Download LABORATORY MANUAL CHEMISTRY 121

LABORATORY MANUAL
CHEMISTRY 121
FIFTH EDITION
2007
Dr. Steven Fawl
LABORATORY MANUAL
CHEMISTRY 121
FIFTH EDITION
Dr. Steven Fawl
Science, Mathematics, and Engineering Division
Napa Valley College
Napa, California
NAPA VALLEY COLLEGE
COURSE: Chemistry 121
INSTRUCTOR: Dr. Steven Fawl, Room 1843, 253-3195
LECTURES: TW
LAB DEMO: Mon.
TESTS:
Th.
LABS:
MT
OFFICE HRS: MTW
11:00 to 12:20
11:00 to 11:50
11:00 to 11:50
1:30 to 5:20
12:20 to 1:20 Room 1843
EXAM DATES:
The following are tentative dates for your exams and the material each exam will cover.
Exam #1 - Thursday, February 25th - Kinetics
Exam #2 - Thursday, March 25th - Equilibrium
Exam #3 - Thursday, May 13th - Thermodynamics
Final Exam - Comprehensive - Wednesday, May 26th, 10:30-12:30
DESCRIPTION: A continuation of CHEM 120. Topics include solutions, acid-base and redox
equilibria, thermodynamics, kinetics, pH, buffers, solubility product, complex ions, thermodynamics,
electrochemistry, biochemistry and nuclear chemistry.
COURSE CONTENT:
1. Chemical Kinetics
Rate Laws
Activation Energy
Mechanisms
Catalysis
2. Chemical Equilibrium
LeChatlier's Principle
Homogenous Systems
Heterogeneous Systems
3. Acids & Bases
Strong and Weak acids
pH
Buffers
Titration Curves
4. Applications of Aqueous Equilibria
Solubility
pH controlled Solubility
Complex Ions
Amphoterism
5. Spontaniety, Entropy and Free Energy
Effect of Temperature
Work and Efficiency
6. Electrochemistry
Nernst Equation
Standard State Potentials
7. Radioactivity
Nuclear Stability
Half Life
Nuclear Transformations
8. Organic Chemistry
Nomenclature
Functional Groups
Free-Radical Halogenation
Substitution and Elimination Reactions
COURSE OBJECTIVES:
1.
2.
3.
4.
5.
6.
7.
8.
Explain the development of chemical principles and concepts based on experiments.
Analyze and solve complex or extended problems involving mathematical skills as well
as an ability to place these problems in an environment, biological, economic or social
context.
Design a laboratory experiment by defining the problem, collecting data, obtaining
results, deriving conclusions, and preparing a report to communicate the information to
others in writing.
Explain the concepts related to rates of reaction, activation energies, mechanisms of
reactions, as applied to the kinetic molecular theory.
Relate equilibrium information from chemical systems to the free energy, enthalpy and
entropy.
Determine the equilibrium constants and show how the spontaneity of the system is
related to the driving force of the reaction.
Apply the equilibrium system concepts to acid/base, solubility, redox, and complex ion
formation reaction systems.
Indicate how an electrochemical cell can be used to establish the standard free energy of a
chemical reaction, and measure the pH of a system.
STUDENT LEARNING OUTCOMES
1.
2.
3.
Communicate chemical and physical processes at the molecular level and how they relate to the
macroscopic environment.
Solve both qualitative and quantitative chemistry problems while demonstrating the reasoning
clearly and completely.
Implement laboratory techniques correctly using appropriate safety procedures and express them
clearly in written laboratory reports.
GRADING POLICY: Three exams and a final, quizzes, plus laboratory scores will count toward
the final grade according to the following schedule,
3 Exams = 300pts (100pts each)
Final = 200pts
Quizes = 100pts (10 @ 10pts each)
Lab
= 120pts
Total = 720pts
Grading is based on the class average = B-. The approximate breakdown of grades is,
100-85% A / 84-70% B / 69-60% C / 59-50% D / <50% F
ALL of the labs must be completed to pass the course regardless of the overall performance of the
student or else an "F" will be given.
LABS ARE CONSIDERED LATE IF THEY ARE TURNED IN ANY TIME AFTER THE
FRIDAY THAT THEY ARE DUE. LABS THAT ARE TWO WEEKS LATE WILL RECEIVE
NO CREDIT. Special arrangements must be made if a lab must be missed!
LAB EXPERIMENTS AND DATES
DATE
EXPERIMENT ONE . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2/1
Kinetics of the acid hydrolysis of trans-[Co(en)2Cl2]Cl
EXPERIMENT TWO . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2/8
Determination of the t½ of a Radioactive Isotope
HOLIDAY - Presidents Day . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2/15
EXPERIMENT THREE . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2/22
Chemical equilibria and Le Chatelier's principle
EXPERIMENT FOUR . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3/1
Hydrogen ion concentration and pH of aqueous solutions
EXPERIMENT FIVE . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3/8
Neutralization and hydrolysis
EXPERIMENT SIX . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3/15
Carbonic acid and its salts
EXPERIMENT SEVEN . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3/22
Titration Curve for KHP and Determination of pKa1 and pKa2
HOLIDAY - Spring Break . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3/29
EXPERIMENT EIGHT . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4/5
Determination of The Equilibrium Constant For FeSCN2+
EXPERIMENT NINE . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4/12
Determination of the Heat of Reaction
EXPERIMENT TEN . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4/19
Hess’ Law - Heats of Solution
EXPERIMENT ELEVEN . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4/26
Electrochemical cells
EXPERIMENT TWELVE . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5/3
Mystery Experiment
LABORATORY
There will be twelve lab assignments, approximately one per week. Attendance in lab the first
week of each session is necessary. Safety goggles are mandatory in the laboratory and can be
purchased in the campus bookstore. Lab Reports are due on Friday the week after the experiment
is done.
***** THERE IS A SUBSTANTIAL PENALTY FOR LATE LABS ****
***** HALF POINTS IF YOU ARE LATE UP TO ONE WEEK *****
****NO POINTS IF YOU ARE LATE BY MORE THAN ONE WEEK ****
Special arrangements must be made if a lab must be missed
LAB DRESS- GOGGLES MUST BE WORN AT ALL TIMES! Closed toe shoes are required.
Wear clothes that you are not afraid to ruin! Acids and other caustic solutions may cause holes in
your clothing if you are not careful. All goggles must meet with the approval of the instructor
(goggles bought from the bookstore are approved). Regular prescription glasses cannot be worn
instead of safety goggles.
Following is a description of how to do a lab report for this class. It is suggested that you re-read
this information on a regular basis, and that you follow this outline exactly.
SAFETY AND TECHNIQUE
RULES
Safety in the laboratory is extremely important. It is expected that you know laboratory safety
rules. It is important that if you feel uncomfortable with your knowledge of these rules that you
take the time to learn them. The following list is NOT complete. The Media Center, room 1028,
there is a video tape available in which a Napa Valley College instructor explains in detail safety
and technique rules. There is NO excuse for not following safety rules.
1)
Be attentive to instructions and follow them carefully. Read the board at the back of the
class room when you first come to class, any changes in procedure will be written there.
2)
If you ever have any questions about the procedure, apparatus, or chemicals it is important
that you ask the Instructor or Instructional assistant.
3)
Do not perform any unauthorized experiments. Anyone found doing so faces permanent
expulsion from class.
4)
Do not handle chemicals or materials not assigned to you or not called for in the experiment.
5)
Learn the location and proper use of the fire extinguisher, safety shower, eye and face wash.
Keep the first aide sink area clear at all times
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6)
Coats, books, etc, should be kept in the space provided for them at the back of the lab.
Many of the chemicals used in the lab can ruin or stain paper and clothing.
7)
Never taste chemicals, nor pipet by mouth. Always use pipet bulbs or wheels.
8)
Smell chemicals by fanning a little vapor towards you.
9)
Experiments in which dangerous or obnoxious fumes are produced must be done in the fume
hood. Be sure to stop these reactions as soon as possible.
10) No eating, drinking or smoking in the lab.
11) Never point test tubes at yourself or others.
12) In the event of any injury, spill or glass breakage inform the Instructor immediately.
13) Goggles must be worn at all times when in the lab.
14) Chemicals may not be taken out of the lab (not even to the I.A.'s desk.)
15) Chemicals may not be stored in lockers.
16) Avoid unnecessary contact with ALL chemicals.
17) Do not leave lit burners unattended.
18) Every time you use a chemical read it's label carefully. If any discrepancies inform the IA or
instructor immediately.
19) All containers which contain a chemical or in which a reaction occurs must be labeled.
20) When labeling a storage container include name and/or formula of chemical, any appropriate
warnings, concentration, date and your name.
21) NEVER place anything inside a reagent bottle, no spatulas, droppers, nor pipets. If the
reagent is a clumpy solid inform the IA. Proper technique is to "roll" containers from side to
side to remove solids and to pour liquids into smaller containers (such as a beaker) first.
22) NEVER return unused chemical (liquids or solids) back to the original container - offer
excess to another student or dispose of it appropriately.
8
23) Be conservative of reagents, place only the amount you need into a labeled container (such
as a beaker). Do not take the reagent bottles to your work area - leave them where every one
can find them.
24) Use tap water to wash glassware - you should rinse with DI water - please be conservative.
25) To dilute acids and bases, Add the Acid (or Base) to the Water.
26) Dispose of liquids and solids appropriately, read the board, or your experimental procedure
for special instructions, otherwise dispose of liquids and soluble solids down the sink with
lots of water, insoluble materials (such as paper towels) should be put in the waste basket.
KEEP THE SINKS CLEAN
27) It is very important to keep the lab clean. Before you leave each time be sure to:
a) return equipment to its proper place
b) clean up your workspace with the sponge
c) put away your labware
d) lock your locker
There is NO reason for a messy lab. Everything you need to keep your lab neat and clean is
available. Dirty counters, paper towels left in the sink or troughs, labware left out, messes
left under the fume hood, chemical spills left on the balance, are BAD technique and as such
will not be tolerated.
28) You may not be in the laboratory at any time other than your scheduled laboratory period
unless you have the permission of the instructor in charge as well as your course instructor.
Do not visit friends during their lab time and do not invite your friends or family to visit
you.
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EXPERIMENT ONE
KINETICS OF THE ACID HYDROLYSIS OF trans-[Co(en)2Cl2]Cl
In aqueous solution the green complex trans-dichlorbis (ethylenediamine) cobalt(III) chloride
dissociates into chloride ion and trans-[Co(en)2Cl2]+. This cobalt complex ion then reacts in acid
solution to yield a mixture of the cis and trans forms of [Co(en)2(H2O)Cl]2+. The progress of this
reaction can be followed visually because the products are red while the reactant is green. The
purpose of this exercise is to determine the rate law for the hydrolysis of trans-[Co(en)2Cl2]+ and
to determine the activation energy for the hydrolysis reaction by carrying out the reaction at
several different temperatures from 45C to 85C.
The equation for the hydrolysis reaction is,
trans-[Co(en)2Cl2]+ + H2O  cis/trans-[Co(en)2(H2O)Cl]+ + Cl(dark green)
(burgundy to red violet)
The reactant is green while the product is supposed to be burgundy red to violet, but because of
problems we have had in recent years, the product is sometimes yellow or orange. To determine
the rate law and rate constants for this reaction, we shall measure the half-life for each
experiment. When 50% of the reactant has been converted to product, the mixture (50% green
and 50% red) has a characteristic color best described as "gun-metal gray", but other colors are
possible and your instructor will inform you of them..
EXPERIMENT: You may work in pairs. Weigh out two samples (one about 0.12 g and one 0.02
g) of trans-[Co(en)2Cl2]Cl, and dissolve them in separate test tubes each containing 8.0 mL of icecold 1 M sulfuric acid, then place these and all following solutions in an ice bath. With a pipet
transfer exactly 2 mL (save the rest for the next part) from each of these solutions to separate test
tubes, and heat these in boiling water for 5 minutes to completely hydrolyze the complex to
trans-[Co(en)2(H2O)Cl]+. The solutions should be red, but may be yellow or orange. Cool both
the tubes to room temperature, and add 2 mL of the corresponding green unhydrolyzed complex.
You should now have two pairs of solutions that are gun metal gray, or some other color. The
exact color is not important as long as you have a color representing the half-way point in the
reaction. Keep these tubes in an ice bath so that they do not undergo additional reaction during
the laboratory period. Adjust the temperature of a large beaker of water to 55 +/- 1C. One
student should place the tubes which contain the remaining unhydrolyzed green complex into the
hot water while a second student records the time. As the solutions change color a comparison
should be made to our standards prepared previously. When the colors match, the time and
temperature should be recorded. From this you should be able to determine the order of the
reaction and the rate law.
Prepare a solution of 0.15 g of trans-[Co(en)2Cl2]Cl in 30 mL of 1 M H2SO4, transfer 2 mL to a
test tube and boil it for 5 minutes. As before, cool it and add 2 mL of unreacted solution to make
4 mL of half converted standard. Using 4 mL portions measure the half life of the reaction at 85,
10
75, 65, and 45C. Record the times and temperatures in your notebook. Be sure to record t1/2, K,
ln K and 1/T in your data table.
CALCULATIONS
(You may use this sheet in your lab report)
1) Write down the rate law for the acid hydrolysis of trans- [Co(en)2Cl2]Cl. What is the order of
the reaction?
2) The relationship between the rate constant k and the half-life t1/2 for first and second order
reactions is given by,
k = 0.693/t1/2 (First order)
or
k = 1/([A]t1/2) (Second order)
Using the order obtained in (1) and the proper half-life equation calculate k at each temperature of
your reaction. Show one example here, include all the k's in your data table.
3) Make a plot of ln(k) vs. 1/T, where k is the rate constant at each temperature and T is the
temperature in Kelvin.
4) The relationship between the rate constant and a quantity called the activation energy is given
by the Arrhenius equation,
ln(k) = -Ea/RT + ln(A)
Where R = 8.314 J/mol-K, A is called a pre-exponential term (a constant), Ea is the activation
energy, and k is the rate constant. This equation is of the general form
y = mx + b
where m is the slope and b is the intercept of a line. Your plot of ln(k) vs. 1/T should yield a
slope of -Ea/R and an intercept of ln(A). From the slope and intercept calculate Ea and A. Ea and
A are your results.
11
PROBLEMS
1) In this experiment you assumed that the t1/2 was very slow at ice bath temperatures of about
0C. Use your data to calculate the t1/2 at this temperature. Was the assumption valid?
2) While doing a different experiment a student found that the half-life of his reaction (not this
experiment) decreased as the concentration of his reacting species increased. What if anything
was wrong?
3) It takes 3 minutes to soft-boil an egg at 100C. Assuming a soft boiled egg represents the
half-life of the denaturation of egg albumin, comment on why most people add salt to the water
when boiling an egg. (What does the salt do to the water?
12
Name ______________________
Date _______________
EXPERIMENT ONE
KINETICS OF THE ACID HYDROLYSIS OF trans-[Co(en)2Cl2]Cl
DATA
Concentration
Time
Rate Law =
Temperature
Time
Rate Const. K
ln K
1/Temp
Attach Graph
Slope =_____________________________________
Ea = __________________________
Intercept = _________________________________
A = ___________________________
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EXPERIMENT TWO
DETERMINING THE HALF-LIFE OF AN ISOTOPE
DISCUSSION
One type of nuclear reaction is called radioactive decay, in which an unstable isotope of an
element changes spontaneously and emits radiation. The mathematical description of this process
is shown below.
Af = Aie-kt
In this equation, k is the decay constant, commonly measured in sec-1 (or another appropriate unit
of reciprocal time) similar to the rate law constant, k, in kinetics analyses. Ai is the activity (rate
of decay) at t = 0. The SI unit of activity is the bequerel (Bq), defined as one decay per second.
This equation shows that radioactive decay is a first-order kinetic process.
One important measure of the rate at which a radioactive substance decays is called half-life, or
t1/2. Half-life is the amount of time needed for one half of a given quantity of a substance to decay.
Half-lives as short as 10-6 second and as long as 109 years are known.
In this experiment, you will use a source called an isogenerator to produce a sample of radioactive
barium. The isogenerator contains cesium-137, which decays to produce barium-137. The newly
made barium nucleus is initially in a long-lived excited state, which eventually decays by emitting
a gamma photon and becomes stable. By measuring the decay of a sample of barium-137, you will
be able to calculate its half-life.
EXPERIMENT
Each group should have a computer interfaced to a radiation monitor and have the Vernier “Lifetime” program running. You will also be given a shallow aluminum cup that will hold your
radioactive sample. Place the radiation monitor on top of, or adjacent to, the cup to get a
maximum rate of sample detection.
Your instructor will deliver a small sample of radioactive Barium-137 to each group. As soon as
delivery is complete, click the
button in the upper right corner of the data window.
Collect data for thirty minutes. Do not move the radiation monitor or the cup during the data
collection. Be careful not to spill the solution in the cup.
When the data collection is complete, you may dispose of the barium solution by pouring it down
the sink.
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DATA TABLE
Experimental Data Fit to Af = Ai exp(- kt) + B
Ai =
k=
B=
t1/2 =
CRC t1/2 for 137 Ba =
CALCULATIONS
The solution you obtained from the isogenerator may contain a small amount of long-lived cesium
in addition to the barium. To account for the counts due to any cesium, as well as for counts due
to cosmic rays and other background radiation, you can determine the background count rate from
your data. By taking data for 30 minutes, the count rate should have gone down to a nearly
constant value, aside from normal statistical fluctuations. The counts during each interval in the
last five minutes should be nearly the same as for the 20 to 25 minute interval. If so, you can use
the average rate at the end of data collection to correct for the counts not due to barium.
BACKGROUND RADIATION
Select the data on the graph between 25 and 30 minutes by dragging across the region with your
mouse. Click on the statistics button on the toolbar. Read the average counts during the intervals
from the floating box, and record the value in your data table as the average background counts.
Use the corrected count rates to derive an exponential function for the first fifteen minutes of the
data collection.
DETERMINING VALUES FOR k and Ai
Click and drag the computer mouse across the region between 0 and 15 minutes. Make sure that
all of the data points in this region are in the shaded area. Click the curve-fit icon on the tool bar.
Select natural exponent from the equation list. In the "Coefficients" box, type in the background
radiation value as the "B" value. This will account for the background radiation inherent in the
experiment. Click Try Fit . The function will be shown and plotted along with your data. Click
to see a full graph of the function and data. Record the fit parameters Ai, k, and B in your data
table.
15
Print a copy of your graph and the table of data.
From the fit parameters, the half-life t1/2 for 137Ba and place it in your data table
PROBLEMS
1) What fraction of the initial activity of your barium sample would remain after 25 minutes?
2) Was it a good assumption that the counts in the last five minutes would be due entirely to
non-barium sources?
3) Would any of your t1/2 or k values change if you had been given more or less sample? Explain.
16
EXPERIMENT THREE
CHEMICAL EQUILIBRIA AND LE CHATELIER'S PRINCIPLE
All chemical reactions proceed toward an equilibrium position, some more rapidly than others. In
this experiment we shall consider only reactions which occur fairly rapidly, reaching equilibria in
a few minutes or less. Such reactions are said to be rapid and reversible.
The equilibrium position varies for different reactions. For example acetic acid dissociates only a
small extent.
CH3CO2H  H+ + CH3COOAt equilibrium in 0.1 M acetic acid, only about 1% of the acetic acid molecules are dissociated
into hydrogen ions and acetate ions. We say that this equilibrium "lies far to the left."
The equilibrium position is the same whether one starts with reactants or products. For example,
when equivalent amounts of H+ and CH3CO2- are mixed, they combine to a large extent to form
CH3COOH molecules, and at equilibrium only about 1% of the H+ and CH3COO- remain
unreacted. The final concentrations of H+, CH3COO-, and CH3CO2H are exactly the same as in a
solution of acetic acid of the same concentration.
LE CHATELIER'S PRINCIPLE
The equilibrium position may be shifted to the left or to the right by changing experimental
conditions such as concentration, pressure or temperature. A very useful chemical principle,
called Le Chatelier's Principle. enables one to predict in which direction the equilibrium will
shift. The principle states that;
"If a change is made in any of the factors influencing a system at equilibrium, reaction
will occur in the direction which tends to counteract the change made."
To illustrate this principle, we shall examine the effect of increasing the concentration of CH3CO2on the equilibrium of the reaction above. If this change causes the equilibrium to shift to the left,
then the concentration of H+ will decrease; if the equilibrium shifts to the right, then the
concentration of H+ will increase.
EQUILIBRIUM SHIFT WHEN A CONCENTRATION IS CHANGED
EXPERIMENT: Place a 5 mL portion of 1M acetic acid in each of two test tubes, a 5 mL portion
of 0.1 M acetic acid in a third test tube, and 5 mL of de-ionized water in a fourth. Add 2 drops of
methyl orange indicator to each test tube and record in your laboratory notebook the color of the
contents of each test tube. To one of the test tubes containing 1 M acetic acid, add 2 mL of 4 M
sodium acetate (CH3CO2Na) dropwise. Observe and record the changes in color that occur.
17
EQUILIBRIUM BETWEEN A SOLID SALT AND A SOLUTION
As a second example of a rapid reversible reaction we shall investigate the equilibrium between a
solid salt and a saturated solution of the salt. The solubility of solid silver acetate (CH3CO2Ag) is
0.06 mole per liter at room temperature, and the solubility product is therefore (0.06)2 or 3.6 x10-3.
CH3CO2Ag (s)  Ag+ + CH3CO2- Ksp = 3.6 x 10-3
EXPERIMENT: Prepare some solid silver acetate as follows: Place 5 mL of 0.1 M silver nitrate
(AgNO3) in a centrifuge tube and add 2 mL of 4 M sodium acetate. Stir the mixture thoroughly
using a glass rod. Record the color of the precipitate that is formed. Place the centrifuge tube in
the centrifuge and balance the centrifuge by placing a second centrifuge tube containing 7 mL of
water opposite the first tube. Turn on the centrifuge for about 30 secs. When the centrifuge has
stopped, remove the tubes then decant the liquid and discard it.
To the solid which remains, add 3 mL of de-ionized water. Place the tube in a beaker of boiling
water and note the amount of solid in the tube. Cool the tube to room temperature by placing it in
a beaker of tap water for 5 minutes. Again note the amount of solid. Centrifuge the mixture as
above and decant the liquid into a clean test tube for use in another experiment.
To the remaining solid add 2 mL 6 M HNO3. Stir with a glass rod until all of the solid dissolves.
THE COMMON ION EFFECT
EXPERIMENT: To the test tube containing the saturated silver acetate solution add 2 mL of 4
M sodium acetate. Stir the mixture with a glass rod. Observe any reaction that occurs. Write a
net reaction to account for the observed change. In your conclusion state briefly how this
experiment illustrates Le Chatelier's Principle. This application of Le Chatelier's principle is
called the common ion effect.
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Name ______________________
Date _______________
EXPERIMENT THREE
LeChatelier’s Principle
Data
Equilibrium Shift When Concentration is Changed
Solution
Color w/ Methyl Orange
1.0 M Hac
0.1 M HAc
Pure Water
Volume of NaAc added to
HAc
Color
0.0 mL
0.5 mL
1.0 mL
1.5 mL
2.0 mL
Equilibrium between Solid and Solution & Common Ion Effect
Procedure
Result (amount)
AgNO3 + NaAc Before Boiling
AgNO3 + NaAc After Washing and Boiling
AgAc(s) + 2 mL of HNO3
AgAc solution + 2 mL NaAc
Net Ionic Reaction =
19
PROBLEMS
1) What happens to the H+ concentration in acetic acid when sodium acetate is added? What is
the reaction taking place to account for this? Will the [H+] be higher or lower after the sodium
acetate/acetic acid system has come to equilibrium?
2) What substance disappeared when HNO3 was added to the wet solid formed by adding AgNO3
to sodium acetate? Write the reaction that has occurred (net ionic equation).
3) How would the solubility of AgCl be affected by addition of 6 M HNO3? Why do AgCl and
AgAc behave differently with addition of H+?
4) How would the equilibrium have shifted in the common ion experiment upon addition of Silver
Nitrate? of Sodium Nitrate?
20
5) Excess solid Ca(OH)2 is in equilibrium with its ions,
Ca(OH)2  Ca2+ + 2 OHState whether the amount of Ca(OH)2 increases, decreases, or stays the same upon addition of;
NaOH
KNO3
HCl
Ca(OH)2
CaCl2
6) Does the solubility of solid Ca(OH)2 in water depend on the amount of solid in contact with the
saturated solution?
7) The solubility of Ca(OH)2 in water is 0.165 g per 100 mL of water. Calculate the solubility
product of Ca(OH)2.
21
EXPERIMENT FOUR
HYDROGEN ION CONCENTRATION AND pH OF AQUEOUS SOLUTIONS;
DISSOCIATION CONSTANTS OF ACETIC ACID
AND AMMONIA SOLUTIONS
Water is a weak electrolyte. It dissociates slightly to give H+ and OH-.
H2O  H+(aq) + OH-(aq)
In pure water [H+] = [OH-] = 1.0 x 10-7 M and thus the amount of dissociation is extremely small.
The dissociation constant for water can be written,
Kw = [H+] [OH-] / [H2O] = 1.0 x 10-14 @ 25C
Chemists use activities in the formulation of Kw instead of concentrations. By convention the
activity of water is set equal to one and the activity of dilute solutions is set equal to the
concentrations of the ion. Therefore our dissociation constant can be written as,
Kw = [H+] [OH-] = 1.0 x 10-14
We can use the dissociation constant of water to calculate the concentration of H+ or of OH- in any
aqueous solution if the concentration of the other ion is known.
THE USE OF pH TO MEASURE ACIDITY
It is generally easier to express [H+] in terms of pH. We shall define pH = -log[H+]. Thus the pH
of pure water would be pH = -log[H+] = -log[10-7] = 7.0. Acidic solutions have a pH less than 7.0
and basic solutions have a pH larger than 7.0. THE LOWER THE pH THE HIGHER THE
ACIDITY.
DETERMINATION OF HYDROGEN ION CONCENTRATION AND pH
In this section you will determine the [H+] and pH of two unknown aqueous solutions. In the back
of your LabManual you will find a table which will tell you the relationship between color and pH
for several different indicators. While the data in this table are useful as a general guide, it should
be realized that different individuals respond differently to the same color, and they may also use
different words to describe a particular color. For the accurate determination of pH with an
indicator, one must always make a direct comparison of the color of the unknown solution with
the colors of solutions of known pH values. Furthermore, the solutions being compared should be
as similar as possible in factors which may affect the color such as size of test tube, volume of
solution, and number of drops of indicator. It is usually helpful to observe the colors by placing
the test tubes side by side on a piece of white paper, and then looking down through the solutions
from the top.
22
EXPERIMENT: Obtain your unknowns from the instructional assistant. Test 2mL your first
unknown with 2 drops of bromothymol blue to ascertain whether it is acidic or basic. Then
proceed with the appropriate indicators to determine the approximate pH. Record the name of
each indicator
used, its color with your unknown and the pH range indicated into your data table. For these tests
use 2 mL of your unknown solution and 2 drops of indicator.
When you have decided the approximate pH, you will need to verify it. To do this you must
prepare a solution of this pH and test it with the appropriate indicators. If your solution is in the 1
M to 10-3 M H+ concentration range (0 to 3 pH), prepare it from 6 M HCl. Solutions in the range
from 10-4 to lO-11 M H+ (pH 4 to 11) are available as buffers from the instructional assistant.
Solutions in the 10-11 to 10-14 M range of H+ concentration range (pH 11 to 14 ) are prepared from
6 M NaOH. Use serial dilution to prepare solutions from HCl and NaOH. For these final critical
comparisons use 5 mL of both the standard and the unknown (in different test tubes) and 3-4
drops of indicator. Compare the color of the standard solution to the color of your unknown with
at least 2 and preferably 3 different indicators. If they are the same, then you know the pH and
[H+] of your unknown solution.
AMMONIA SOLUTIONS
Ammonia is a gas at room temperature that is very soluble in water. The resulting solution is basic
because of the reaction,
NH3(g) + H2O  NH4+(aq) + OH-(aq)
Kb = [NH4+] [OH-] / [NH3]
EXPERIMENT : Prepare a 1 M solution of NH3 from the 6 M NH3 found in the lab. Determine
the [OH-] in the 1 M NH3 by using the indicators alizarin yellow and indigo carmine. Record the
colors in your data table. Copy this table into your data section and complete it for 1 M NH3. This
table and all similar tables must be copied into your data section.
Substance
H+
OH-
NH3
NH4+
Equilibrium
Concentration
Using these data calculate the dissociation constant Kb for NH3. Include this in your results.
23
REACTIONS OF NH4+ AND OHPredict what will happen when solutions of ammonium chloride (NH4Cl), and sodium hydroxide
are mixed.
EXPERIMENT: Prepare 12 mL of 1 M NaOH by adding 2 mL of 6 M NaOH to 10 mL of water.
Mix thoroughly and then place 5 mL of the 1 M NaOH in each of two test tubes. Add several
drops of indigo carmine to each test tube. Record the color in your data table. Slowly add 5 ml of
1 M NH4Cl to one test tube and record the color changes that occur. Does the OH- concentration
increase or decrease as you add NH4Cl to the NaOH solution?
Write an equation for the net reaction that occurs. Include this rxn and all similar reactions in your
data section.
DISSOCIATION CONSTANT OF ACETIC ACID
In the following experiment we shall determine the dissociation constant of acetic acid by mixing
solutions containing equal numbers of moles of acetic acid and sodium acetate. The resulting
buffer solution will contain approximately equal concentrations of HAc and NaAc.
[HAc]  [Ac-]
It then follows from the law of chemical equilibrium that the concentration of H+ in the solution is
equal to the dissociation constant,
HAc  H+ + Ac-
Ka = [ H+ ] [ Ac- ] = [ H+ ]
[HAc]
Furthermore, if we define pKa = -logKa, then in this solution -logKa = -log[H+], or pKa = pH.
EXPERIMENT : Add 10 mL of 1 M HAc to 10 mL of 1 M NaAc. Determine the pH and [H+] of
the resulting solution using the appropriate indicators. Record the name of each indicator and its
color in this solution in your notebook. In your notebook, copy and complete the following table
for the solution (which does not contain 0.5 M HAc and 0. 5 M NaAc).
Substance
H+
OH-
HAc
Ac-
Na+
Equilibrium
Concentration
Using these data, calculate the dissociation constant Ka and pKa for acetic acid. Include this
number in your results.
24
PROBLEMS
1. A solution is prepared by dissolving 0.050 mole of HCl and 0.030 mole of NaOH in sufficient
water to give a final volume of 2.0 liters. Calculate the concentration of all the ionic (Na+ etc.)
species in the final solution.
2. Instead of determining the dissociation constant of ammonia by our experimental procedure we
could mix equal volumes of 1 M NH3, and 1 M NH4Cl. Show that for such a solution Kb = [OH-],
and pKb = 14 - pH.
3. Calculate the H+, and OH- of a 0.10 M NH3 solution. (Use Kb = 1.8 x 10 -5 in this and the
following problems.
25
4. A solution is prepared by mixing 0.050 mole of NH4Cl and 0.010 mole of NaOH in sufficient
water to give a final volume of 200 mL. Calculate the concentration of all the molecular and ionic
species in the resulting solution.
5. A solution is prepared by mixing 0.050 mole of NH3 and 0.20 mole of HCl in sufficient water
to give a final volume of 500 mL. Calculate the concentration of all the molecular and ionic
species in the resulting solution.
26
EXPERIMENT FIVE
NEUTRALIZATION AND HYDROLYSIS
The neutralization of an acid by a base to form a salt and water proceeds to an equilibrium
position which varies depending on the strengths of the acid and base that are used. In this
assignment we shall investigate the equilibrium positions reached for the reaction of;
A. a strong acid and a strong base,
H+ + OH-  H2O
B. a weak acid (HX) and a strong base,
HX + OH-  X- + H2O
C. a strong acid and a weak base,
H+ + BOH  B+ + H2O
or, in the special case of ammonia,
H+ + NH3  NH4+
In each of the above cases the same equilibrium position may be reached by mixing equivalent
amounts of the reactants (as in a titration) or by preparing an aqueous solution of the product (the
salt of the acid and base) of equivalent concentration. Since it is much easier experimentally to
prepare the salt solutions, we shall prepare them and measure their pH and [H+], from which we
can infer the equilibrium position. We shall therefore study the above reactions in the reverse
direction. The reverse of the neutralization reaction is called the "hydrolysis" reaction, because in
the typical case a water molecule is split by the reaction.
HYDROLYSIS OF THE SALT OF A STRONG ACID AND STRONG BASE
EXPERIMENT: Test about 5mL of a 50 mL portion of distilled water with bromothymol blue
indicator. It should give a green color. If it is yellow, boil the water until it gives a green color
with bromothymol blue. If your solution is blue inform the instructional assistant. Prepare an
approximately 1 M NaCl solution by dissolving about 1.2 grams of sodium chloride (compare
with the sample in the laboratory) in 20 mL of your distilled water. Determine the pH and [H+] of
your NaCl solution by using appropriate indicators, including bromothymol blue. Use the same
procedure you used in the preceding assignment for determining an unknown pH. (Record names
and colors of indicators in your laboratory notebook.) Copy in your laboratory notebook and
complete the following table for a 1 M NaCl solution.
Substance
H+
OH-
Na+
Equilibrium
Concentration
27
Cl-
HCl
NaOH
Because sodium ion is the ion of a strong base, there is no tendency for Na+ to combine with OH-,
nor for H+ to combine with Cl-, and therefore the following reactions do not occur.
Na+ + H2O  NaOH + H+
Cl- + H2O  HCl + OH(More precisely, the equilibrium positions of these two reactions are so far to the left that [NaOH]
= 0 and [HCl] = 0.
HYDROLYSIS OF THE SALT OF A WEAK ACID AND A STRONG BASE.
The neutralization reaction for acetic acid, CH3CO2H, by a strong base is
CH3CO2H + OH-  CH3CO2- + H2O
The hydrolysis of acetate ion, CH3CO2-, is represented by the reverse of the above reaction:
CH3CO2- + H2O  CH3CO2H + OHEXPERIMENT: Determine the pH and [H+] of 1 M CH3CO2Na by using appropriate indicators,
including thymol blue and phenolphthalein. (Record names and color of indicators in your
laboratory notebook.) From your measured value of the pH and Kw, calculate [H+] and [OH-] in 1
M CH3CO2Na. Using this value for [OH-], calculate [CH3CO2-]. Copy in your laboratory
notebook and complete the following table for 1 M CH3CO2Na.
Substance
H+
OH-
Ac-
HAc
Na+
NaOH
Equilibrium
Concentration
HYDROLYSIS OF THE SALT OF A STRONG ACID AND A WEAK BASE
The neutralization reaction for ammonia, NH3, by a strong acid is
NH3 + H+  NH4+
The "hydrolysis" reaction of ammonium ion, NH4+, is represented by the reverse of the above
reaction;
NH4+  NH3 + H+
EXPERIMENT: Determine the pH and [H+] of 1 M ammonium chloride by using appropriate
indicators, including bromocresol green and bromothymol blue. (Record names and colors of
indicators in your laboratory notebook.) From your measured value of the pH and Kw, calculate
28
[H+] and [OH-] in 1 M NH4Cl. Using this value for [H+], calculate [NH3] and [NH4+]. Copy this
into your laboratory notebook and complete the following table for 1 M NH4Cl.
H+
Substance
OH-
NH4+
Cl-
NH3
HCl
Equilibrium
Concentration
HYDROLYSIS OF SODIUM CARBONATE SOLUTIONS
The neutralization reaction for bicarbonate ion, HCO3- , by a strong base is
HCO3- + OH-  CO32- + H2O
The hydrolysis of carbonate ion, CO32- is represented by the reverse of the above reaction
CO32- + H2O  HCO3- + OHEXPERIMENT: Determine the pH and [H+] of 1 M sodium carbonate by using appropriate
indicators, including alizarin yellow and indigo carmine. (Record names and colors of indicators
in your laboratory notebook.) From your measured value of the pH and Kw, calculate [H+] and
[OH-] in 1 M Na2CO3. Using this value for [OH-], calculate [HCO3-] and [CO32-]. Copy in your
laboratory notebook and complete the following table for 1 M Na2CO3.
Substance
H+
OH-
Equilibrium
Concentration
29
Na+
CO32-
HCO3-
Name ___________________
Date_________________
NEUTRALIZATION AND HYDROLYSIS WORKSHEET
Data: Strong Acid and a Strong Base: Complete the following table for 1M NaCl
pH =
Substance
H+
OH-
Na+
Cl-
HCl
NaOH
Equilibrium
Concentration
Weak Acid and Strong Base: Complete the following table for 1 M CH3CO2Na.
pH =
Substance
H+
OH-
Ac-
Na+
NaOH
NH3
HCl
HAc
Equilibrium
Concentration
Weak Base and Strong Acid: Complete the following table for 1 M NH4Cl.
pH =
Substance
H+
OH-
NH4+
Cl-
Equilibrium
Concentration
Weak Diprotic Acid and Strong Base: Complete the following table for 1 M Na2CO3.
pH =
Substance
H+
OH-
Equilibrium
Concentration
30
Na+
CO32-
HCO3-
PROBLEMS
1) a) Is the [H+] of the NaCl solution greater than, less than or equal to that of your boiled distilled
water? b) Which indicator would you choose for the most exact endpoint for the titration of an
HCl solution with NaOH?
2) a) Would you expect a solution of sodium acetate to be acidic, basic or neutral? b) What is the
pH of sodium acetate? c) Which indicator would you choose for the titration of CH3CO2H with
NaOH?
3) a) Would you expect a solution of ammonium chloride to be acidic, basic or neutral? b) What
is the pH of 1 M NH4Cl? c) Which indicator would you choose for the titration of NH3 with HCl?
4) Would you expect a solution of sodium carbonate to be acidic, basic or neutral?
5) The Ka for acetic acid is 1.8 x 10-5, and the Kb for ammonia is also 1.8 x 10-5. Calculate the
equilibrium constant for the
(a) neutralization of a strong acid by a strong base,
(b) neutralization of acetic acid by sodium hydroxide,
(c) neutralization of ammonia by hydrochloric acid.
6) Ammonium acetate, NH4CH3CO2, is a strong electrolyte. When a 1 M solution of ammonium
acetate is tested with bromothymol blue, the solution is green. What substances are present in a 1
M NH4CH3CO2 solution, include all ions and molecules? Write an equation for the net reaction
for the hydrolysis of NH4CH3CO2. Calculate the equilibrium constant for this reaction, using the
information given in problem #5.
31
EXPERIMENT SIX
CARBONIC ACID AND ITS SALTS
Carbon dioxide is a colorless gas at room temperature which is somewhat soluble in water. Water
in equilibrium with gaseous CO2 at 1 atm pressure contains 0.034 M H2CO3.
CO2(1 atm) + H2O  H2CO3 K = 0.034
Solid carbon dioxide is a convenient source of carbon dioxide (dry ice), but in this lab you will
use a carbon dioxide generator to make saturated solutions of H2CO3. Carbonic acid is a typical
diprotic acid in that it dissociates stepwise as shown below,
H2CO3  H+ + HCO3-
Ka1 = 4.3 X 10-7
HCO3-  H+ + CO32-
Ka2 = 4.7 X 10-11
Thus we see that there are two possible salts, NaHCO3 (sodium bicarbonate) and Na2CO3 (sodium
carbonate). In the following experiments we will investigate the properties of H2CO3, NaHCO3,
and Na2CO3.
CARBONIC ACID SOLUTIONS
EXPERIMENT: Prepare about 100 mL of a saturated solution of carbon dioxide by adding a few
mL of concentrated H2SO4 to some solid sodium bicarbonate inside a CO2 generator found in the
lab. Allow the resulting gas to bubble through 100 mL of distilled water for about 5 minutes. Test
the resulting solution with bromocresol green and methyl orange indicators. Record the pH and
label the solution 0.034 M H2CO3. Note- the pH of the solution is due almost entirely to the first
acid dissociation.
EXPERIMENT: Bubble some CO2 through a 10 mL sample of 1M CaCl2. Does a precipitate
form? If we assume that the solution is saturated with CO2, what is the concentration of H2CO3 in
the solution? Calcium carbonate is very slightly soluble in water.
CaCO3  Ca2+ + CO32- Ksp = 4.7 x 10-9
Calculate the maximum concentration of CO32- that could exist in a 1 M CaCl2 solution. Is the
concentration of CO32- in your saturated H2CO3 solution larger or smaller than this value? Using
these data complete the following table. Use the value of Ka2 to calculate [CO32-].
32
Substance
H2CO3
H+
HCO3-
OH-
CO32-
Equilibrium
Concentration
SODIUM CARBONATE SOLUTIONS
Sodium carbonate, Na2CO3 is a solid that is very soluble in water and it is a strong electrolyte
which dissociates to give Na+ and CO32- in aqueous solutions. Sodium carbonate solutions can
hydrolyze stepwise in the following manner,
CO32- + H2O  HCO3- + OHHCO3- + H2O  H2CO3 + OHEXPERIMENT: Place 5 mL of 1 M Na2CO3 into a test tube and add 5 mL of 1 M CaCl2.
Centrifuge the mixture and note the amount of solid CaCO3 produced. Record your results. Write
an equation for the net reaction.
SODIUM BICARBONATE SOLUTIONS
Sodium bicarbonate, NaHCO3, is a solid that is very soluble in water and it is a strong electrolyte
which dissociates to give Na+ and HCO3- in aqueous solutions. Sodium bicarbonate solutions may
also be prepared by the reaction of solutions containing 1 mole of H2CO3 and 1 mole NaOH. The
equation for the net reaction is,
H2CO3 + OH-  HCO3- + H2O
Since HCO3- is a weak acid, it should dissociate to give H+ and CO32-. However the H+ produced
would react with another HCO3- to form the weak acid H2CO3. As a result we predict that the
principal equilibrium in a NaHCO3 solution would be,
2 HCO3-  H2CO3 + CO32- *
The position of this equilibrium is far to the left, the concentrations of H2CO3 and CO32- being
much smaller than the concentration of HCO3-.
EXPERIMENT: Place 5 mL of 1 M NaHCO3 and 5 mL of 1 M CaCl2 in a centrifuge tube.
Centrifuge the mixture and note the amount of precipitate obtained. Decant the solution into a
second centrifuge tube. Now, place the tube containing the decanted solution in a beaker of
boiling water. Keep it in the boiling water for 10 minutes, and then cool it by placing it in a beaker
of water at room temperature. After it has cooled to room temperature centrifuge it. Note the
amount of precipitate obtained. Are the results in accord with what you would predict from a shift
33
in equilibrium above?* Compare the amounts of CaCO3 obtained before and after boiling the
mixture with that obtained when you used the Na2CO3 solution. Record your observations.
EXPERIMENT: Determine the pH and H+ concentration of 1 M NaHCO3 by testing 5 mL
portions with the following indicators: bromothymol blue, cresol red, and phenolphthalein.
Record your observations. Calculate [OH-]. What would happen to the [OH-] if the 1 M NaHCO3
solution is boiled for a few minutes and then cooled?
EXPERIMENT: Place 5 mL of 1 M NaHCO3 in a test tube, add 2 drops of phenolphthalein, and
place the test tube in a beaker of boiling water for 5 minutes. Record the color changes. Does
[OH-] increase or decrease?
RESULTS: In the result section of your lab report please list all the ions and molecules (except
for water) for the following solutions a) the carbonic acid solution (made by bubbling CO2
through water); b) the 1 M Na2CO3 and c) the 1 M NaHCO3 (without the CaCl2 and before
heating). Now for each of the 3 solutions list the ions in order of decreasing concentration. It will
be much simpler if you group them in pairs.
34
Name ___________________
Date_________________
CARBONIC ACID AND ITS SALTS
Data: Carbonic Acid Solutions
Bromocresol Green Color ____________ Methyl Orange Color ____________ pH = _______
Substance
H2CO3
H+
HCO3-
OH-
CO32-
Equilibrium
Concentration
Sodium Carbonate Solutions
Amount of CaCO3 produced (mm) = _____________________
Net ionic reaction =
Sodium Bicarbonate Solutions
Amount of CaCO3 initially _____________ Amount of CaCO3 after heating________________
pH of NaHCO3
Bromothymol Blue Color _______________ Cresol Red Color ______________
Phenolphthalein Color __________________ pH = __________ [OH-] = __________________
Phenolphthalein Color After Boiling ________________________ [OH-] increase decrease
Ion and Molecules Present: From highest to lowest concentration:
0.034 M H2CO3
1 M Na2CO3
1 M NaHCO3
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
_________
35
PROBLEMS
1) In some of the following cases the two substances cannot coexist in the same solution because
they will react together. For these cases write a net equation for the reaction that would occur. In
other cases write "No Reaction".
(a) OH- and HCO3-
(e) Ca2+ and HCO3-
(b) OH- and CO32-
(f) H2CO3 and CO32-
(c) OH- and H2CO3
(g) H+ and HCO3-
(d) H+ and H2CO3
(h) H+ and CO32-
2a) Using the equilibrium constants given in the lab calculate the value of the equilibrium
constant for 2 HCO3-  H2CO3 + CO32-.
b) Calculate the concentration of H2CO3 and CO32- in a 1 M solution of NaHCO3.
3a) Calculate the value of the equilibrium constant for H2CO3  2 H+ + CO32-
b) What is your experimental value for the pH of a solution of 1 M NaHCO3?
c) Show that for HCO3- that pH = (pK1 + Pk2)/2
36
EXPERIMENT SIX
Titration Curve for Potassium Acid Phthalate
INTRODUCTION:
In this experiment a pH meter equipped with a combination electrode (a glass electrode and a
calomel reference electrode in the same housing) and drop counter will be used to construct the
titration curve for KHP titrated with NaOH. From the titration curve and the concentration of the
base you can determine the concentration of the unknown KHP solution and both the first and the
second dissociation constants of phthalic acid.
When molar concentration units are used, the concentration quotient
(1)
Kc2 = [H+][P2-]
[HP-]
is only approximately constant over the whole accessible range of concentrations. When activities
are used, however, the ionization constant defined by
(2)
Ka2 = (aH+)(aP2-)
(aHP-)
is truly constant at all concentrations of the ions.
The activity may be regarded, for our purposes, as an effective concentration. In ionic solutions the
nearest neighbor shell around a cation will contain a (slight) predominance of anions, and vice
versa. This partial screening of the ions by their counterions makes the effective concentration a of
an ionic species somewhat smaller than its stoichiometric concentration. The activity and
concentration of an ion i are related by
(3) ai = Cifi
where fi is an empirical parameter, called the activity coefficient, which depends on the size of the
ion and the total ionic strength of the solution (cf. Chapter 8).
Electrochemical experiments, such as the measurement of pH, yield ion activities rather than ion
concentrations. At the ionic-.strengths to be used in this experiment the activity coefficients of the
phthalate anionic species differ appreciably from unity, and a serious misestimate of Ka results if
they are ignored,
The appropriate equations are
(4) Ka2 = (aH+)(aP2-) = (aH+)[P2-](fP2-)
(aHP-)
[HP-](fHP-)
37
Taking the log of both sides,
(5) log Ka2 = log (aH+) + log [P2-]/[HP-] + log (fP2-)/(fHP-)
Now since
(6) pH = - log(aH+) and pKa2 = -log Ka2, it follows that,
(7) pH = pKa2 + log [P2-]/[HP-] + log (fP2-)/(fHP-)
Thus when exactly half the KHP has been neutralized (when [P2-] = [HP-]) the measured pH and
pKa are roughly the same, except for the activity coefficient term.
At the equivalence point on the titration curve of a weak acid with strong base, the slope of the
curve is a maximum. We can use this fact to find the equivalence point very precisely. You will
construct a "differential" titration curve, a plot of (ΔpH/ΔV) vs. ΔV, which will be very sharply
peaked. The peak of this curve occurs at the equivalence point.
PROCEDURE:
D. Supply your instructor with a clean, dry, 125 mL Erlenmeyer flask labeled with your name and
locker number. You will receive 55 mL of your unknown KHP solution in this flask.
B. Set up a titration assembly, consisting of a pH meter, a magnetic stirrer, and a buret filled with
your standard 0.1 M NaOH..
C. Just before you are ready to prepare a titration curve standardize the pH meter according to the
directions given in a separate handout.
D. Using a pipet deliver 25 mL of the unknown KHP solution into a clean, dry 100 mL beaker
which already contains a clean magnetic stirring bar. Add 20 mL of distilled water and 2 drops
of phenolphthalein indicator. Put the beaker onto the magnetic stirrer so that the bar rotates
freely, but there is still room alongside it for the electrode to be inserted near the bottom of the
beaker. Mix the solution thoroughly with the stirrer.
E.
Wash the standardized electrode with your wash bottle, blot it dry, and position it throuh the
large hole of the drop counter. Read the pH meter and the buret, and record both readings.
F.
Turn on the magnetic stirrer so it is rotating slowly. You are now ready to begin collecting
data. Click the Collect button. No data will be collected until the first drop goes through the
drop counter slot. Adjust the drop rate to about 1 drop every 2 seconds. When the first drop
passes through check to see that the data has begun recording.
G. Continue watching your graph to see when the large increse in Ph takes place - this will be the
equivalence point. When this jump in pH occurs, add about 3 more milliliters of NaOH, click
STOP, and stop the titration. Remove the electrode, wash it thoroughly, and put it back into
38
its storage bottle. Remove the stir bar from the beaker and pour the solution down the sink.
Clean and dry the stir bar.
H. Find the equivalence point. The best method for determining the equivalence point is to take
the second derivative of of the pH vs. volume data, a plot of Δ2pH/Δvol2.
a. Open Page 3 by clicking o the Page window of the menu bar.
b. Analyze the second derivative plot and record the volume of NaOH at the equivalence
point.
I.
From this curve determine the pH at half and three-fourths neutralization of the KHP. Go to
the worksheet and use it to compute an average value for pKa2, taking into account the
activity coefficients. α’s (Na+ = 4, P2- = 6, HP- = 6, K+ = 3). The Debye-Huckel equation is
below.
J.
Having Ka2 in hand you can obtain Ka1 from the initial pH of the diluted unknown. In a
solution of pure KHP in water we know that, in concentration units,
(8) [K+] = [P2-] + [HP-] + [H2P] (conservation of mass) and
(9) [K+ ] + [H+] = 2[P2-] + [HP-] + [OH-] (electroneutrality)
Also, the equilibrium constants are related to the concentration quotients by
(10)
When activity coefficients are included, the solution of pure KHP has a [H+] given by;
(11)
In your diluted KHP solutions the Ka1 term in the denominator is only about 3% of the second
term and can be neglected, similarly the Ka1Kw term in the numerator is very small by comparison
with the first term. Moreover the activity coefficient of the molecule H2P is very close to one.
Under these special circumstances Eq. 11 becomes
(12)
and
39
Since we have earlier defined pH as -log(aH+), Eq. 12 reduces to the simple form
(13) 2pH = pKa1 + pKa2 + log fP2The ionic strength of your diluted starting solution is easily calculated from the concentration of
your unknown found in step 11. The activity coefficient of P2- at this ionic strength can be
interpolated from the table in your text. With this value, the measured pH, and pK2 which you
have calculated in step 12, pK1 is in hand.
REPORT: Turn in your titration curve, the concentration of the undiluted KHP unknown, and the
calculations of pKa1 and pKa2 from steps 12 and 13, using equations 7 and 13.
40
Initial
Conc. of Unk.
Vol. KHP used
Vol. NaOH used
Conc.NaOH used
Name
For Initial Solution
Calculation of PKa1
3/4 Neutralization
1/2 Neutralization
Calculation of PKa2
¾ Neut
½ Neut
41
μ
μ
pH
Vol.
NaOH
-
not needed
fHP
pH
Total
Vol.
2-
fP
-
fHP
Mol.
Na+
pKa1
2-
fP
Conc.
Na+
p
Mol.
K+
pKa2
Conc.
K+
Avg.pKa2
Mol.
HP-
EXPERIMENT SEVEN : Titration Curve of KHP
Conc.
HP-
Mol.
P2-
Date
Conc.
P2-
Ionic
Strength
EXPERIMENT SEVEN
Determination of the Equilibrium Constant For FeSCN2+
INTRODUCTION
When 2 reactants are mixed, the reaction typically does not go to completion. Rather, they will
react to form products until a state is reached whereby the concentrations of the reactants and
products remain constant. This is a dynamic state in which the rate of formation of the products is
equal to the rate of formation of the reactants. The reactants and products are in chemical
equilibrium and will remain so until affected by some external force. The equilibrium constant Kc
for the reaction relates the concentration of the reactants and products.
In this experiment we will study the equilibrium properties of the reaction between iron (III) ion
and thiocyanate ion:
Fe3+ (aq) + SCN– (aq) ---> FeSCN2+ (aq)
When solutions containing Fe3+ ion and thiocyanate ion are mixed, the deep red thiocyanatoiron
(III) ion (FeSCN2+) is formed. As a result of the reaction, the starting concentrations of Fe3+ and
SCN- will decrease: so for every mole of FeSCN2+ that is formed, one mole of Fe3+ and one mole of
SCN- will react. The equlibrium constant expression Kc, according to the Law of Chemical
Equilibrium, for this reaction is formulated as follows:
[FeSCN2+] / [Fe3+][ SCN- ] = Kc
Square brackets ([]) are used to indicate concentration in mols/liter, i.e., molarity (M).
The value of Kc is constant at a given temperature. This means that mixtures containing Fe3+ and
SCN– will react until the above equation is satisfied, so that the same value of the Kc will be
obtained no matter what initial amounts of Fe3+ and SCN– were used. Our purpose in this
experiment will be to find Kc for this reaction for several mixtures made up in different ways, and
to show that Kc indeed has the same value in each of the mixtures.
The reaction is a particularly good one to study because Kc is of a convenient magnitude and the
red color of the FeSCN2+ ion makes for an easy analysis of the equilibrium mixture using a
spectrophotometer. The amount of light absorbed by the red complex is measured at 447 nm, the
wavelength at which the complex most strongly absorbs. The absorbance, A, of the complex is
proportional to its concentration, M, and can be measured directly on the spectrophotometer using
the Beer-Lambert law:
A = M
Where  = extinction coefficient, M = Molarity, and  = path length (1cm)
42
EQUIPMENT
5 test tubes
10 ml graduated cylinder
10 ml graduated measuring pipette (0.1 ml graduations)
stirring rod
small labels or markers
5 cuvettes for the spectrophotometers
250 ml bottle acetone for rinsing
HAZARD: As always wear Safety glasses while performing this experiment
CONTAIMINATION NOTES: If your flask is wet before you prepare your standard/sample
solutions ensure that the flask is wet with diluant (in this case it is the 0.0200M Fe(NO3)3 in 1.0 M
HNO3 ).
EXPERIMENT: Beer-Lambert Data
A solution of 0.0200M Fe(NO3)3 in 0.5 M HNO3 has been prepared for you.
Dilute 5.0, 10.0 and 15.0 ml portions of 2.00 x10-4 M KSCN to 100 ml with the 0.0200M
Fe(NO3)3 in 0.5 M HNO3. This will give you 3 solutions that can be assumed to be 1.0x10-5,
2.0x10-5 and 3.0x10-5 Min FeSCN2+.
Measure the absorbances of these solutions at 447nm, using a solution of 0.0200M Fe(NO3)3 in 0.5
M HNO3 as the reference solution. Measure the absorbances of these solutions at this wavelength.
Plot absorbances vs [FeSCN2+] using Graphical Analysis. Find the slope and intercept for this plot.
EXPERIMENT: Determination of Kc
The mixtures will be prepared by mixing solutions containing known concentrations of iron (III)
nitrate, Fe(NO3)3, and potassium thiocyanate, KSCN. The color of the FeSCN2+ ion formed will
allow us to determine its equilibrium concentration. Knowing the initial composition of a mixture
and the equilibrium concentration of FeSCN2+, we can calculate the equlibrium concentrations of
the rest of the pertinent species and then determine Kc.
Label five regular test tubes 1 to 5, with labels or by noting their positions in your beakers. Pour
about 30 mL 0.02 M Fe(NO3)3 in 0.5 M HNO3 into a dry 100 mL beaker. Pipet 5.00 mL of that
solution into each test tube. Then add about 20 mL 2.00 x 10-4 M KSCN to another dry 100-mL
beaker. Pipet 1,2,3,4, and 5 mL from the KSCN beaker into each of the corresponding test tubes
labeled 1 to 5, then pipet the proper number of milliliters of 0.5 M HNO3 into each test tube to
bring the total volume in each tube to 10.00 mL.
43
The volumes of reagents to be added to each tube are summarized in the table below.
Reagents (in mLs)
Fe(NO3)3
KSCN
HNO3
1
5.00
1.00
4.00
TestTube #
2
3
4
5.00
5.00
5.00
2.00
3.00
4.00
3.00
2.00
1.00
5
5.00
5.00
0.00
Mix each solution thoroughly with a glass stirring rod. Be sure to dry the stirring rod after mixing
each solution to prevent cross-contamination.
Measure the absorbance of each mixture at 447 nm as demonstrated by your instructor and put the
data in the worksheet provided.
Determine the concentration of FeSCN2+ from your calibration curve. Record the value on your
Report form. Repeat the measurement using the mixtures in each of the other test tubes.
44
Name _______________________
Date____________________
Determination of the Equilibrium Constant For FeSCN2+
Beer-Lambert Data
Data
Data For Beer-Lambert Plot
10 mL
5 mL
15 mL
Absorbance
Data for Equilibrium Calculation
Data
1
Test Tube #
3
2
Absorbance
Conc of FeSCN2+
Conc of Fe3+
Conc of SCN-
Value of Keq
Average Keq =
Staple Beer-Lambert plot to the back of this sheet.
45
4
5
EXPERIMENT NINE
THE DETERMINATION OF A HEAT OF REACTION
The purpose of the experiment is to determine the quantity of heat liberated in the reaction,
Mg(s) + 2 HCl (aq)  H2 (g) + MgCl2 (aq)
The heat liberated in this reaction can be trapped and utilized to melt ice. Ice is less dense than water
and as a consequence there is a volume change when ice melts. This volume change can be measured
and used to calculate how much energy was transferred to the ice. This in turn is used to measure the
amount of energy released in the reaction (the enthalpy).
EXPERIMENT: Students may work in pairs with one student calling out the time and the other
student reading the volume and recording both the time and volume. The calculations are to be made
individually. Fill a styrofoam container with crushed ice from the ice provided in the lab and place
in it a small flask containing 3 mL of 6 M HCl and 6 mL of water. While this solution is cooling,
assemble the bottle and stopper as shown in the lab discussion. Fill the bottle with tap water and push
the stopper in tightly. Water will spurt out the top of the pipet. Watch the water level in the pipet over
a 5 minute period. If it does not drop, the apparatus is free of leaks and you may continue with the
experiment. If the level drops, try to repair the leak with the aid of your instructor or instructional
assistant. ( Do not use grease on the stopper.) When your assembly is leak-proof, fill the bottle with
crushed ice to the brim. Add ice water from the styrofoam container to fill the bottle completely and
insert the stopper. Pour the 9 mL of HCl solution that you cooled to 0 C in the reaction test tube.
Immediately place the bottle in the styrofoam container and surround the bottle with crushed ice and
water to just below the rim of the test tube. Insert the cork in the test tube loosely. Allow the
apparatus to stand for 15 minutes during which time the temperature should become constant at 0 C.
During this time interval weigh a 0.1 g sample of magnesium ribbon to the nearest milligram. Record
the mass in your notebook.
After the apparatus has been allowed to stand for 15 minutes, adjust the water level in the pipet to a
reading of 0.8 mL or greater by placing a small piece of latex tubing on the end of the pipet and filling
it with water. By pushing on the stopper it is possible to raise the level of the water inside of the pipet
until it meets the water in the tubing. Releasing the stopper will pull the water from the tubing into
the pipet thus filling it. This technique may require practice. Begin reading the pipet volume once the
water level has reached some convenient mark on the pipet. Read the pipet volume each minute for
5 minutes and record these readings in your notebook. If the volume change in the first 5 minutes is
greater than 0.05 mL then you have a leak and must readjust the stopper and start again at the
beginning of this paragraph. If the volume change during this 5 minute period is 0.05 mL or less, roll
the magnesium ribbon into a loose coil, and drop it into the HCl solution. Place the cork LOOSELY
over the test tube and continue to record the volume each minute for a period of 15-20 minutes.
46
Plot your data using a computer graphing program, or ask for a sheet of graph paper. The results
should be similar to those shown in the figure below.
CALCULATIONS: Draw parallel lines through the initial and final points. The volume change due
to the heat of reaction of your sample of magnesium is the vertical distance between the parallel lines.
It takes 875 calories to melt one mL of ice. To calculate the energy required to melt your volume of
ice take your measured volume and multiply by 0.875 Kcals/mL. The result of this calculation gives
you the amount of energy released when about 0.1 gram of magnesium reacts. To calculate the amount
of energy released when one mole of magnesium reacts you must calculate the number of moles of
magnesium used in your experiment. Calculate the number of moles of magnesium used in your
experiment (the atomic mass of magnesium is 24.3 g/mole). By dividing the number of calories
calculated earlier and dividing this by the number of moles of magnesium used one obtains the
enthalpy for the following reaction,
Mg (s) + 2 HCl (aq)  MgCl2 (aq) + H2 (g) + Enthalpy
Check your calculations and report the result of this calculation in your notebook.
PROBLEM
The standard heat of formation of HCL (aq) is -40 Kcal/mole. From this, using Hesses law, calculate
the standard heat of formation of MgCl2 (aq). (Hint: What is meant by the standard heat of formation?)
47
EXPERIMENT ELEVEN
ELECTROCHEMICAL CELLS
In this experiment you will investigate reactions in which electrons are transferred from one reactant
to another. Since electrons are being transferred we can force them to travel through an electrical
connection and thereby be measured. We call these set-ups electrochemical cells. We can observe
the effect of electron transfer in other ways. In the first experiment we will observe the formation of
elements by electron transfer.
EXPERIMENT: For each of the following experiments, record your observations and write an net
equation for each reaction. Also determine which of the two substances is the better oxidizer
(oxidizers get reduced and are usually called oxidizing agents).
Place a small piece of;
a) Metallic lead in 5 mL of 1 M Cu(NO3)2.
b) Metallic zinc in 5 mL of 1 M Pb(NO3)2.
c) Metallic copper in 5 mL of 0.10 M AgNO3.
d) Metallic zinc in 5 mL of 1 M HCl.
e) Metallic lead in 5 mL of 1 M HCl.
f) Metallic copper in HOT 6M HNO3. (Silver also reacts with HNO3).
g) 1 mL of CCl4 and 3 mL of 1 M KI and 3 mL of Bromine water.
h) Metallic copper in 5 mL of Bromine water. Heat in the hood to remove any excess bromine
then add 2 mL of 6 M NH3. Using NH3 allows you to see the blue color of the copper ions in
solution.
i) 1 mL of CCl4 and 3 mL of 1 M KI and 3 mL of 6 M HNO3.
Some of these reactions may be slow, allow them to stand for 15 minutes and look carefully at the
surface of metal of some of the slower reactions.
Write all of the half reactions as reductions. Now arrange these reactions in such a way that a reaction
will occur when you take a reaction and reverse any of the reactions above it (reversing a reaction
makes it an oxidation). It is sometimes difficult to see the lead reaction with H+, a reaction does occur
so place the lead appropriately. What additional experiments would you need to perform to place
bromine and iodine in their proper position?
48
ELECTROCHEMICAL CELLS
In this portion of the experiment you will be measuring the voltage produced by three electrochemical
cells. The electrodes used in electrochemical cells are usually metals and occasionally carbon rods.
The oxidation reaction occurs at the anode and reduction at the cathode so that in electrochemical cells
electrons flow from the anode to the cathode and the anode will be negatively charged compared to
the cathode. You may work in pairs.
EXPERIMENT: Put about 40 mL of 1 M CuSO4 into a 250 mL beaker and about 40 mL of 1 M
ZnSO4 into a porous cup. Place the porous cup into the 250 mL beaker containing the CuSO4. Put a
copper strip into the copper sulfate solution and a zinc strip into the zinc solution. Attach either a
voltmeter to the zinc and copper electrodes and measure the voltage. Draw the cell and determine
which electrode is acting as the anode and cathode. Write the reaction occurring at each electrode.
EXPERIMENT: Repeat the experiment above using Bromine water and a carbon electrode to replace
the copper sulfate solution and the copper electrode. Measure the voltage. Draw the cell and determine
which electrode is acting as the anode and cathode and write the reaction occurring at each electrode.
MEASUREMENT OF A SOLUBILITY PRODUCT
Electrochemical cells can be used to determine the equilibrium constant of chemical reactions. In this
lab you will calculate the Ksp of Cu(OH)2 by the following set of half-reactions,
Cu(s) + 2 OH- <-----> Cu(OH)2(s) + 2eCu2+(1 M) + 2e- <----> Cu(s)
Cu2+(1 M) + 2 OH- <-----> Cu(OH)2 Eo = ?
From the measured value of Eo one could evaluate the Ksp for Cu(OH)2(s) using,
E = (0.0592/n) log (1/Ksp)
where n is the number of electrons transferred in the reaction.
EXPERIMENT: Clean the electrochemical cell used above. Place 40 mL of 1 M CuSO4 in the beaker
and 40 mL of 1 M NaOH in the cup. Add 4 drops of 1 M CuSO4 to the 40 mL of NaOH. Place a
copper strip into each cells and measure the voltage as previously. Note the negative terminal. Which
electrode is the anode? Calculate the Ksp for Cu(OH)2(s).
49
PROBLEMS
1) Use your table of oxidizing and reducing agents to predict whether the following reactions will
occur as written, or will occur in the opposite direction.
a) I2 + Zn(s) = Zn2+ + 2 Ib) 2 NO3- + 8 H+ + 6 Cl- = 2 NO(g) + 4 H2O + 3 Cl2
c) Pb(s) + Cu2+ = Pb2+ + Cu(s)
d) Pb2+ + 2 Br- = Pb(s) + Br2
2) You are given the half-reaction for pure water,
2 H+ (10-7M) + 2e- = H2 (g, 1atm) E = -0.41 eV
Now, rank the following metals according their standard electrode potentials as found in the CRC
Handbook, (Na, K, Mg, Al, Zn, Fe, Sn, Pb, Cu, and Ag) by placing the most positive one at the top
of the list and then determine which of the following metals should,
a) React with water evolving gas.
b) React with 1 M HCl but not with water to produce hydrogen gas.
c) What reagent and conditions would you use to dissolve the metals that cannot be dissolved in 1 M
HCl?
3) Calculate the voltage of an electrochemical cell which uses the reaction,
Cu2+(1 M) = Cu2+(0.001 M)
4) Would your results have changed if you had added more drops of CuSO4 to your Ksp cell? Why
or why not?
50
EXPERIMENT TWELVE
The Mystery Experiment
51
52
pH Color Chart
0
1
Malachite Green Y
Methyl Violet
3
5
4
B
9
10
11
12
13
V
14
C
R
Y
Y
Methyl Red
B
Y
R
Bromothymol Blue
B
Y
Cresol Red
R
Thymol Blue
R
Alizarin Yellow R
8
V
Bromocresol Green
Phenolphthalein
7
6
Y
Methyl Orange
Indigo Carmine
2
Y
V
Y
B
C
R
O
Y
B
B = Blue Y = Yellow V = Violet C = Clear O = Orange R = Red
Y
Exam I
Rate Laws
Activation Energies
Mechanisms
Radioactive Decay
Kinetics and Activation Energy
1) Rate information was obtained for the following reaction at 25C and 33C;
Cr(H2O)63+ + SCN- ---> Cr(H2O)5NCS2+ + H2O(l)
[Cr(H2O)63+]
[SCN-]
2.0x10-11
2.0x10-10
9.0x10-10
2.4x10-9
1.0x10-4
1.0x10-3
2.0x10-3
3.0x10-3
0.10
0.10
0.15
0.20
2.0x10-11
1.0x10-4
0.20
Initial Rate
@ 33C
a) Write a rate law consistent with the experimental data.
b) What is the value of the rate constant at 25C?
c) What is the value of the rate constant at 33C?
d) What is the activation energy for this reaction?
e) What is the reverse rate law?
2) The mechanism for the decomposition of phosgene
COCl2(g) ---> CO(g) + Cl2(g)
is thought to be,
fast eq.
slow
fast eq.
Cl2(g) <---> 2 Cl(g)
COCl2(g) + Cl(g) ---> COCl(g) + Cl2(g)
COCl(g) <---> CO(g) + Cl(g)
Based on this mechanism, what is the rate law for this reaction?
3) The following experimental data were obtained for the reaction at 250 K,
F2 + 2 ClO2 -----> 2 FClO2
[F2]/M
0.10
0.10
0.20
[ClO2]/M
0.010
0.040
0.010
Rate/(M/sec)
1.2x10-3
4.8x10-3
4.8x10-3
3a) Write a rate law consistent with this data.
Rate =
3b) What is the value and units of the rate constant?
4) The bromination of acetone is acid catalyzed;
CH3COCH3 + Br2 ---> CH3COCH2Br + H+ + BrThe rate of disappearance of bromine was measure for several different concentrations of
acetone, bromine and hydrogen ions;
Rate
6x10-5
6x10-5
1.2x10-4
3.2x10-4
8x10-5
[Acetone]
[Br2]
0.30 0.050
0.30 0.100
0.30 0.050
0.40 0.050
0.40 0.050
[H+]
0.050
0.050
0.100
0.200
0.050
a) What is the forward rate law for this reaction?
b) What is the value and units of the forward rate constant?
c) What is the reverse rate law?
d) If the equilibrium constant is 1.3x103, what is the value of the reverse rate constant?
5) At low temperatures, the rate law for the reaction,
CO(g) + NO2(g) ---> CO2(g) + NO(g)
can be determined by the following data;
[NO]/10-3 M [CO]/10-3 M [NO2]/10-3 M Initial Rate/10-4
1.20
1.20
1.20
2.40
1.50
3.00
0.75
1.50
0.80
0.80
0.40
0.40
3.60
7.20
0.90
0.90
Write a rate law in agreement with the data.
Which of the following mechanisms is consistent with this rate
law? (Circle A, B, or C)
A) CO + NO2 ---> CO2 + NO slow
B) NO2 <---> NO + O equil.
O + CO --> CO2 slow
C) 2 NO2 <---> 2 NO + O2 equil.
CO + O2 ---> CO2 + O slow
O + NO ---> NO2
fast
D) None of the above
6) Given the following rate data, calculate the rate law at 25C.
Rate/10-4
[I-]/M
[OCl-]/M
0.20
0.40
0.20
0.050
0.050
0.100
6.10
12.2
24.4
What is the rate law?
Rate =
What is the value of the rate constant?
What are the units for the rate law?
7) The following experimental data were obtained for the reaction at 250 K,
F2 + 2 ClO2 -----> 2 FClO2
[F2]/M
0.10
0.10
0.20
[ClO2]/M
0.010
0.040
0.010
Rate/(M/sec)
1.2x10-3
4.8x10-3
4.8x10-3
7a) Write a rate law consistent with this data.
Rate =
7b) What is the value and units of the rate constant?
8) The following experimental data were obtained for the reaction;
2 NO + 2 H2 ---> N2 + 2 H2O
Initial Rate/10-5
0.60
2.40
0.30
[NO]/10-2 M
[H2]/10-2 M
0.50
1.00
0.25
0.20
0.20
0.40
a) Write a rate law in agreement with the data.
Rate =
b) What is the value of the rate constant?
c) What is the new rate when [NO] = 3x10-2 M and [H2] = 1.2x10-2 M?
d) Which of the following mechanisms are consistent with the rate law above?
i) 2 NO + H2 ---> N2O + H2O (slow)
ii) 2 NO <----> N2O2 (equil)
N2O2 + H2 ----> N2O + H2O (slow)
iii) NO + H2 <---> H2O + N (equil)
N + NO ----> N2O (slow)
e) Platinum acts as a catalyst for this reaction. What term must be added to the rate law to
account for the presence of the catalyst?
f) Would platinum be a homogeneous or heterogeneous catalyst for this reaction?
If the temperature is increased by 6.8C the reaction rate increases 1.65 times. What is the
activation energy of the reaction?
9)The following experimental data were obtained for the reaction at 25C,
2 NO + H2 ---> N2O + H2O
Initial [NO]
(x103 M)
6.40
12.8
6.40
Initial [H2]
(x103 M)
Initial Rate
(M/s x 105)
2.20
1.10
4.40
2.60
5.20
5.20
Which of the following mechanisms are consistent with the rate law for this reaction?
a)
2 NO + H2 ---> N2O + H2O (slow)
b)
2 NO <----> N2O2 (equil)
N2O2 + H2 ----> N2O + H2O (slow)
c)
NO + H2 <---> H2O + N (equil)
N + NO ----> N2O (slow)
10) At low temperatures, the rate law for the reaction,
CO(g) + NO2(g) ---> CO2(g) + NO(g)
is;
Rate = k [NO2]2
Which of the following mechanisms is consistent with this rate
law? (Circle A, B, or C)
a) CO + NO2 ---> CO2 + NO
slow
b) 2 NO2 <---> N2O4
fast equil.
N2O4 + 2 CO --> 2 CO2 + 2 NO slow
c) 2 NO2 ----> N2O4
N2O4 <----> 2 NO + O2
O2 + CO2 ---> 2 CO2
slow
fast equil.
fast
11) Given the following rate data, calculate the rate law at 25C.
Rate/10-4
[I-]/M
[OCl-]/M
6.10
12.2
36.6
0.20
0.40
0.60
0.050
0.050
0.100
14.4
0.20
0.050
@ 25C
@ 33C
What is the rate law?
Rate =
What is the value of the rate constant at 25C and at 33C?
What is the activation energy for this reaction?
What is the half-life of this reaction at 45C if [I-] = [OCl-] = 0.25 M?
12) Thiosulfate (S2O32-) can react with triiodide (I3-) according to the following reaction at 25C,
2 S2O32- + I3- <----> S4O62- + 3 I- Keq = 3.75x105
Experimentally the forward rate law can be determined from the following data,
Rxn
#1
#2
#3
Rate/Msec-1
2.56x10-4
1.28x10-4
1.92x10-4
[S2O32-]/M
[I3-]/M
0.040
0.020
0.060
0.12
0.12
0.06
a) What is the forward rate law?
b) What is the value of the rate constant? (Include units)
c) What is the reverse rate law?
d) What is the value of the reverse rate constant?
e) What are the minimum number of steps required in the mechanism of the forward rate law?
Circle one
1
2
3
h) If reaction #1 is heated to 35C the rate increases to 4.8x10-4 M/sec. What is the activation
energy of this reaction?
i) At what temperature would the reaction rate double for reaction #1?
13) A cook finds that it takes 30 minutes to boil potatoes at 100C in an open sauce pan and only
12 minutes to boil them in a pressure cooker at 110C. Estimate the activation energy for
cooking potatoes, which involves the conversion of cellulose into starch. Remember that there
is an inverse relationship between time and the rate constant.
14) When N2O4 decomposes it forms NO2,
N2O4 ----> 2 NO2
If the half-life of this reaction is 1386 seconds, how much of a 10 gram sample would be left
after 1500 seconds? (Is the reaction first or second order?)
15) An archaeologist measured the amount of radioactivity of a piece of cloth used to wrap an
Egyptian mummy. The cloth was found to have a decay rate of 9.1 dpm. If the decay rate is 15.3
dpm in living tissue, how old is the mummy? t1/2 = 5730 years.
16) Uranium is radioactive and decays into lead. This process can be used to date rocks. A
piece of zirconium was dated using this process and it was found that this rock contained
3.2x10-3 grams of uranium and 2.2x10-5 grams of lead. If the half-life of uranium is 4.41x109
years how old is the rock?
17) Assuming that the loss of ability to recall learned material is a first-order process with a halflife of 35 days. Compute the number of days required to forget 90% of the material that you
learned in preparation for this exam.
18) Under acidic conditions sucrose (table sugar) can be broken down into its individual sugars,
glucose and fructose. At 27C it takes 54.5 minutes to convert half the sucrose to glucose and
fructose and at 37C it takes 13.7 minutes. Estimate the activation energy for the breakdown of
sucrose.
18b) The above reaction is known to second order which means that the half-life is dependent on
the initial concentration. Will this fact effect your calculation of the activation energy above?
Why or why not. (Hint: What must you assume about the initial concentration of [A] when using
the Arrhenius equation?)
19) A 10 gram sample of 131I was sent from a pharmaceutical company to a hospital for use in the
treatment of hyperthyroidism. If the half life of 131I is 8.07 days, how much of the sample would
be left after a 2 day mail delivery?
20) The denaturation of the virus that causes the rabbit disease Myxomatosis can be followed by
heating the virus under a microscope. It is observed that the reaction is first-order and that it
takes 22.35 minutes at 50C and 0.35 minutes at 60C for the virus to denaturate. Estimate the
activation energy for the denaturation of the Myxomatosis virus.
Exam II
Chemical Equlibria
LeChatlier’s Principle
Acid Dissociation
Base Hydrolysis
Titration Curves
Solubility Product
pH Controlled Solubility
Amphoterism
Equilibria and LeChatliers Principle
1) For each of the following sets of compounds write the equilibrium reaction that would occur
when the compounds are mixed together.
a) HNO3 and NaBenz
b) NH4Cl and KOH
c) NaCN and NaOH
2a) If H2 and Cl2 are added to a container, both at 2 atm, what will the pressure of HCl be after
the system reaches equilibrium?
H2 + Cl2 ---> 2 HCl
K = 150
2b) If the equilibrium pressure of HCl is 2 atm what pressures of H2 and Cl2 must have been
added to the container originally?
3) If 100 grams of NaF and 70 grams of KOH are added to 250 mL water what is the equilibrium
concentration of all ionic species. Ka = 6.76x10-4
4) What is the pH of 100 mL of 1.2 M HAc after 20 mL of 2 M NaOH is added? Ka = 1.8x10-5
5) What is the pH of a solution that is made from 0.10 M HBenz, 0.25 M NaBenz, and 0.25 M
KOH? Ka = 6.46x10-5
6) What is the pH of a solution where 100 mL of 0.50 M NH4OH is completely neutralized by 65
mL of HCl? Kb = 1.8x10-5
7) Ca(IO3)2 is a slightly soluble precipitate. What would the concentration of IO3- be in a solution
saturated with Ca(IO3)2? Ksp = 1.5 x 10-15
8) What is the pH of a 0.10 M solution of NaAc? Ka = 1.8x10-5
Ac- + H2O <---> HAc + OH-
9) When ammonia is heated it decomposes to N2 and H2 according to the following reaction,
2 NH3 <----> N2 + 3 H2
Given 3 atm of NH3 and an equilibrium constant of 3x10-3, what will the final pressure (total
pressure) be in the system?
10) What is the pH of a solution made by adding 0.2 mole of NaAc to 250 mL of 1 M acetic
acid? Ka = 1.8x10-5
11) In class we noted that a 0.10 M solution of calcium ions can be precipitated with 0.10 M
NaOH but not in 0.10 M NH4OH. The solubility product for Ca(OH)2 is 7.88x10-6. Please
explain why 0.10 M Ca2+ will not precipitate in 0.10 M NH4OH (Kb = 1.8x10-5).
12) What is the pH of the resulting solution when 0.25 mole of NaCH3CO2 and 0.15 mole HCl
are added to 200 mL water? The Ka for CH3COOH is 1.8x10-5M.
13) What is the pH of a solution made by adding 0.30 mole NH3(aq) to 0.50 mole NH4Cl and
0.25 mole KOH? Kb for NH3(aq) = 1.8x10-5M
14) What is the concentration of all ionic and molecular species when you add 30 mL of 0.5 M
NaOH to 120 mL of 0.75M HCN? Ka for HCN is 4.8x10-10.
15a) What is the pH of the following solutions when the following amounts of 1.20 M NaOH
are added to 160 mL of 3 M dl Aspartic acid?
H2Asp <---> H+ + HAspHAsp- <---> H+ + Asp2-
1.38x10-4
1.51x10-10
0 mL
200 mL
400 mL
700 mL
800 mL
900 mL
15b) An Aspartic acid buffer of pH 5 was made by adding some NaHAsp to 0.5 moles of H2Asp.
How much 1.2 M NaOH or 1.2M HCl must you add to this buffer in order to make a new buffer
of pH 4.5?
16) What is the pH of the solution when 75 mL of 0.8 M dl-Histidine is titrated with the
following volumes of 1.20 M NaOH?
H3His <---> H+ + H2HisH2His- <---> H+ + HHis2HHis2- <---> H+ + His3-
pKa1 = 2.40
pKa2 = 6.04
pKa3 = 9.33
0 mL
30 mL
50 mL
75 mL
100 mL
125 mL
135 mL
150 mL
175 mL
17) How much NH4Cl and NH4OH must you add to 250 mL of water to make a buffer of pH =
8.0? Kb = 1.8x10-5
18) How many moles of sodium acetate must be added to 100 mL of 0.25 M acetic acid to make
a buffer of pH = 4.0? pKa = 4.74 for acetic acid.
19) Another way of making the buffer from 2 above would be to titrate the acid with a base.
What volume of 0.40 M NaOH must be added to 100 mL of 0.25 M acetic acid to make a buffer
of pH = 4.0? pKa = 4.74 for acetic acid.
20) How much 0.5 M NaOH must be added to 0.75 M H3PO4 to make a buffer of pH = 8? Ka1 =
2.12, Ka2 = 7.21, Ka3 = 12.67
Solubility Products
21) Are the following molecules acidic, basic or neutral in aqueous solution?
NaF
Cr(NO3)3
KCl
NH4CN
acid
acid
acid
acid
basic
basic
basic
basic
neutral
neutral
neutral
neutral
can't tell
can't tell
can't tell
can't tell
22) Are the following compounds soluble or insoluble in water?
NaIO3
Cr(OH)3
PbSO4
CaCl2
Soluble
Soluble
Soluble
Soluble
Insoluble
Insoluble
Insoluble
Insoluble
FeSO4
ScCl3
Na2O
PbSO4
Soluble
Soluble
Soluble
Soluble
Insoluble
Insoluble
Insoluble
Insoluble
23) What is the solubility of PbBr2 in pure water? Ksp = 4x10-5
24) What is the solubility of Ca(OH)2 in 0.05 M NaOH? Ksp = 5.5x10-6
25) What is the solubility of Ca3(PO4)2 in water? Ksp = 5.87x10-8
26) What is the solubility of Ca3(PO4)2 in 0.5 M Na3PO4? Ksp = 5.87x10-8
27) What is the solubility of AgCl in a solution of 1 M HCl? Ksp = 1.8x10-10
28) At what pH will the concentration of Cu2+ exceed 0.02 M given the following equilibrium?
Cu(OH)2 ----> Cu2+ + 2 OH- Ksp = 2.2x10-20
29) How much NH3(aq) must you add to 100 grams of AgCl in order to dissolve all of the AgCl.
Assume a liter of solution and calculate the concentration of NH3(aq).
AgCl ---> Ag+ + Cl- Ksp = 1.8x10-10
Ag(NH3)2+ ----> Ag+ + 2 NH3(aq) Keq = 1.6x10-9
30) Copper(I) ions in aqueous solution react with NH3 according to,
Cu+ + 2 NH3 <----> Cu(NH3)2+ Keq = 6.3x1010
Calculate the solubility of CuBr (Ksp = 5.3x10-9) in a solution in which the equilibrium
concentration of NH3 is 0.185 M.
31) What is the solubility of PbCl2 in a solution of 0.10 M H2S at a pH=0? Ksp =1.6x10-5 for
PbCl2, Ksp = 8x10-28 for PbS, Keq = 1x10-20 for H2S (to 2 H+ + S2-).
Rxn = Pb2+ + H2S ---> PbS(s) + 2 H+
32) Calculate the maximum concentration of Fe3+ in an 0.10 M H2S solution buffered at pH=6.
Ksp = 1.6x10-21 for Fe2S3 and Keq = 1x10-20 for H2S (to 2 H+ + S2-).
33) What is the solubility cadmium sulfide (CdS) in a saturated solution of H2S = 0.10 M at pH
= 3?
CdS <--> Cd2+ + S2- Ksp = 8x10-27
H2S <--> 2H+ + S2Keq = 1.1x10-20
34) What is the solubility (silver ion concentration) of silver sulfide (Ag2S) in a saturated
solution of H2S = 0.10 M at pH = 3?
Ag2S <--> 2 Ag+ + S2- Ksp = 1.6x10-49
H2S <--> 2H+ + S2Keq = 1.1x10-20
35) How much of each precipitate will form and what will the concentration of lead be if
0.75 mole sodium oxalate (Na2C2O4 = Na2Ox) is added to 100 mL of solution containing 0.30 M
Mg(NO3)2 and 0.5 M Pb(NO3)2?
MgOx(s) <---> Mg2+ + Ox2- Ksp = 8.6x10-5
PbOx(s) <---> Pb2+ + Ox2- Ksp = 4.8x10-10
36) What is the solubility of silver carbonate (Ag2CO3) in a solution saturated with carbonic acid
(0.034M) and which is buffered at pH = 9?
Ag2CO3 <---> 2 Ag+ + CO32H2CO3 <--> 2 H+ + CO32-
Ksp = 8.1x10-12
Ka = 2.02x10-17
Chemistry 121
Second Exam
Name
April 10, 2008
CLOSED BOOK EXAM - No books or notes allowed. ALL work must be shown for full credit.
You may use a calculator.
Question
Credit
1(10)
2(15)
3(45)
4(20)
5(10)
Total
1) Which of the following compounds is the most soluble in water? Show your work.
BaF2
Ksp = 1.84x10-9
Mg3(PO4)2
Ksp = 1.04×10-15
Ag3PO4
Ksp = 1.2x10-15
2) Please write the equilibrium reaction that will occur when the following compounds are
mixed. Write all reactions as dissociations where appropriate.
NaOH + HAc
Na3PO4 + HCl
HCl + KOH
3) Calculate the pH of 80 mL of 1.5 M Carbonic acid upon addition of the following amounts of
0.6 M NaOH
H2CO3 <---> H+ + HCO3HCO3- <---> H+ + CO32-
Ka1= 4.3x10-7
Ka2= 4.7x10-11
a) 0 mL
b) 100 mL
c) 150 mL
d) 200 mL
e) 300 mL
f) 365 mL
g) 400 mL
h) 450 mL
i) How much 0.8 M NaOH must be added to 80 mL of 1.5 M H2CO3 to make a buffer of pH =
10?
4) What is the solubility of Ag2S in a saturated solution of H2S = 0.01 M at pH = 3?
Ag2S <---> 2 Ag+ + S2- Ksp = 6.0x10-50
H2S <---> 2 H+ + S2- Keq = 6.84x10-23
5a) Iron can be separated from copper using pH controlled solubility in H2S. CuS is not soluble
in acid solutions and FeS is soluble. Draw and approximate ion separation curve for CuS and
FeS.
5b)
Define amphoterism and give an example.
Exam III
Thermodynamics
1 , 2nd, and 3rd Laws
Enthalpy, Entropy, and Gibbs Energy
Heats of Formation
Hess’ Law
Spontaniety
Gibbs Equation
Electrochemical Cells
Balancing Redox Reactions
Nernst Equation
Batteries
Electroplating and Hydrolysis
st
Thermodynamics
1) A 100 gram block of copper was heated with a bunsen burner and then thrown into a cup of
water (250 mL) initially at 25C. When the system reached thermal equilibrium the final
temperature of the water and the block was 61C. What was the initial temperature of the block
of copper? Cp(Cu)=24.5 J/mol-K, Cp(H2O)=75.2 J/mol-K, FWT(Cu) = 63.55 g/mol, FWT(H2O)
= 18 g/mol, density of water = 1 g/ml
2) Given the following information calculate the heat of formation of C2H4.
C2H4 + 3 O2 ------> 2 CO2 + 2 H2O ΔH = -414 kJ/mol
C + O2 -----> CO2 ΔH = -393.5 kJ/mol
H2 + ½ O2 -----> H2O ΔH = -241.8 kJ/mol
3) When ethanol (C2H5OH) burns in oxygen it forms CO2 and water. If the heat of combustion
is -1237.7 kJ/mol, the heat of formation of water is -241.8 kJ/mol, and the heat of formation of
carbon dioxide is -393.5 kJ/mol, what is the heat of formation of C2H5OH?
4) Given a block of ice at 0C how much of the ice will melt if you put a 100 gram block of
aluminum on it that is initially at 100C? Cp(Al)=24.1 J/mol-K, Cp(H2O)=75.2 J/mol-K, Heat
of Fusion = 6.01 kJ/mol for ice, FWT(Al) = 26.98 g/mol, FWT (water) = 18 g/mol
5) 10 grams of diethyl ether was placed into a bomb calorimeter and then combusted in excess
oxygen. The balanced combustion reaction was as follows;
C4H10O + 6 O2 -----> 4 CO2 + 5 H2O
The heat released by this reaction heated 2500 mL of water 8.6C. What is the heat of formation
of the diethyl ether? Cp(H2O) = 75.2 J/mol-K, ΔHf(CO2) = -393.5 kJ/mol, ΔHf(H2O) = -241.8
kJ/mol, FWT (Diethyl ether) = 74 g/mol, FWT (water) = 18 g/mol.
6) If a 25 gram block of copper at 45C is added to 50 grams of liquid ammonia at -60C,
calculate the final temperature of the system. For NH3: FWT(NH3) = 17 g/mol, ΔHvap = 23.4
kJ/mol, Cp(liq)= 35.1 J/mol.K and for Cu : Cp = 24.4 J/mol.K, FWT(Cu) = 63.55 g/mol
7) You have 50 grams of ice at -2.5 C sitting in a dixie cup. If you heat 5 nickels and throw them
into the cup and 8 grams of the ice melts, how hot were the nickels? Each nickel weighs 5 grams
and has a heat capacity of 24.3 J/mol-K. Assume that nickels are pure nickel with an atomic
mass of 57.81 g/mol.
8) What is the ΔG and ΔS for the dilution of 4.5 M H2SO4 to 0.03 M H2SO4 at 25C? The
enthalpy for this reaction is - 210 kJ/mol.
9) Calculate the ΔG for the making of 6M NaOH from solid NaOH. Saturated solutions of
NaOH have a concentration of 19.25 M.
10) Mathematically we write the first law of thermodynamics as U = q + w. The work, w, is
pressure-volume work. What is the sign on this work (+ or -) and why does it have this sign?
11) Please give me a definition of a heat capacity.
12) What is the ΔG for the following reaction at 25C?
BaBr2(aq) + CuSO4(aq) ----> BaSO4(s) + CuBr2(s)
ΔH
BaBr2
CuSO4
BaSO4
CuBr2
S
-186.47 kJ/mol
-201.51 kJ/mol
-345.57 kJ/mol
-25.1 kJ/mol
42.0 J/mol K
19.5 J/mol K
7.0 J/mol K
21.9 J/mol K
Circle all that apply about this reaction,
spontaneous
exothermic
entropy increases
the reaction occurs quickly
not spontaneous
endothermic
entropy decreases
the reaction occurs slowly
What is the equilibrium constant for this reaction?
13) Given the following set of thermodynamic data calculate the ΔG, ΔH, ΔS, and Keq for
the following reaction at 25C (Enthalpies are in kJ/mol and entropies are in J/mol-k)
CoCl2(s) <----> Co2+ + 2 Cl-
CoCl2
Co2+
Cl-
ΔH
-325.5
-67.36
-167.4
ΔS
106.3
-155.2
55.1
Electrochemical Cells
14) Please balance each of the following redox reactions;
MnO4- + Fe2+ —> Fe3+ + Mn2+
CrO42- + I- —> I2 + Cr3+
Co2+ + NO3- —> Co2O3 + NO
I3- + S2O32- —> I- + S4O62H2S + Cr2O72- —> Cr3+ + S
Fe2+ —> Fe3+ + Fe
NiOOH + OH- —> Ni(OH)2 + O2
15) Consider the following electrochemical reaction,
PbI2 + 2 e- ---> Pb + 2 IPb ---> Pb2+ + 2 ePbI2 ----> Pb2+ + 2 IHow would the voltage change if you,
a) Add KI(s) to the oxidation cell
b) Add AgNO3 to the reduction cell
c) Add heat
d) Add water to the reduction cell?
e) Add NaNO3(s) to the oxidation cell
f) Enlarge the oxidation electrode?
Inc.
Inc.
Inc.
Inc.
Inc.
Inc.
N.C.
N.C.
N.C.
N.C.
N.C.
N.C.
Dec.
Dec.
Dec.
Dec.
Dec.
Dec.
16) Consider the following electrochemical reaction,
Au ---> Au3+ + 3eAuCl4- + 3e- ---> Au + 4 Cl-----------------------------AuCl4- ---> Au3+ + 4 ClHow would the voltage change if you,
a) Enlarge the Au electrodes?
Inc.
N.C.
Dec.
b) Add KCl(s)?
Inc.
N.C.
Dec.
c) Add AgNO3?
Inc.
N.C.
Dec.
d) Add water to the reduction cell?
Inc.
N.C.
Dec.
e) Add NaNO3(s)
Inc.
N.C.
Dec.
f) Cool the reaction vessel?
Inc.
N.C.
Dec.
17) Given the following half-reactions draw an electrochemical cell, calculate the voltage of the
cell and lable the anode and cathode. You may use a carbon electrode for the I- solution.
Fe3+ + e- ----> Fe2+
I2 ----> 2 I-
+0.771 eV
+0.535 eV
18) What is the value of the half reaction,
Cu2+ + e- ----> Cu+
given that,
Cu2+ + 2e- ----> Cu E = +0.3402
Cu+ + e- ------> Cu E = +0.522
19) Given the following half-reactions draw an electrochemical cell that would work. Calculate
the voltage of the cell and label the anode and cathode, tell which electrode is positive and which
is negative, and where the oxidation and reduction reactions are occuring. In addition indicate
the direction of electron flow, and the concentration of any ionic species in solution.
Co3+ + e- ---> Co2+ +1.842 eV
Pb2+ + 2e- ---> Pb -0.126 eV
20) Given the following half-reactions, properly set up a WORKING electrochemical cell.
AgI + e----> Ag + I- ΔE = -0.1519 V
PbI2 + 2e----> Pb + 2 I- ΔE = 0.3580 V
Label the anode, the cathode, the direction of electron flow, and indicate the concentrations of all
ionic species. Indicate what the electrodes are made of.
20b) Draw the cell diagram for the above cell.
20c) What is the Keq for the above cell?
20d) Look very carefully at the overall reaction, and explain whether your answer for 7c) is
reasonable or not.
21) Balance the following half-reactions in a BASE and then use these reactions to set up a
WORKING electrochemical cell.
Bi2O3(s) ---> Bi(s)
Hg2O(s) ---> Hg(l)
-0.46 Volts
-0.123 Volts
Label the anode, the cathode, the direction of electron flow, and indicate the concentrations of all
ionic species. Indicate the composition of the electrodes.
21b) Draw the cell diagram for the above cell.
21c) What is the Keq for the above cell?
22) How long will it take to deposit 1 ounce of gold (31.1 g) at a constant current flow of 10
amps? MW (Au) = 197.8
23) How long will it take to deposit 1 lb. of Palladium (453.6 g) from a solution of PdCl4 at a
constant current of 25 amps? MW (Pd) = 106.4
Chemistry 121
Third Exam
Name____________________
May 15, 2007
CLOSED BOOK EXAM - No books or notes allowed. ALL work must be shown for full
credit. You may use a calculator.
Question
Credit
1 ( 18)
2 (40)
3(12)
4(15)
5(15)
Total
R = 8.314 J/mol-K
 =  - 0.0592/n log Q
ΔG = -n
 = 96486 C/mol eΔG = ΔG + RT ln Q
PV = nRT
C = amp x sec
ΔG= -RTlnKeq
R = 0.08205 L-atm/mol-K
1) Balance each of the following oxidation-reduction half-reactions;
IO3- —> I3- (in acid)
Mn(OH)2 —> Mn2O3 (in base)
2) Given the following set of half-reactions, set up a WORKING electrochemical cell.
Hg2Cl2(s) + 2 e- —> 2Hg + 2Cl2 H+ + 2 e- —> H2
 = 0.27 V
 = 0.000 V
Label the anode, the cathode, the direction of electron flow, and indicate the concentrations of all
ionic species. Indicate the composition of the electrodes.
2b) Draw the cell diagram for the above cell.
2c) What is the voltage of the cell when all ionic species have a concentration of 0.10 M?
3) How long will it take to galvanize (coat with Zinc) a nail with 2 grams of zinc using a current
(amperage) of 0.8 amps? FWT of Zn2+ = 65.41 g/mol
4) You are given 100 mL of water at 25C. You add a 20 gram ice cube at 0C and 50 mL of
boiling water at 100C. What is the final temperature of the system?
5) What is the ΔG and ΔS for the dilution of 4.5 M H2SO4 to 0.03 M H2SO4 at 25C? The ΔH =
- 210 kJ/mol.for this reaction.
Final Exam
Comprehensive Exam
~ 30% Cut and Paste from Old Exams
~30% Organic Chemistry
~40% New Questions on Old Topics
Wednesday, May 26th, 10:30-12:30
Chemistry 121
Final Exam
Name
May 27, 1998
CLOSED BOOK EXAM - No books or notes allowed. ALL work must be shown for full credit.
You may use a calculator.
Question
Credit
Question
1(15)
5(20)
2(20)
6(20)
3(20)
7(25)
4(40)
8(40)
Total
Total
Credit
TOTAL POINTS =
I want my final to count 300 pts.______________________________
(This will drop your lowest exam)
Signature
R = 8.314 J/mol-K
ΔE = ΔE - 0.0592/n log Q
ΔG = -nΔE
ln(Ai/Af) = kΔt
 = 96486 C/mol eΔG = ΔG + RT ln Q
PV = nRT
1/Af - 1/Ai = kΔt
C = amp x sec
ΔG= -RTlnKeq
R = 0.08205 L-atm/mol-K
ln(k1/k2) = Ea/R (1/T2 - 1/T1)
1) Write down the equilibrium that occurs when the following compounds are mixed.
a) NaOH + NaBenz
b) NH4OH + NH4Cl
c) NaCN + HCl
2) Last night I poured myself a glass of water from the tap and took a drink. The water was not
cold and didn't taste good so I put a couple of ice cubes in it. The ice cubes melted and the water
tasted better. If my cup had 250 mL of water at 22C, and I put two ice cubes into the water that
weighed 20 grams each, what was the final temperature of the water if the ice cubes were initially
at 0C? ΔHfus = 6.01 kJ/mol, Cp(H2O) = 75.2 J/molK
3) Assume that the cooking of an egg is a first order process. Assume further that it takes 3
minutes to half-cook (soft boil) an egg at 100C. If the activation energy of the cooking of an
egg is 75 kJ/mole, how long would it take to soft boil an egg at the top of Mt. Everest where
water boils at 80C?
4a) Calculate the pH at each of the following points for the titration of 100 mL of 2.00 M
Phthalic acid with 5 M NaOH
H2Phthal ---> H+ + HPhthalHPhthal- ---> H+ + Phthal2-
pKA1 = 2.89
pKA2 = 5.41
INITIAL pH
20 mL NaOH
40 mL NaOH
50 mL NaOH
80 mL NaOH
100 mL NaOH
4b) How many milliliters of 5 M NaOH must you add to 100 mL of 2 M Phthalic acid to make a
buffer of pH = 3.5?
5) At low temperatures, the rate law for the reaction,
CO(g) + NO2(g) ---> CO2(g) + NO(g)
can be determined by the following data;
Initial Rate/10-4 [NO]/10-3 M
3.60
7.20
0.90
0.90
1.20
1.20
1.20
2.40
[CO]/10-3 M
1.50
3.00
0.75
1.50
[NO2]/10-3 M
0.80
0.80
0.40
0.40
Write a rate law in agreement with the data.
Work out the rate law for each of the following mechanisms (show your work) and then circle
the one that is consistent with the rate law determined above. (Circle A, B, or C)
A) CO + NO2 ---> CO2 + NO slow
B) NO2 <---> NO + O equil.
O + CO --> CO2 slow
C) 2 NO2 <---> 2 NO + O2 equil.
CO + O2 ---> CO2 + O slow
O + NO ---> NO2
fast
6) What is the solubility of Barium Carbonate in a solution saturated with H2CO3 at a pH of 10?
[H2CO3]sat'd = 0.034M.
BaCO3 <---> Ba2+ + CO32- Ksp = 8.1x10-9
H2CO3 <---> 2 H+ + CO32- Keq = 2.02x10-17
7) A cell has the following cell diagram;
Ag / 1M AgNO3 // 1M HCl / AgCl / Ag
7a) What is the reduction half reaction?
7b) What is the overall reaction?
7c) Based on the cell diagram please draw the electrochemical cell.
7d) If the Ksp for this cell is 1.8x10-10 then what is the cell voltage when all the ionic species are
0.01 M?
8) Please name or draw the following compounds.
OH
CH3
OH OH
C
H2N
C
C
Cl
O
C
C
C
C
C
C
C
O
C
C
C
C
C
Cl
C
CH3
butanone
3,3-dichloro butanoic acid
trans 1,3-cyclopentadiol
o-aminotoluene
C
Chemistry 121
Final Exam
Name
May 26, 2000
CLOSED BOOK EXAM - No books or notes allowed. ALL work must be shown for full credit.
You may use a calculator.
Question
Credit
Question
1(15)
6(20)
2(15)
7(15)
3(15)
8(30)
4(15)
9(20)
5(25)
10(30 )
Total
Total
Credit
TOTAL POINTS =
I want my final to count 300 pts._________________________________
(This will drop your lowest exam)
Signature
R = 8.314 J/mol-K
ΔE = ΔE - 0.0592/n log Q
ΔG = -nΔE
ln(Ai/Af) = kΔt
 = 96486 C/mol eΔG = ΔG + RT ln Q
PV = nRT
1/Af - 1/Ai = kΔt
C = amp x sec
ΔG= -RTlnKeq
R = 0.08205 L-atm/mol-K
1) Which of the following compounds is the most soluble in water? Show your work.
BaSO4
Fe(OH)2
Ag3PO4
Ksp = 1.98x10-10
Ksp = 1.64x10-14
Ksp = 1.2x10-15
2) Please write the equilibrium reaction that will occur when the following compounds are
mixed. Write all reactions as dissociations where appropriate.
NaOH + HF
Na3Asp + HCl
HCl + KOH
3) The bromination of acetone is acid catalyzed;
CH3COCH3 + Br2 ---> CH3COCH2Br + H+ + BrThe rate of disappearance of bromine was measured for several different concentrations of
acetone, bromine and hydrogen ions;
Rate
6x10-5
2.4x10-5
1.2x10-4
3.2x10-4
8x10-5
[Acetone]
0.30
0.30
0.30
0.40
0.40
[Br2]
0.050
0.100
0.050
0.050
0.050
[H+]
0.050
0.050
0.100
0.200
0.050
a) What is the forward rate law for this reaction?
b) What is the value and units of the forward rate constant?
4) What is the solubility of Cu2S in a saturated solution of 0.10M H2S that is buffered at pH=10?
Cu2S <--> 2 Cu+ + S2- Ksp = 2x10-47
H2S <--> 2H+ + S2Keq = 1.1x10-20
5a) Set up a WORKING electrochemical cell that will determine the solubility product for HgO.
Note that this reaction is spontaneous in the reverse direction. Your overall reaction should be;
HgO + H2O —> Hg2+ + 2 OH-
Label the anode, the cathode, the direction of electron flow, and indicate the concentrations of all
ionic species. Indicate the composition of the electrodes.
5b) Draw the cell diagram for the above cell.
5c) If the  = 1.056 volts, what is the voltage of the cell when concentration of the ionic species
in the reduction cell are 0.5 M and the ionic species in the oxidation cell are 0.2 M?
6a) Assume that the cooking of an egg is a first order process. Assume further that it takes 3
minutes to half-cook (soft boil) an egg at 100C. If the activation energy of the cooking of an
egg is 75 kJ/mole, how long would it take to soft boil an egg at the top of Mt. Everest where
water boils at 80C?
6b) If the rate of a reaction doubles with a 15C change in temperature, what is it's activation
energy? State your assumptions.
7) A block of copper intially at 500C was placed on top of a 100 gram block of ice at 0C. All
of the ice melted and turned into water. The water and the block of copper ended up at 20C.
How much did the block of copper weigh? ΔHfus = 6.01 kJ/mol, Cp(Cu) = 24.5 J/mol-K,
Cp(H2O) = 75.2 J/mol-K, FWT (Cu) = 63.54 g/mol, FWT (H2O)= 18 g/mol
8) Calculate the pH at each of the following points for the titration of 100 mL of 3 M Phthalic
acid with 5 M NaOH
H2Phthal ---> H+ + HPhthalpKA1 = 2.89
HPhthal- ---> H+ + Phthal2- pKA2 = 5.41
INITIAL pH
20 mL NaOH
60 mL NaOH
75 mL NaOH
90 mL NaOH
8b) How much 5 M NaOH would you add to 100 mL of 3 M H2Phthal to make a buffer of pH =
5?
9) How much heat must be added to 100 g of ice at -5C to turn it into 60 grams of water and 40
grams of steam, both at 100C? Cp H2O = 75.7 J/mol-K, Cp Ice = 37.7 J/mol-K, Cp Steam =
35.1 J/mol-K, ΔHfus = 6.01 kJ/mol, ΔHvap = 40.7 kJ/mol