . . 7 - • " . ..• -' 

c. Using the Nernst equation, calculate E for the 0.10 M cell, 
,' -A ;: , y .••'•: , •.' -7. 7 . • '-.y . . y : • .. . ;. '.' • •:..:' . 

•:; : 
y . .'•: .7 . .:.:'• A -.• .•;.•-. .-: ; • .. . . -y . •'• . 
y . • 
•. . y • 


d. Does the calculated E for the 0,1 M cell equal the E* for 
the standard 1.0 M cell? why? 
: • " 


.• 7 •'•:'•'. | •'•:. • . ..... . 7. " ' y y '••••'•'••• 

. 

••• '. •. 


2. How measure E as the concentrations of Pb2 + and Cu2* are 
changed. Fill out the table and show your calculations for the E 
values. Use leftover filter paper pieces from Part A 


.• "•. 

tCu*t] tPb**l\ E¿¿^(V) ;>cYic 


• i 
... '-• 
•• 7. ••• ' 

-y 


7 '' • ••>• y-.

1.00 M O.IQ M ; 
1 • • ; •. 

:'•• • 

yyy A 
• 

-•'••• ' 

• .•• 
• 

.. y 

1.O0 M 0*0010 M 

. • ;. 

•: 
.: :

^ 

• 

•

[Pb2 +3 0. 10 H: 
. • 

• ': ; 
•

fk . -. •-. 

• . 
• 

• ••: 

• •. 

..•• : .... . ••... ' • •"• -y .. -. . • 

•. y y.y- ik 
• y.y . •' . ••.'....•• '•. 7 • . • • • • ' . ' • ':....:: ....-. ' - y . •:• 

• : • . • • . • • • • • • " . • . . • • • . • . -• . • ' • ••'•• : 
• .• • •••• . ; • • • • • • • ' • - .:• • • • " • " . •;, . . . . . . ••. :••• ' 
• : • . . : •. , . . • : *:' '.;•-. • • • . • ' • '•-;• •'" . ;."••-•. 
y-.. .• "• "•• •• •'•••• y . .'• y-. '••> •. • -y y •.•••'•• . 

• ' : • • • 
• •;• ; ;;•;•; • y y yy y ;••' 7. 

. ••' • . •:-.-• •-• 
• 


: ••• • ••••y"' -' ' •'• . " ' ' • • -• :

• . ••.. ••• 
•;. ..•• .:.• •••-• • " y 
.: • '••••'. 

y yA yy . -A y-y.


;: • 

y''. :y':

• • • . • ' 

• . • . . • • • • 
' " •• •• 

. ' '•' 

i 

:••••" • . y y 

-. •".• • •<• .: ,• •• 
• 

•. • . • • 
•••', yy-, r '.•••'• y

• .• • : : 
'y , y yy'S'y. • y'' -y'-' ••• 'yyiyï ,.7:: • • 

\ 

ELECTROCHEMISTRY Page 11 


• :....• 

..':,• • '.:..•-.. ' 

• -••'; .'•' • y y

. _ . . . . . . :••• ••:; y y '••'•'.'.y :'•'.••-• • •••• • "• yy.-• . ... 


si + 1 =


calculations for E when [Cu2+] - 1.00 M and [Pbi+] - o.ooio M: 

3, It is possible to set up a voltaic cell that is driv>; by a 
concentration difference. A potential will continue to exist 
between the two half cells until their concentrations are equal. 
Investigate this by setting up the following cell; 


7Cu(,)/Cu2+(aq1 0(10H) M Cu'*<aq, 1,00 M) / Cu(t ) 


a. Measure E for this cell. (Bernetuber to switch the meter leads 
if needed to obtain a positive voltage*} 
• -7 • R7*v y-y. - y ' 

b. Under standard conditions, what would E* for this cell be? 
• -\. 

:--yyyy. ' .y y;7 

c. Deviations from standard conditions (one molar concentration 
for ions, one atmosphere pressure, and standard phases at 25e C) 
cause the déviation of/your .Observed £ in paragraph C.3.a. from E* 7 
Use the Nernst Equation to calculate E for this cell. Compare your 
measured and calculated values for E* 
•ay Are there other Conditions besides standard conditions that 
..Willygive anE«o for this cell? What are they? y 
ELECTROCHEMISTRY Page 13 



D. Electrolytic Cells: 
l* Electrolysis of Copper. 


a. Partially fill one well in your reaction tray with 1.00 
M Cu{N05)a. Clean the oxide layer off a nail with 1 M HCl and place 
it into the solution. Wait for about 30 seconds, what occurred? 
•• . ••: 


b. Write the balanced redox equation for the reaction you 
just observed: 
• • • y •'•••' . • • ; • " 


.. . : : . • .' . •. 


c* Remove the nail from solution and describe it. 


d. What is the coating on the nail? 
. •• .-.••• '•• 
• y 
•• 

e. What haIf-reactions took place to produce this coating? 
Oxidation haIf-reaction: 


: Réduction haIf-reaction: 


; •:fi, •.. • • S ïnce wé. didrt' t add any energy to make th i s react i on 
occur. It's a spontaneous process. As in Part A, we could have 
separated the two half-reactions of the above redox reaction and 
produced a galvanic cell; however, we'll now make an electrolytic 
cell and reverse this spontaneous process. To do this, hook up the 
.9-volt battery to the nail and to a graphite electrode, then immerse 
both in the 1.00 H Cu (NO-,) 2 solution. Will you need to hook up the 
positive or negative lead to the nail? Why? (Hint, the positive 
lead withdraws electrons into the battery while the negative lead 
"pumps" electrons from the battery.) Stop the experiment shortly 
after the nail begins to bubble violently. Before stopping the 
experiment, observe the nail in the solution. Does anything happen 
when you withdraw it from the solution? 10-15 seconds after that? 

g. To show that the Cu(H03)a solution will not react 
spontaneously with the graphite, place a piece of graphite in the 
solution for about one minute, withdraw the graphite and describe 
significant changes, if any. 
ELECTROCHEMISTRY Page 13 



h, Did your electrolytic cell cause the spontaneous 
teaction above (steps a-6} to be reversed? Explain how you know 
this.' 


i. Why did the nail regain the copper color about 5-10 
seconds after withdrawing it from solution? 
• •. • 


j. Write the half-reaction which occurred on the graphite. 
k. Did your nail perform as an anode or cathode? Explain 
your-answer.-y 
1. Write the half-reaction which occurred on the nail once 
the plated metal w«s removed. 
2, Electrolysis of Water 


Hydrogen is an excellent fuel that someday may be used in place 
of jet fuel for aircraft or gasoline for cars. Wow it is used in 
the space program to propel the space Shuttle into orbit. Hydrogen 
though, does not exist by itself in abundant quantities on earth 

(most of it has escaped the earth's gravitational pull over the 
ages) . One method cf producing it is through the electrolysis of 
water, that isf using electricity to break water into H2 and o2. 


. 


• • • • 


: .: 


• 

ELECTROCHEMISTRY Page 14 



1S1 


a*•'•':' Make ah- electrolytic cell us ing ybui: 9-voit battery and 
graphite electrodes to electroiyzé water. Use Q.10 M MiwOj solution 
and add 4-6 drops of Yamada Indicator. Why must you use water with 


.10 M NaH03 in it?yy. 


b. Connect the battery and list your observations. 
c. What is the color around the cathode? What gas is evolving 
from the cathode? (Hint: To determine which electrode is the 
cathode, look at the oxidation and reduction haIf-reactions for the 
e1ectrolysis o f water • in your text.7 Use the color of Yamada 
indicator to determine which half-reaction is being produced at each 
electrode. Remember that Yamada gives red in acid and blue in 
Wi;*•*:•••-).' y7y7-y \ • 7\v;;.-;;.'.7; '. y^ y y y yy:yyyyy 
d. Write the half-reaction occurring at the cathode. 
e. What color exists at the anode and what gas evolves there? 
Write the half-reaction occurring at the anode. 
E* cathodic Protection* 


1. Which of the metals you have been given could be used to 
cathodically protect an iron nail? (Which metals oxidize easier 
than iron?} • ••••: y 
72. Of these metals, which do yóü think would best protect iron 
from corrosion? -7:7. 
3* Wrap a piece of the metal you identified in paragraph E.Z. 
around the head of a nail. Place an "unprotected" nail and the 
"protected" nail: in plastic cupé containing enough de-ionized H¡¿q7to 
cover the nails completely and add-4-6 drops of Yamada indicator. 
What happens? (The reaction may take ú^ to 10 minutes to occur*) 


ELECTROCHEMISTRY Page 15 



: 4. Repeat this for another. Qrie;of the meta1s Identi fled in. 
paragraph E.l. Observations? 


£y Application Experiments: 


1. Recovering Iodine From Solution 
Given an aqueous solution/.of Hal,, we; can set up an electrolytic 
cell/tó collect'"-la*•'/•":'"'• y 

a* When Wal is placed into water it dissociates into what 
ions? 


b. We want to convert the iodide ion, I", to I2. Write the 
half-reaction for this process. Is this an oxidation or reduction 
reaction? .7-y -7y: 77yV;..y >'•",: 
•• -^ y.. . 

• • • "'.'•' "" 
'"•"' • : ^ 
.y . .'• •"••-•_..' 
. . • 
' '--:.' -. 777- 7;, 

c. in order to make the electrolytic cell, you'll use two 
graphite electrodes connected to the battery. Use the graphite 
electrodes from the electrolysis experiment. Do not use the Yamada 
indicator in this experiment. (Graphite is a common "inert" 
electrode. It plays no part in the balanced redox equation; it 
s imply acts a» à réaction sur face,.éIIowing-the-] électrons to f 1 ow in 
the cellii ) y/ 
d. Which electrode on the battery (positive or negative) 
cause the haIf-reaction in paragraph F.l.b.? 
'•y.yy -"•"•'•' ' '. . .• ...••' •'• ' : . . •• • • -^. • ••••.-•, • \ .-"• \ •• :; . v •• -¡ 

••'; .•.•.:". 


e* will it be the anode or cathode? why? , ' y 


"..••• • ; 


:


.y.-.---'/' ' ' ' 

ELECTROCHEMISTRY Page 16 



f. Perform the experiment using O.io M prál. Report your 
procedure, observations and results. 


black color emanating from one of the electrodes is the I2. 
The yellow color is I3" {which you saw in previous experiments.) 


2. Optional Experiments. Additional experimentation is allowed if 
cleared by your instructor. Use the next page to design your 
experiment before you ask your instructor's permission. Your 
experimental design must include a background statement (what 
chemical theory do you plan to use), experimental section detailing 
what you will do, and a conclusion section. Some ideas to help you 
get started: 
a. Electroplate an object, what voltage is produced during 
your experiment? 


b. Voltaic cells in series vs. single cells* 
c. Voltaic calls connected In parallel vs. single cells or 
series cells. y : 


•:<?•'• : 
BLSCTBOCHBKIBTRY Page 17 



. ' • • • . . 

OPTIOHAL EXPERIMENT DESIGN PAGE 


• :•••• : s • . ','•'. 
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yy•.. 

4i"i "<•; ."• J""-' 

Pag e 10 


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' • y • ' 
' .•-" vV^-X':y..~^77: 

.. • •• •

I. Ifwewere toconstruct abattery inwhich the oxidation of Pb 
Would beone half-reaction, which metal inYOUR TABLE of 
half-reactions would give usthe highest voltage? Justify your 
answer. 
y• :••
7 
.. -. • 7"^, . : yy . y :•::'. y •••••.• y:\7 y 
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• 
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2* Usihg your data collected for the CuandPbgalvanic cell, draw 
a schematic ofthis ceil. Label allimportant parts* givethe 
oxidation andreduction half-reactions for the appropriate 
electrodes andshow the correct flow ofelectrons andions. Draw 
thé line notation forthe cell. Refer toFigure 1and/or your text 
for help* (You can draw itlike Figure 1ordraw aside view of 


••..• 


your experimento 


-..••*,-. y ••-. .• ¡„ • • ..., v ' 7 -. • • ^ 
•,;•... . .,•*:. >.•. ••••: 


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' . . / • • •."..' • 
. ' ••'• • • • ' • • " " y'i \ y... •. 'y-y •* 

3. Define standard state for asolid, liquid, gas, andion in 
solution. Are wejustified incomparing the experimental results to 
the theoretical values for Ep? Explain your answer. 
;; •.:•:._• 7 -;. •. . ; .:.• 7y; /v'.';77-'•.''^ • •
•• ' y, •'..'. A'--7.
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yy. . • : ."•••: ' .-••...•.••.... /.....•,..•,..•••.. 
•.•• . y y y : y -y . ••'•/•: 

4. Eased onYOUR TABLE ofStandard Reduction Potentials: yy 
. • • • • • .. .. . ' '••••• . 1 . • • • • ' • • ' . • ' • . ' . •.iT.^y.-y.y... . ; ..i, y.'-:.:••• ..;•.-.LT'cf..'-:. 

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v, m..,^ vl c i t--m M.mm-mr v.* v ^vuw w v.v.av^w.^ . <,>T^ ^ 
..-!;."•-!'/ i'-;'• :'< '777

a. Which metal has the greatest potentialn -n 
tobeoxidized? 
:^i^ . •; V . . . • . . -. . • . '••'.-' . ". • " " . * " • • • • • -_ • • . " . " ' -: • . • . • y:' y. 

Which metal has the lowest potential tobeoxidised? 


• --•••-.yxj

c. Give anexample where these metals are used by society.Do 
the applications fit the behavior ofthe metal toward oxidation? 
(See Chapter 22inthe text for adiscussion onthe uses ofthese 
metals.) 
y . y •• . • ' . • ' • • • . • • • .-• • •.• • •••••• • ...•• • •.•:• . • .:•,-.•: . if--y y 
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yyy,^yyyyyyy .• 7. . . 
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7 v^ i*y y y •'•• ' y — 

• •-••. • ¿y 


•*'••-.-•-••. ••• . 
:!••• !••. i-: • 1

ELECTROCHEMISTRY Page 19 

r ' : • ••' -• ••••' • • • • • : ' • •"."' • . . ... " • ' . • -. • • • : • •'•' • • . ' " 7 . -. ' 
• . • • • • . -. " 


.y-yyy yyyyy-y-y y ••':.. yy y y.;. '-yy:yy .-• yy-y? ••••••••••. .• yy y 7-....-A-• yyyy y 

lue 


5. Acids can oxidize some metals. Based on oxidation potentials, 
y •• • 


.-• 


would HCl oxidize Mg and En? Justify your answer using half 
reactions. 


.-. • . • • • • -. •• . y y y -y --. -. •• •• y ••••.-• • * : • ••• : ••••••• y y 

:• 7 -y ' • • ' • • • "•-.• . ' • -• . ' • • • : • --. • ..••••• • . -• • -• " 
• . : . • . ' 

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,.•7:77/ yyyyy y-y - y-y-y 7 7 . y-.. .•••• -A • .y . -,;•••••, ./ ;... , ; . y 

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. . 
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"y. .'.-.' y:'. ..A. A . ' " ; • 7? : • '• ' • A ' • • • 7 • -''•. • : y 

6. For the voltaic cell in Figure 1, would you expect the mass of 
the Cu electrode to increase or decrease? Would the mass of the Ag 
electrode increase or decrease? Explain. 
... 
• .• • y 

• . • •• -• : -: --: . • • • . . • . . . . • • . • • • . • 
. . . •: . . . -. • • • -y . • •. • • 

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-.-. • 

7. 
Briefly describe the differences between a voltaic cell and an • •• 
. • . •. •
electrolytic cell. 


.-.. -y.• .7 

y \ . y••/.'"• .-•'•• .'"•'• '77' ..7-7' -•77 

.- •'•!•: • • ' •-."• ' • /.; • • 
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...... 

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• .• , ' • .y . • '••• y • 
' • ' . . ' 
• A . ' .• • • . ' • • • . • ' ' . • ' . ' ' ' . : ' • • • • ' . • •.••• • • • 
8. Why were there difference s between your measured E and th e .. 
"••'•"_ .•" • --:

calculated EK il l Pa*rt !C? 

•. y • . • . • . • • • • 
-. . -• • • :••••• . 
7:;•-•••• 
-• -y-• • : • • : 
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'•"-' y . 7 "-••' -"•'•:. :...'. •" 
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/ELECTROCHEMISTRY Page 20 

•..•7. M.-• 
• -• • ! •.::•":••-.•: >••••.-'•.'. • .••• . y y y. , ' y 7 .-V-..: ' .• .' v • : • " • , , . ' • , " '.'• • ;: ' -: : •• •• y • .": • ,.-.'•.... . y y - .•':•'..•...•• . • ; * : ' • •;• •/'.•••• 


L • 


Sv J. f/;we;bubbled C17 gas through our 0•
•• 10 M HáX so1ution;7dpu1d> we 
obtain I2? Explain using concepts you have learned in 
electrochemistry. 

*• ' ".-*. 


• . • • 
•••• •••••. -..


. > " . " • • • • . . " 


•.-..• • 
• ; ••:'.• .:: 7. 7^ 
A . ., ;• • .. .. ' 
. — .. ; yy • 


"y .yy


10. Explain how the Nernst Equation can be used to determine the 
equilibrium constant for a balanced redox reaction. 
:•: •'•• :• • • • , ' • ' 

.•<,-.. 

11, Should m change in temperature affect E' for a cell? 


ELECTROCHEMISTRY Page 21 


. ; ' \*' ' ' V •.' ".. '• •".'.•'.'.':"'• •••:• • 'Í. •• -••••-'•:• ••-•• '•'" • " '••'"•• v . '-'.••' :':• . :' -. . •••. : '.' .• ••.... . •• •• • 


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ELECTftOCHEHÏSTHlf LABORATDHY 


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ELECTROCHEMISTRY page 22 



ELECTROCHEMISTRY SMALL-SCALE LAB 
PRELAB £U5STIORS 


These questions must ba completed Lài^i-E coming to lab. If you 
cannot answer all of them» reread this handout, and try them agair,, 


1, 
a. What is oxidation? 


b. 
Reduction? 
2. 
a- Draw a Voltaic Cell based on the following: 
Zn(V ,./ 2h {NO3) i {£-q y } j CU (HO3 ) j( , q ) / CU( *) 
• 

-

. • 

• 


; • 


• 


" 


•'•.. • 
' 
b. Label the anode, cathode, salt bridge and Indicate the 
flow of electrons and ions. 
c. 
Write the oxidation and reduction reactions. 
d. 
Write the overall net reaction. 
3. Calculate the voltage you could get out of a voltaic cell made 
of silver with its ion, Ag*, and tin with its ion, Sna+ . (Hí^tí you 
do not need to balance the redox reaction. Use the Standard 
Reduction Potentials Table on page 740 of your text. 
a. 
Write the reduction haIf-réactions. • 
Deduction Half-reaction E* (V) 
Silver: 
Tin: 


ELECTROCHEMISTRY Page 23 


"•y y • 

• 

• . 
190' 


b. Write the oxidation and reduction reactions that occur in 
this Cell: 
. 


Oxidation reaction: , 


Seduction reaction: 


c. Write the overall net reaction. 
d. Calculate E°. 
Your Efl should be positive sines this is a spontaneous 


electrochemical cell. If it is not, you need to switch the 


direction of your oxidation and reduction reactions. 


'•'•:•'•' '•'•• • .'' -." 

•;'• 

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ELECTROCBEMISTRY Pag e 2 4 

•• • • . : • • : • • • • . • • • ' 

• V .

lyi • 

ORGAMIC LABORATORY 


i . .•..•• . .• • • 
-• •• •. . •


chemistry iai 
IKTRODOCTlON 


Organic chemistry is the chemistry of carbon compounds. Prior to 
1Ô2B, the term "organic" was used to describe chemicals that had come 
from animals or plants. in that year, Friedrich Wohler showed that 
organic chemicals could be synthesized from inorganic sources. He 
synthesized urea simply by heating the salt ammonium cyanate. 


NH4OCH 


(inorganic) (organic) 


Organic chemistry is a very large field of chemistry. Over 7 
million organic chemicals are known. Our lives are surrounded by, and 
sustained by organic chemistry. For this reason it is alto referred to 
as the chemistry of life. The large number of organic compounds can be 
understood more easily by subdividing them into groups. These groups 
can arise from the classification of organic compounds based on their 
physical characteristics and reactions. These two factors are 
influenced by the "functional groups" within each molecule. 


Thépry 7/7; 77.yy ::y: "'•:>'..'•' 

A functional group is a group of atoms that act as a site where 
chemical reactions may occur. Reactions may also occur in the 
proximity of functional groups* in this lab, you will investigate 
several important functional groups. These Include amines, carboxyiic 
acids* alcohols, aldehydes, ketonesF and esters. Examples of each are 
shown in Figure l. These eight functional groups are found in the 
majority of organic compounds in your body and in nature. 


:y y


H 
H 
0 :7 V: o • 
7,11 

7.Í7 yy It 

R-C-H R-C-R" 


R-H-H 
AHÏHE ALDEHYDE KETONE 


'')&.':' y .707 
J! 7 7 

. 1 7


R-C-OH S-C-OE* 


R-C-OH 


•X.V 
CARBOXÏLIC ESTER 
7*'y7 y: 


:


ALCOHOL A«|> - 'y 'A y/A:/;,/,/..._ 

•7i|-o7: : 
'1 


R-H-C-R' R-C-X 


" . : • " • '' 


V-.X-. AHISE ACID HALIDE 


y y : ..." • • / )• -:,..(: •• 

i "..-. • 

PIQURE i 
Organic Functional croups 


Organic Laboratory Paga 1


'•> y y 

:*&* 


•yy .y y®. 
Amines are characterized as bases, and generally are polar. This 
polarity is due to a lone pair of electrons on the nitrogen and 
differences in the electronegativities of other substituents on the 
nitrogen. All amino acids contain an amine group fas veil as a 
carboxyiic acid group). Many pharmaceuticals also contain amines. An 
example is amphetamine (e.g.r or-methylben^eneethanatnine) , which 
stimulates thé central nervous system and is addictive. 


m2 

o - Nethylbenzeneethananine 


An Important reaction of amines is the condensation reaction 
with carboxyiic acids to form an amide, This reaction can be used to 
produce polymeric chains such as nylon: 


0 H 

0 

1 

1 1 

UH3-C-OH + H-li-CHa-CH3 '+ • HjO 
: *7c 
è i t <e * e I4. . *t)f y I •• «pi t n * : : •"'• :':«tNJÍ' »"t«m¡' * .-/. 


Alcohols are common organic substances that are encountered 
everyday* Ethyl alcohol (grain alcohol) is found In beverages, over 
the counter drugs, and cosmetics. Methyl alcohol (wood alcohol} is 
used as a fuel, secondary alcohols react to form ketones in the 
presence of oxidizing agents. Potassium dichromate may be used as the 
oxidizing agent and the reaction Is accompanied by a color change 
because dichromate is reduced to Cr3*. Ths general reaction is: 


otf 


-


KjCr-jO, / H2S04


R-CH-R' R-C-R 


This reaction will be done as a demonstration during this lab 
using isopropyl alcohol (rubbing alcohol) and acetone (the dimethyl 
ketone used in fingernail polish remover). You should note this is 
an oxidation-reduction reaction. What is being reduced? Write the 
reduction half-reaction in the box below: 


Aldehydes are good solvents for polar substances such as 
alcohols. Formaldehyde is a gas that is used in aqueous solutions 
called formalin. Formaldehyde is used to preserve biological 


organic laboratory Page 2 



:• y ••":•'• 77/. y y •••.. >•• •• • • y 193 

specimens and as the basis of many glues, resins, and polymers. 
Acetaldehyde is useâ in its trimeric form as a sedative. Aldehydes 
react to form carhoxylic acids in the presence of oxidizing agents. 
A qualitative test for aldehydes is the Tollens test. Tollens 
Reagent contains the silver ammonium ion, Ag(NH3)2 
H7 and will 
distinguish an aldehyde from a ketone, even though their functional 
groups are similiar. As the aldehyde is oxidized, the silver is 
reduced to a solid. A silver mirror will then form on the surface of 
the test tube (if the glass is clean). A finely divided black 
precipitate of silver appears if the glass isn't clean. The general 
reaction is; 


O 7.7 ' • Ó*\ 
• y.yry y . ïi • 


" *NH H


Ê-C-R'.:.+ 2 Aa{KHs')-a*.-+. 3 Oui-
:-—r^-' R-C*Q'+ ï *&(w) +3 + ¿2 ° 

Carboxyiic acids have characteristically sour tastes and odors. 
Acetic acid is an example of a carboxyiic acid, and is primarily 
responsible for the smell of vinegar. Carboxyiic acids! are easily 
converted to esters (R-CO-0-R'ï which have characteristically 
pleasing odors. 


Several functional groups are classified as derivatives of 
carboxyiic acids. These include esters, amides, acid chlorides/ and 
acid anhydrides. 


Esters are guite common in nature and are easily identified by 
their unigue odors. Isoamyl acetate has a banana smell, while methyl 
butyrate smells like apples. Ethyl butyrate is very structurally 
similar to methyl butyrate, but smells like pineapples instead of 
apples* Three/other : esters with distinct odors are isobutyl 
propionate, n-propyl acetate, and octyl acetate which smell like rum, 
pears, and oranges respectively. 


Aspirin (acetylsalicylic acid), an ester, is perhaps the most 
widely Used drug in the world. Since IS99, when aspirin was 
introduced as a mild analgesic and antipyretic, it has become the 
layman's first line of defense against west minordiécomforts such as 
colds and headaches, it has four helpful effects: 


1. Analgesic. It relieves pain rapidly, inexpensively, and 
effectively* 
2. Antipyretic. This means it brings down fever by increasing 
sweating and the flew of blood near the skin'3surface. 
3>'": Ant irheumatic7 : it /reduces the in f 1 arama t i on of and pá ili in 
the Joints, permitting mobility. 

Organic Laboratory Page 3 



4. Uricosuric, it decreases the deposits of urate that form in 
the joints. 


T'y-"77;/7, 
C-OH 


ij^ Aspirin 
O-C-CH


•••". T":' 


:;0:7 


The reaction between carboxyiic acide and alcohols is a simple 
way to prepare esters. An example reaction is shown below* Water is 
lost to yield the ester. 


•0^7A'7;/: ^ 
CH3-C-OH + HOCH3 -^y* CH5-C-OCH3 + HaO 

Amides• Acetaminophen is an amide more commonly known by the 
trade names Tylenol* and Datril*. It has important pain relief and 
fever reduction properties, but does not relieve inflammation as does 
aspirin. Its structure is¡ 


One of the many uses of acid ehleridee (and in general any acid 
halide) is in condensation reactions with amines to form polymers. 
During the polymer demonstration your instructor will make nylon 
using an acid chloride (adipoyl chloride) and an amine (hexamethylene 
diamine)* The condensation reaction has the general form: 


Q 6 

'••'•:• 77IÍ7 H 
yR^C-CJt + HO-R' ••' 7 > R-C-OR* 4 HCl 
-


The condensatio n polymerisatio n o f nylo n is : 

':::H/;': 7"H-••
Or. • '-•' 7 ¡ "• -Y yy-, 


Cl-C-CGH^-C-Cl + HîM-(CHa)6-Htt2 ^ > -(-C-(CH2>4-C-S-(CH3)fl-K-}n - + HCl 

••J. _ 


V tt-yy ' ;/•.'.••'•' "•' 7'.'.. " :•. I l \ '.:•• 7 

yyiïy', y&y 'y-'-yy 7 7 : '.'Q-'.= yO,.7. -yy 

• Jlpoy l thl'Jtrlil * h*X4ni*> t ly I ib e nylo n 6-6 
• •••>•'. .""•' J í i n i a E polyamide ' < :.' • y 

"yyi):•'•:"' y-- Organic Laboratory Page 4 7-y 77 -77 


!••..!: . f>1> y. j 


yyy [yy /• íyy*??y. yyj.yy<y 

EXPERIMENTAL 195 


TAv Demonstrationt Retoñe Préparation Prom an Alcohol 


1* Your instructor will demonstrate the oxidation of isopropyl 
alcohol to acetone using potassium dichromate in sulfuric acid. The 
unbalanced reaction is: 


OH 


: ' y '•;*•••'•••• ''s * 

H_c^c-CH; + cr„o H* 


2~7 CH3-C^CH3 +. Cr** + H20 
y I ' 

H 


2. What color change did you see? What change in oxidation 
number is this color change due to? 
3. Write the balanced equation using the redox half-reaction 
method. You will need to know that the central carbon atom in 
isopropyl alcohol has an oxidation state of 0 (zero), while the same 
carbon in acetone has an oxidation state of +2, 
. 

. • • 


7 


7 


Procédures 


1. Synthesis of Methyl Salicylate: Condensation of a 
carboxyli^ acid with an alcohol. 
a* Place a small amount (about the size of an aspirin 
tablet) Of .salicylic acid in a large test tube. Add 5 ml* of methanol' 
7(CAUTIONS methanol is flammable) and two drops of concentrated 
sulfuric acid. .:• ' y v.77yy z>c 


b. Heat in a water bath set up by your instructor for 
three to five minutes* For safety considerations, no flame should be 
used to warm the reactants. 
c. Carefully smell the product. If needed, the odor may 
be made more apparent by pouring the product onto 25 mL of ice* 
Describe the smell. 
Laboratory Page. 5: 


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d. Complete the following reaction of methanol and 
salicylic acid: 


1. •' •-. 
•v6-\: 
M^ 7 

H * ••••:.:: 

:' •• ;

L II 77:.••* 01,0»-+> 
^^ N0H 


"7 i • .'::.'" 

• • • •—^ ^ 

2. Synthesis of Isoamyl Acetate: 
a* Place three drops of isoamyl alcohol (CAUTION: 
flammable}* two drops of glacial acetic acid, and one drop of 3*0 M 
sulfuric acid into a test tube. (CAUTION: These are concentrated 
acids, avoid skin contact) 


b. Carefully heat the test tube in a hot water bath for 
three to five minutes. (CAUTION: Keep the flammable solution away 
from flame. Hot sulfuric acid is extremely corrosive.) 
c. Carefully smell the product. Record the smell below; 
d. Complete the following reaction for the synthesis you 
have just done: 


3. Synthesis of n-Propyl Acetate: 
a. Add one mL of n-propanol (cH3CH?CH2OH) to one mL acetic 
acid in a test tube. Add three drops of 3.0 M sulfuric acid to this 
mixture. yy :>y 
b. Heat the solution in boiling water for three to five 
minutes. (cAUTidtf: Keep the flammable solution away from the flame. 
¡ 

sulfuric acid is extremely corrosive.) 
7y' 


y 

Organic Laboratory Page 6 


yyyyy-: .. ;, A y / y y/-." .7 -/ 


çv .Carefully; sniff the products. Record the smell below:: 


, d. •/;•'. Wr ité the Condensât ion réaction bétwééri àcet le acid and 
;h-propanbi^ in>the7b^ 


4, / Tolléhs,: .Test ;; 'péteàtiéh - of Aldehydes 


a. Add;:• One mL of $\ si Iyer nitrate s oluti on ( CAUTION: Do 
•not get ¡silver nitrate on your skiny to a clean test tube. Add one 
drop of 1.0 M Sodium hydroxide tó the silver nitrate* A brbwut 
precipitate of silver oxide should form. 
b. Write the react ion ; ifór the: format ion of th is;tar own 
precipitatei 7 
c. New add three drops Of 6.0 K ammonium hydroxide to the 
tube. Swirl the test tube (careful not to spill any solution) until 
iai1 ybif ; the pr éc i pi t ate i s dissolved. ; Th lé is/ the To 1 lens Reagent .A 
d. Perform the Tollens test for an aldehyde by adding one 
drop of formalin to the Tollens Reagent* Describe what happened: 
7é7 When Tollen* Reagent: reacts with formalin, ah 
oxidation-reduction occurs* 


What is/it he reduc tion, ha 1 ¿-react ion? 


What is.the Oxidation half-reaction? 


Laboratory Page7 


•• ••.y^ • •' , 

•:'..• "..-• • 

198 


5. sniff Tests of Other Organic Compounds: 
a. carefully smell samples of cinnamic aldehyde, 
benzaldehyde, and methyl benzoate* Describe the smell of each 
compound below. Try to relate the smell to a naturally occurring 
substance. 
(1) Cinnamic aldehyde. 
• • .• 


. •


(2} Benzaldeyde. 


(3) Methyl benBoate. • 
• • " 


b. Complete the following reaction for the synthesis of 
isobutyl propionate. 
. 


:' • . 


CH, O • y 

1 • .:• 
"'• " • • • ' . •
HjC-Cn-CH^-QH + BO-C-CHJ-CHJ ^ 


• . 
• . 
• 


• 


(1) Carefully sniff a sample of isobutyl propionate. 
Record your interpretation of the odor below: 
• . :-•.•• 
• . 
• 


. 


c. Complete the following reaction for the synthesis of 
ithyl butyrates 
.'•••-• •-.. 

• •.. '. '

, - y : -. / .' 

....-• .".. -• 

". y •'••• ' ••' 

. ' ' • • • . 

0^ 
•-• y-y'y y " ':- 7 .. 

: • . ' '• ' ' -' • . • . ' • / • -•' 

II . 
H3C-CHa-0H BO-C-CHj-CHj-CHg -]

•.+ y~» ' ..y A ' 
' •'. '"' 

v.
: 
'.'• 

• 
y 
.,7 •• y y • 

."•• ' • 

• • 7 ' •'•••• -•-• : 

Organic Laboratory Page 8 



(I) Name the alcohol and acid reactants: 
(2) Carefully sniff the odor of ethyl butyrate and 
describe it below: 
7 7 yyy.'.yy-: -.-y /.-A:/-' ../../ • :••• = y. ;. :/.. ••:..;• .••.•• •,•-.y-. . .. 

.-•-'". 


Organ 16 ;. Laboratory Page 9 



ORGANIC FREIiAB ZXERCISE 7 7 

Ivy Circle the functional groups;on the following molecules: 


HjN-(CHj)i-NMa CH a-CH2 -CH2-C0OH 
Hexamethylene diamine Butyric Acid 


• y-y-, o y y O /-O/' 
•; H •_'.*•'••. 
y yyyyy^y 

- (rC-C6»4^ç^tîTCHa-CH2 -0-} „• 
H7N*C ÍH4-C-O-CH a-CHi 

Mylar,


Ben ¡toca in * • : 

2. Refer to table 26*4 in your text to identify an ester that could 
be associated with the following corresponding odor: 
¿89» 1SXSE 


Bananas 


Rua 
Pineapple 


3. Draw the condensation polymerization of nylon from hexametnyiene 
diamine and adipoyl chloride: 
A-LrJt oowmmit T MHTIMG smct i (*•*' i m^ai/HM* 

-.-/-:--••.. >. 

Organic Page 10 

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Títlo : Che mira ¡Is, chemistry 131 manual 
fslM99Q 
Author: 
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