LABORATORY MANUAL CHEMISTRY 121 - Napa Valley … 121/Complete … · N A P A V A L L E Y C O L L E...

LABORATORY MANUALCHEMISTRY 121

FIFTH EDITION

2007Dr. Steven Fawl

LABORATORY MANUAL

CHEMISTRY 121

FIFTH EDITION

Dr. Steven Fawl

Science, Mathematics, and Engineering DivisionNapa Valley College

Napa, California

N A P A V A L L E Y C O L L E G E

COURSE: Chemistry 121

INSTRUCTOR: Dr. Steven Fawl, Room 1843, 253-3195

LECTURES: TW 11:00 to 12:20LAB DEMO: Mon. 11:00 to 11:50TESTS: Th. 11:00 to 11:50LABS: MT 1:30 to 5:20OFFICE HRS: MTW 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 - KineticsExam #2 - Thursday, March 25th - EquilibriumExam #3 - Thursday, May 13th - ThermodynamicsFinal Exam - Comprehensive - Wednesday, May 26th, 10:30-12:30

DESCRIPTION: A continuation of CHEM 120. Topics include solutions, acid-base and redoxequilibria, thermodynamics, kinetics, pH, buffers, solubilityproduct, complex ions, thermodynamics,electrochemistry, biochemistry and nuclear chemistry.

COURSE CONTENT:

1. Chemical KineticsRate LawsActivation EnergyMechanismsCatalysis

2. Chemical EquilibriumLeChatlier's PrincipleHomogenous SystemsHeterogeneous Systems

3. Acids & BasesStrong and Weak acidspHBuffersTitration Curves

4. Applications of Aqueous EquilibriaSolubilitypH controlled SolubilityComplex IonsAmphoterism

5. Spontaniety, Entropy and Free EnergyEffect of TemperatureWork and Efficiency

6. ElectrochemistryNernst EquationStandard State Potentials

7. RadioactivityNuclear StabilityHalf LifeNuclear Transformations

8. Organic ChemistryNomenclatureFunctional GroupsFree-Radical HalogenationSubstitution and Elimination Reactions

COURSE OBJECTIVES:

1. Explain the development of chemical principles and concepts based on experiments.2. 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 socialcontext.

3. Design a laboratory experiment by defining the problem, collecting data, obtainingresults, deriving conclusions, and preparing a report to communicate the information toothers in writing.

4. Explain the concepts related to rates of reaction, activation energies, mechanisms ofreactions, as applied to the kinetic molecular theory.

5. Relate equilibrium information from chemical systems to the free energy, enthalpy andentropy.

6. Determine the equilibrium constants and show how the spontaneity of the system isrelated to the driving force of the reaction.

7. Apply the equilibrium system concepts to acid/base, solubility, redox, and complex ionformation reaction systems.

8. Indicate how an electrochemical cell can be used to establish the standard free energy of achemical reaction, and measure the pH of a system.

STUDENT LEARNING OUTCOMES

1. Communicate chemical and physical processes at the molecular level and how they relate to themacroscopic environment.

2. Solve both qualitative and quantitative chemistry problems while demonstrating the reasoningclearly and completely.

3. Implement laboratory techniques correctlyusing appropriate safetyprocedures and express themclearly in written laboratory reports.

GRADING POLICY: Three exams and a final, quizzes, plus laboratory scores will count towardthe final grade according to the following schedule,

3 Exams = 300pts (100pts each)Final = 200ptsQuizes = 100pts (10 @ 10pts each)Lab = 120ptsTotal = 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 thestudent or else an "F" will be given.

LABS ARE CONSIDERED LATE IF THEY ARE TURNED IN ANY TIME AFTER THEFRIDAY THAT THEY ARE DUE. LABS THAT ARE TWO WEEKS LATE WILL RECEIVENO CREDIT. Special arrangements must be made if a lab must be missed!

LAB EXPERIMENTS AND DATES

DATEEXPERIMENT ONE . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2/1Kinetics of the acid hydrolysis of trans-[Co(en)2Cl2]Cl

EXPERIMENT TWO . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2/8Determination of the t½ of a Radioactive Isotope

HOLIDAY - Presidents Day . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2/15

EXPERIMENT THREE . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2/22Chemical equilibria and Le Chatelier's principle

EXPERIMENT FOUR . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3/1Hydrogen ion concentration and pH of aqueous solutions

EXPERIMENT FIVE . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3/8Neutralization and hydrolysis

EXPERIMENT SIX . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3/15Carbonic acid and its salts

EXPERIMENT SEVEN . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3/22Titration Curve for KHP and Determination of pKa1 and pKa2

HOLIDAY - Spring Break . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3/29

EXPERIMENT EIGHT . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4/5Determination of The Equilibrium Constant For FeSCN2+

EXPERIMENT NINE . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4/12Determination of the Heat of Reaction

EXPERIMENT TEN . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4/19Hess’ Law - Heats of Solution

EXPERIMENT ELEVEN . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4/26Electrochemical cells

EXPERIMENT TWELVE . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5/3Mystery Experiment

L A B O R A T O R Y

There will be twelve lab assignments, approximately one per week. Attendance in lab the firstweek of each session is necessary. Safety goggles are mandatory in the laboratory and can bepurchased in the campus bookstore. Lab Reports are due on Friday the week after the experimentis 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 inyour 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 worninstead of safety goggles.

Following is a description of how to do a lab report for this class. It is suggested that you re-readthis information on a regular basis, and that you follow this outline exactly.

S A F E T Y A N D T E C H N I Q U E R U L E S

Safety in the laboratory is extremely important. It is expected that you know laboratory safetyrules. It is important that if you feel uncomfortable with your knowledge of these rules that youtake 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 safetyand 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 theclass 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 importantthat you ask the Instructor or Instructional assistant.

3) Do not perform any unauthorized experiments. Anyone found doing so faces permanentexpulsion 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 fumehood. 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 orinstructor 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 appropriatewarnings, concentration, date and your name.

21) NEVER place anything inside a reagent bottle, no spatulas, droppers, nor pipets. If thereagent is a clumpy solid inform the IA. Proper technique is to "roll" containers from side toside 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 - offerexcess to another student or dispose of it appropriately.

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23) Be conservative of reagents, place only the amount you need into a labeled container (suchas a beaker). Do not take the reagent bottles to your work area - leave them where every onecan 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 procedurefor special instructions, otherwise dispose of liquids and soluble solids down the sink withlots 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 placeb) clean up your workspace with the spongec) put away your labwared) lock your locker

There is NO reason for a messy lab. Everything you need to keep your lab neat and clean isavailable. Dirty counters, paper towels left in the sink or troughs, labware left out, messesleft under the fume hood, chemical spills left on the balance, are BAD technique and as suchwill not be tolerated.

28) You may not be in the laboratory at any time other than your scheduled laboratory periodunless 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 visityou.

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EXPERIMENT ONEKINETICS OF THE ACID HYDROLYSIS OF trans-[Co(en)2Cl2]Cl

In aqueous solution the green complex trans-dichlorbis (ethylenediamine) cobalt(III) chloridedissociates into chloride ion and trans-[Co(en)2Cl2]

+. This cobalt complex ion then reacts in acidsolution to yield a mixture of the cis and trans forms of [Co(en)2(H2O)Cl]2+. The progress of thisreaction can be followed visually because the products are red while the reactant is green. Thepurpose of this exercise is to determine the rate law for the hydrolysis of trans-[Co(en)2Cl2]

+ andto determine the activation energy for the hydrolysis reaction by carrying out the reaction atseveral different temperatures from 45C to 85C.

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 ofproblems we have had in recent years, the product is sometimes yellow or orange. To determinethe rate law and rate constants for this reaction, we shall measure the half-life for eachexperiment. When 50% of the reactant has been converted to product, the mixture (50% greenand 50% red) has a characteristic color best described as "gun-metal gray", but other colors arepossible 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 0.02g) of trans-[Co(en)2Cl2]Cl, and dissolve them in separate test tubes each containing 8.0 mL of ice-cold 1 M sulfuric acid, then place these and all following solutions in an ice bath. With a pipettransfer exactly 2 mL (save the rest for the next part) from each of these solutions to separate testtubes, and heat these in boiling water for 5 minutes to completely hydrolyze the complex totrans-[Co(en)2(H2O)Cl]+. The solutions should be red, but may be yellow or orange. Cool boththe 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. Theexact color is not important as long as you have a color representing the half-way point in thereaction. Keep these tubes in an ice bath so that they do not undergo additional reaction duringthe laboratory period. Adjust the temperature of a large beaker of water to 55 +/- 1C. Onestudent should place the tubes which contain the remaining unhydrolyzed green complex into thehot water while a second student records the time. As the solutions change color a comparisonshould be made to our standards prepared previously. When the colors match, the time andtemperature should be recorded. From this you should be able to determine the order of thereaction 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 atest tube and boil it for 5 minutes. As before, cool it and add 2 mL of unreacted solution to make4 mL of half converted standard. Using 4 mL portions measure the half life of the reaction at 85,

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75, 65, and 45C. 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 ofthe reaction?

2) The relationship between the rate constant k and the half-life t1/2 for first and second orderreactions 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 ofyour 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 thetemperature in Kelvin.

4) The relationship between the rate constant and a quantity called the activation energy is givenby 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 activationenergy, 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 aslope of -Ea/R and an intercept of ln(A). From the slope and intercept calculate Ea and A. Ea andA are your results.

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PROBLEMS

1) In this experiment you assumed that the t1/2 was very slow at ice bath temperatures of about0C. 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 thisexperiment) decreased as the concentration of his reacting species increased. What if anythingwas wrong?

3) It takes 3 minutes to soft-boil an egg at 100C. Assuming a soft boiled egg represents thehalf-life of the denaturation of egg albumin, comment on why most people add salt to the waterwhen boiling an egg. (What does the salt do to the water?

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Name ______________________ Date _______________

EXPERIMENT ONEKINETICS 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 TWODETERMINING THE HALF-LIFE OF AN ISOTOPE

DISCUSSION

One type of nuclear reaction is called radioactive decay, in which an unstable isotope of anelement changes spontaneously and emits radiation. The mathematical description of this processis shown below.

Af = Aie-kt

In this equation, k is the decay constant, commonly measured in sec-1 (or another appropriate unitof reciprocal time) similar to the rate law constant, k, in kinetics analyses. Ai is the activity (rateof 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, ort1/2. 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 radioactivebarium. The isogenerator contains cesium-137, which decays to produce barium-137. The newlymade barium nucleus is initially in a long-lived excited state, which eventually decays by emittinga gamma photon and becomes stable. By measuring the decay of a sample of barium-137, you willbe able to calculate its half-life.

EXPERIMENT

Each group should have a computer interfaced to a radiation monitor and have the Vernier “Life-time” program running. You will also be given a shallow aluminum cup that will hold yourradioactive sample. Place the radiation monitor on top of, or adjacent to, the cup to get amaximum rate of sample detection.

Your instructor will deliver a small sample of radioactive Barium-137 to each group. As soon asdelivery 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 datacollection. 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 downthe 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 cesiumin addition to the barium. To account for the counts due to any cesium, as well as for counts dueto cosmic rays and other background radiation, you can determine the background count rate fromyour data. By taking data for 30 minutes, the count rate should have gone down to a nearlyconstant value, aside from normal statistical fluctuations. The counts during each interval in thelast five minutes should be nearly the same as for the 20 to 25 minute interval. If so, you can usethe 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 yourmouse. Click on the statistics button on the toolbar. Read the average counts during the intervalsfrom 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 thedata collection.

DETERMINING VALUES FOR k and Ai

Click and drag the computer mouse across the region between 0 and 15 minutes. Make sure thatall 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 backgroundradiation value as the "B" value. This will account for the background radiation inherent in theexperiment. Click Try Fit . The function will be shown and plotted along with your data. Clickto see a full graph of the function and data. Record the fit parameters Ai, k, and B in your datatable.

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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 tonon-barium sources?

3) Would any of your t1/2 or k values change if you had been given more or less sample? Explain.

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EXPERIMENT THREECHEMICAL EQUILIBRIA AND LE CHATELIER'S PRINCIPLE

All chemical reactions proceed toward an equilibrium position, some more rapidly than others. Inthis experiment we shall consider only reactions which occur fairly rapidly, reaching equilibria ina 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 asmall extent.

CH3CO2H H+ + CH3COO-

At equilibrium in 0.1 M acetic acid, only about 1% of the acetic acid molecules are dissociatedinto 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 formCH3COOH molecules, and at equilibrium only about 1% of the H+ and CH3COO- remainunreacted. The final concentrations of H+, CH3COO-, and CH3CO2H are exactly the same as in asolution 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 experimentalconditions 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 willshift. The principle states that;

"If a change is made in any of the factors influencing a system at equilibrium, reactionwill 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 CH3CO2-

on 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 theconcentration 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 portionof 0.1 M acetic acid in a third test tube, and 5 mL of de-ionized water in a fourth. Add 2 drops ofmethyl orange indicator to each test tube and record in your laboratory notebook the color of thecontents of each test tube. To one of the test tubes containing 1 M acetic acid, add 2 mL of 4 Msodium acetate (CH3CO2Na) dropwise. Observe and record the changes in color that occur.

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EQUILIBRIUM BETWEEN A SOLID SALT AND A SOLUTION

As a second example of a rapid reversible reaction we shall investigate the equilibrium between asolid salt and a saturated solution of the salt. The solubility of solid silver acetate (CH3CO2Ag) is0.06 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 thoroughlyusing a glass rod. Record the color of the precipitate that is formed. Place the centrifuge tube inthe centrifuge and balance the centrifuge by placing a second centrifuge tube containing 7 mL ofwater opposite the first tube. Turn on the centrifuge for about 30 secs. When the centrifuge hasstopped, 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 boilingwater and note the amount of solid in the tube. Cool the tube to room temperature by placing it ina beaker of tap water for 5 minutes. Again note the amount of solid. Centrifuge the mixture asabove 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 4M sodium acetate. Stir the mixture with a glass rod. Observe any reaction that occurs. Write anet reaction to account for the observed change. In your conclusion state briefly how thisexperiment illustrates Le Chatelier's Principle. This application of Le Chatelier's principle iscalled the common ion effect.

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Name ______________________ Date _______________

EXPERIMENT THREELeChatelier’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 toHAc

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 =

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PROBLEMS

1) What happens to the H+ concentration in acetic acid when sodium acetate is added? What isthe reaction taking place to account for this? Will the [H+] be higher or lower after the sodiumacetate/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 andAgAc behave differently with addition of H+?

4) How would the equilibrium have shifted in the common ion experiment upon addition of SilverNitrate? of Sodium Nitrate?

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5) Excess solid Ca(OH)2 is in equilibrium with its ions,

Ca(OH)2 Ca2+ + 2 OH-

State 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 thesaturated solution?

7) The solubility of Ca(OH)2 in water is 0.165 g per 100 mL of water. Calculate the solubilityproduct of Ca(OH)2.

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EXPERIMENT FOURHYDROGEN ION CONCENTRATION AND pH OF AQUEOUS SOLUTIONS;

DISSOCIATION CONSTANTS OF ACETIC ACIDAND 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 @ 25C

Chemists use activities in the formulation of Kw instead of concentrations. By convention theactivity of water is set equal to one and the activity of dilute solutions is set equal to theconcentrations 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 anyaqueous 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 pHof pure water would be pH = -log[H+] = -log[10-7] = 7.0. Acidic solutions have a pH less than 7.0and basic solutions have a pH larger than 7.0. THE LOWER THE pH THE HIGHER THEACIDITY.

DETERMINATION OF HYDROGEN ION CONCENTRATION AND pH

In this section you will determine the [H+] and pH of two unknown aqueous solutions. In the backof your LabManual you will find a table which will tell you the relationship between color and pHfor several different indicators. While the data in this table are useful as a general guide, it shouldbe realized that different individuals respond differently to the same color, and they may also usedifferent words to describe a particular color. For the accurate determination of pH with anindicator, one must always make a direct comparison of the color of the unknown solution withthe colors of solutions of known pH values. Furthermore, the solutions being compared should beas similar as possible in factors which may affect the color such as size of test tube, volume ofsolution, and number of drops of indicator. It is usually helpful to observe the colors by placingthe test tubes side by side on a piece of white paper, and then looking down through the solutionsfrom the top.

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EXPERIMENT: Obtain your unknowns from the instructional assistant. Test 2mL your firstunknown with 2 drops of bromothymol blue to ascertain whether it is acidic or basic. Thenproceed with the appropriate indicators to determine the approximate pH. Record the name ofeach indicatorused, its color with your unknown and the pH range indicated into your data table. For these testsuse 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 mustprepare a solution of this pH and test it with the appropriate indicators. If your solution is in the 1M to 10-3 M H+ concentration range (0 to 3 pH), prepare it from 6 M HCl. Solutions in the rangefrom 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 from6 M NaOH. Use serial dilution to prepare solutions from HCl and NaOH. For these final criticalcomparisons use 5 mL of both the standard and the unknown (in different test tubes) and 3-4drops of indicator. Compare the color of the standard solution to the color of your unknown withat 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 basicbecause 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. Determinethe [OH-] in the 1 M NH3 by using the indicators alizarin yellow and indigo carmine. Record thecolors in your data table. Copy this table into your data section and complete it for 1 M NH3. Thistable and all similar tables must be copied into your data section.

Substance H+ OH- NH3 NH4+

EquilibriumConcentration

Using these data calculate the dissociation constant Kb for NH3. Include this in your results.

23

REACTIONS OF NH4+ AND OH-

Predict what will happen when solutions of ammonium chloride (NH4Cl), and sodium hydroxideare 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 severaldrops of indigo carmine to each test tube. Record the color in your data table. Slowly add 5 ml of1 M NH4Cl to one test tube and record the color changes that occur. Does the OH- concentrationincrease 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 yourdata section.

DISSOCIATION CONSTANT OF ACETIC ACID

In the following experiment we shall determine the dissociation constant of acetic acid by mixingsolutions containing equal numbers of moles of acetic acid and sodium acetate. The resultingbuffer 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 isequal 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+] ofthe resulting solution using the appropriate indicators. Record the name of each indicator and itscolor in this solution in your notebook. In your notebook, copy and complete the following tablefor the solution (which does not contain 0.5 M HAc and 0. 5 M NaAc).

Substance H+ OH- HAc Ac- Na+


Using these data, calculate the dissociation constant Ka and pKa for acetic acid. Include thisnumber in your results.

24

PROBLEMS

1. A solution is prepared by dissolving 0.050 mole of HCl and 0.030 mole of NaOH in sufficientwater 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 wecould 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 thefollowing problems.

25

4. A solution is prepared by mixing 0.050 mole of NH4Cl and 0.010 mole of NaOH in sufficientwater to give a final volume of 200 mL. Calculate the concentration of all the molecular and ionicspecies in the resulting solution.

5. A solution is prepared by mixing 0.050 mole of NH3 and 0.20 mole of HCl in sufficient waterto give a final volume of 500 mL. Calculate the concentration of all the molecular and ionicspecies in the resulting solution.

26

EXPERIMENT FIVENEUTRALIZATION AND HYDROLYSIS

The neutralization of an acid by a base to form a salt and water proceeds to an equilibriumposition which varies depending on the strengths of the acid and base that are used. In thisassignment 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 equivalentamounts of the reactants (as in a titration) or by preparing an aqueous solution of the product (thesalt of the acid and base) of equivalent concentration. Since it is much easier experimentally toprepare the salt solutions, we shall prepare them and measure their pH and [H+], from which wecan infer the equilibrium position. We shall therefore study the above reactions in the reversedirection. The reverse of the neutralization reaction is called the "hydrolysis" reaction, because inthe 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 blueindicator. It should give a green color. If it is yellow, boil the water until it gives a green colorwith bromothymol blue. If your solution is blue inform the instructional assistant. Prepare anapproximately 1 M NaCl solution by dissolving about 1.2 grams of sodium chloride (comparewith the sample in the laboratory) in 20 mL of your distilled water. Determine the pH and [H+] ofyour NaCl solution by using appropriate indicators, including bromothymol blue. Use the sameprocedure you used in the preceding assignment for determining an unknown pH. (Record namesand colors of indicators in your laboratory notebook.) Copy in your laboratory notebook andcomplete the following table for a 1 M NaCl solution.

Substance H+ OH- Na+ Cl- HCl NaOH


27

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 + OH-

EXPERIMENT: 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 yourlaboratory notebook.) From your measured value of the pH and Kw, calculate [H+] and [OH-] in 1M CH3CO2Na. Using this value for [OH-], calculate [CH3CO2

-]. Copy in your laboratorynotebook and complete the following table for 1 M CH3CO2Na.

Substance H+ OH- Ac- HAc Na+ NaOH


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 appropriateindicators, including bromocresol green and bromothymol blue. (Record names and colors ofindicators 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.

Substance H+ OH- NH4+ Cl- NH3 HCl


HYDROLYSIS OF SODIUM CARBONATE SOLUTIONS

The neutralization reaction for bicarbonate ion, HCO3- , by a strong base is

HCO3- + OH- CO3

2- + H2O

The hydrolysis of carbonate ion, CO32- is represented by the reverse of the above reaction

CO32- + H2O HCO3

- + OH-

EXPERIMENT: Determine the pH and [H+] of 1 M sodium carbonate by using appropriateindicators, including alizarin yellow and indigo carmine. (Record names and colors of indicatorsin 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- Na+ CO32- HCO3

-


29

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


Weak Acid and Strong Base: Complete the following table for 1 M CH3CO2Na.

pH =

Substance H+ OH- Ac- HAc Na+ NaOH


Weak Base and Strong Acid: Complete the following table for 1 M NH4Cl.

pH =

Substance H+ OH- NH4+ Cl- NH3 HCl


Weak Diprotic Acid and Strong Base: Complete the following table for 1 M Na2CO3.

pH =

Substance H+ OH- Na+ CO32- HCO3

-


30

PROBLEMS

1) a) Is the [H+] of the NaCl solution greater than, less than or equal to that of your boiled distilledwater? b) Which indicator would you choose for the most exact endpoint for the titration of anHCl solution with NaOH?

2) a) Would you expect a solution of sodium acetate to be acidic, basic or neutral? b) What is thepH of sodium acetate? c) Which indicator would you choose for the titration of CH3CO2H withNaOH?

3) a) Would you expect a solution of ammonium chloride to be acidic, basic or neutral? b) Whatis 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 theequilibrium 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 ammoniumacetate is tested with bromothymol blue, the solution is green. What substances are present in a 1M NH4CH3CO2 solution, include all ions and molecules? Write an equation for the net reactionfor the hydrolysis of NH4CH3CO2. Calculate the equilibrium constant for this reaction, using theinformation given in problem #5.

31

EXPERIMENT SIXCARBONIC ACID AND ITS SALTS

Carbon dioxide is a colorless gas at room temperature which is somewhat soluble in water. Waterin 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 willuse a carbon dioxide generator to make saturated solutions of H2CO3. Carbonic acid is a typicaldiprotic acid in that it dissociates stepwise as shown below,

H2CO3 H+ + HCO3- Ka1 = 4.3 X 10-7

HCO3- H+ + CO3

2- Ka2 = 4.7 X 10-11

Thus we see that there are two possible salts, NaHCO3 (sodium bicarbonate) and Na2CO3 (sodiumcarbonate). 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 fewmL of concentrated H2SO4 to some solid sodium bicarbonate inside a CO2 generator found in thelab. Allow the resulting gas to bubble through 100 mL of distilled water for about 5 minutes. Testthe resulting solution with bromocresol green and methyl orange indicators. Record the pH andlabel the solution 0.034 M H2CO3. Note- the pH of the solution is due almost entirely to the firstacid dissociation.

EXPERIMENT: Bubble some CO2 through a 10 mL sample of 1M CaCl2. Does a precipitateform? If we assume that the solution is saturated with CO2, what is the concentration of H2CO3 inthe 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-


SODIUM CARBONATE SOLUTIONS

Sodium carbonate, Na2CO3 is a solid that is very soluble in water and it is a strong electrolytewhich dissociates to give Na+ and CO3

2- in aqueous solutions. Sodium carbonate solutions canhydrolyze stepwise in the following manner,

CO32- + H2O HCO3

- + OH-

HCO3- + H2O H2CO3 + OH-

EXPERIMENT: 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. Writean 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 electrolytewhich dissociates to give Na+ and HCO3

- in aqueous solutions. Sodium bicarbonate solutions mayalso be prepared by the reaction of solutions containing 1 mole of H2CO3 and 1 mole NaOH. Theequation for the net reaction is,

H2CO3 + OH- HCO3- + H2O

Since HCO3- is a weak acid, it should dissociate to give H+ and CO3

2-. However the H+ producedwould react with another HCO3

- to form the weak acid H2CO3. As a result we predict that theprincipal equilibrium in a NaHCO3 solution would be,

2 HCO3- H2CO3 + CO3

2- *

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 asecond centrifuge tube. Now, place the tube containing the decanted solution in a beaker ofboiling water. Keep it in the boiling water for 10 minutes, and then cool it by placing it in a beakerof water at room temperature. After it has cooled to room temperature centrifuge it. Note theamount 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 themixture 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 mLportions 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, andplace 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 (exceptfor 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 beforeheating). Now for each of the 3 solutions list the ions in order of decreasing concentration. It willbe 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- CO3

2-


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 becausethey will react together. For these cases write a net equation for the reaction that would occur. Inother cases write "No Reaction".

(a) OH- and HCO3- (e) Ca2+ and HCO3

-

(b) OH- and CO32- (f) H2CO3 and CO3

2-

(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 equilibriumconstant 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 SIXTitration Curve for Potassium Acid Phthalate

INTRODUCTION:

In this experiment a pH meter equipped with a combination electrode (a glass electrode and acalomel reference electrode in the same housing) and drop counter will be used to construct thetitration curve for KHP titrated with NaOH. From the titration curve and the concentration of thebase you can determine the concentration of the unknown KHP solution and both the first and thesecond 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 activitiesare 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 thenearest neighbor shell around a cation will contain a (slight) predominance of anions, and viceversa. This partial screening of the ions by their counterions makes the effective concentration a ofan ionic species somewhat smaller than its stoichiometric concentration. The activity andconcentration 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 theion and the total ionic strength of the solution (cf. Chapter 8).

Electrochemical experiments, such as the measurement of pH, yield ion activities rather than ionconcentrations. At the ionic-.strengths to be used in this experiment the activity coefficients of thephthalate anionic species differ appreciably from unity, and a serious misestimate of Ka results ifthey 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 andpKa 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 thecurve is a maximum. We can use this fact to find the equivalence point very precisely. You willconstruct a "differential" titration curve, a plot of (ΔpH/ΔV) vs. ΔV, which will be very sharplypeaked. 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 andlocker 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 withyour standard 0.1 M NaOH..

C. Just before you are ready to prepare a titration curve standardize the pH meter according to thedirections given in a separate handout.

D. Using a pipet deliver 25 mL of the unknown KHP solution into a clean, dry 100 mL beakerwhich already contains a clean magnetic stirring bar. Add 20 mL of distilled water and 2 dropsof phenolphthalein indicator. Put the beaker onto the magnetic stirrer so that the bar rotatesfreely, but there is still room alongside it for the electrode to be inserted near the bottom of thebeaker. Mix the solution thoroughly with the stirrer.

E. Wash the standardized electrode with your wash bottle, blot it dry, and position it throuh thelarge 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 collectingdata. Click the Collect button. No data will be collected until the first drop goes through thedrop counter slot. Adjust the drop rate to about 1 drop every 2 seconds. When the first droppasses 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 theequivalence point. When this jump in pH occurs, add about 3 more milliliters of NaOH, clickSTOP, 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 takethe 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 equivalencepoint.

I. From this curve determine the pH at half and three-fourths neutralization of the KHP. Go tothe worksheet and use it to compute an average value for pKa2, taking into account theactivity coefficients. α’s (Na+ = 4, P2- = 6, HP- = 6, K+ = 3). The Debye-Huckel equation isbelow.

J. Having Ka2 in hand you can obtain Ka1 from the initial pH of the diluted unknown. In asolution 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 secondterm and can be neglected, similarly the Ka1Kw term in the numerator is very small by comparisonwith 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 fP2-

The ionic strength of your diluted starting solution is easily calculated from the concentration ofyour unknown found in step 11. The activity coefficient of P2- at this ionic strength can beinterpolated from the table in your text. With this value, the measured pH, and pK2 which youhave calculated in step 12, pK1 is in hand.

REPORT: Turn in your titration curve, the concentration of the undiluted KHP unknown, and thecalculations of pKa1 and pKa2 from steps 12 and 13, using equations 7 and 13.

40

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41

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 willreact to form products until a state is reached whereby the concentrations of the reactants andproducts remain constant. This is a dynamic state in which the rate of formation of the products isequal to the rate of formation of the reactants. The reactants and products are in chemicalequilibrium and will remain so until affected by some external force. The equilibrium constant Kcfor 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) ionand 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+ andSCN- will decrease: so for every mole of FeSCN2+ that is formed, one mole of Fe3+ and one mole ofSCN- will react. The equlibrium constant expression Kc, according to the Law of ChemicalEquilibrium, 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+ andSCN– will react until the above equation is satisfied, so that the same value of the Kc will beobtained no matter what initial amounts of Fe3+ and SCN– were used. Our purpose in thisexperiment will be to find Kc for this reaction for several mixtures made up in different ways, andto 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 thered color of the FeSCN2+ ion makes for an easy analysis of the equilibrium mixture using aspectrophotometer. The amount of light absorbed by the red complex is measured at 447 nm, thewavelength at which the complex most strongly absorbs. The absorbance, A, of the complex isproportional to its concentration, M, and can be measured directly on the spectrophotometer usingthe Beer-Lambert law:

A = M

Where = extinction coefficient, M = Molarity, and = path length (1cm)

42

EQUIPMENT

5 test tubes10 ml graduated cylinder10 ml graduated measuring pipette (0.1 ml graduations)stirring rodsmall labels or markers5 cuvettes for the spectrophotometers250 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/samplesolutions ensure that the flask is wet with diluant (in this case it is the 0.0200M Fe(NO3)3 in 1.0 MHNO3 ).

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 0.0200MFe(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 0.5M 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 willallow us to determine its equilibrium concentration. Knowing the initial composition of a mixtureand the equilibrium concentration of FeSCN2+, we can calculate the equlibrium concentrations ofthe 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. Pourabout 30 mL 0.02 M Fe(NO3)3 in 0.5 M HNO3 into a dry 100 mL beaker. Pipet 5.00 mL of thatsolution into each test tube. Then add about 20 mL 2.00 x 10-4 M KSCN to another dry 100-mLbeaker. Pipet 1,2,3,4, and 5 mL from the KSCN beaker into each of the corresponding test tubeslabeled 1 to 5, then pipet the proper number of milliliters of 0.5 M HNO3 into each test tube tobring 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.

TestTube #

Reagents (in mLs) 1 2 3 4 5

Fe(NO3)3 5.00 5.00 5.00 5.00 5.00

KSCN 1.00 2.00 3.00 4.00 5.00

HNO3 4.00 3.00 2.00 1.00 0.00

Mix each solution thoroughly with a glass stirring rod. Be sure to dry the stirring rod after mixingeach solution to prevent cross-contamination.

Measure the absorbance of each mixture at 447 nm as demonstrated by your instructor and put thedata in the worksheet provided.

Determine the concentration of FeSCN2+ from your calibration curve. Record the value on yourReport 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 For Beer-Lambert Plot

Data 5 mL 10 mL 15 mL

Absorbance

Data for Equilibrium Calculation

Test Tube #

Data 1 2 3 4 5

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

EXPERIMENT NINETHE 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 waterand as a consequence there is a volume change when ice melts. This volume change can be measuredand used to calculate how much energy was transferred to the ice. This in turn is used to measure theamount of energy released in the reaction (the enthalpy).

EXPERIMENT: Students may work in pairs with one student calling out the time and the otherstudent reading the volume and recording both the time and volume. The calculations are to be madeindividually. Fill a styrofoam container with crushed ice from the ice provided in the lab and placein 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 pushthe stopper in tightly. Water will spurt out the top of the pipet. Watch the water level in the pipet overa 5 minute period. If it does not drop, the apparatus is free of leaks and you may continue with theexperiment. If the level drops, try to repair the leak with the aid of your instructor or instructionalassistant. ( Do not use grease on the stopper.) When your assembly is leak-proof, fill the bottle withcrushed ice to the brim. Add ice water from the styrofoam container to fill the bottle completely andinsert 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 andwater to just below the rim of the test tube. Insert the cork in the test tube loosely. Allow theapparatus 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. Recordthe mass in your notebook.

After the apparatus has been allowed to stand for 15 minutes, adjust the water level in the pipet to areading of 0.8 mL or greater by placing a small piece of latex tubing on the end of the pipet and fillingit with water. By pushing on the stopper it is possible to raise the level of the water inside of the pipetuntil it meets the water in the tubing. Releasing the stopper will pull the water from the tubing intothe pipet thus filling it. This technique may require practice. Begin reading the pipet volume once thewater level has reached some convenient mark on the pipet. Read the pipet volume each minute for5 minutes and record these readings in your notebook. If the volume change in the first 5 minutes isgreater than 0.05 mL then you have a leak and must readjust the stopper and start again at thebeginning of this paragraph. If the volume change during this 5 minute period is 0.05 mL or less, rollthe magnesium ribbon into a loose coil, and drop it into the HCl solution. Place the cork LOOSELYover 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 resultsshould be similar to those shown in the figure below.

CALCULATIONS: Draw parallel lines through the initial and final points. The volume change dueto 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 ofice take your measured volume and multiply by 0.875 Kcals/mL. The result of this calculation givesyou the amount of energy released when about 0.1 gram of magnesium reacts. To calculate the amountof energy released when one mole of magnesium reacts you must calculate the number of moles ofmagnesium used in your experiment. Calculate the number of moles of magnesium used in yourexperiment (the atomic mass of magnesium is 24.3 g/mole). By dividing the number of caloriescalculated earlier and dividing this by the number of moles of magnesium used one obtains theenthalpy 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, calculatethe standard heat of formation of MgCl2 (aq). (Hint: What is meant by the standard heat of formation?)

47

EXPERIMENT ELEVENELECTROCHEMICAL CELLS

In this experiment you will investigate reactions in which electrons are transferred from one reactantto another. Since electrons are being transferred we can force them to travel through an electricalconnection and thereby be measured. We call these set-ups electrochemical cells. We can observethe effect of electron transfer in other ways. In the first experiment we will observe the formation ofelements by electron transfer.

EXPERIMENT: For each of the following experiments, record your observations and write an netequation 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 brominethen add 2 mL of 6 M NH3. Using NH3 allows you to see the blue color of the copper ions insolution.

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 thesurface 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 reactionwill occur when you take a reaction and reverse any of the reactions above it (reversing a reactionmakes it an oxidation). It is sometimes difficult to see the lead reaction with H+, a reaction does occurso place the lead appropriately. What additional experiments would you need to perform to placebromine and iodine in their proper position?

48

ELECTROCHEMICAL CELLS

In this portion of the experiment you will be measuring the voltage produced by three electrochemicalcells. 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 cellselectrons flow from the anode to the cathode and the anode will be negatively charged compared tothe 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 MZnSO4 into a porous cup. Place the porous cup into the 250 mL beaker containing the CuSO4. Put acopper strip into the copper sulfate solution and a zinc strip into the zinc solution. Attach either avoltmeter to the zinc and copper electrodes and measure the voltage. Draw the cell and determinewhich 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 replacethe copper sulfate solution and the copper electrode. Measure the voltage. Draw the cell and determinewhich 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 thislab you will calculate the Ksp of Cu(OH)2 by the following set of half-reactions,

Cu(s) + 2 OH- <-----> Cu(OH)2(s) + 2e-

Cu2+(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 beakerand 40 mL of 1 M NaOH in the cup. Add 4 drops of 1 M CuSO4 to the 40 mL of NaOH. Place acopper strip into each cells and measure the voltage as previously. Note the negative terminal. Whichelectrode 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 willoccur as written, or will occur in the opposite direction.

a) I2 + Zn(s) = Zn2+ + 2 I-

b) 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 CRCHandbook, (Na, K, Mg, Al, Zn, Fe, Sn, Pb, Cu, and Ag) by placing the most positive one at the topof 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 MHCl?

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? Whyor why not?

50

EXPERIMENT TWELVEThe Mystery Experiment

51

1 2 3 4 5 13 147 8 9 10 11 1260

Malachite Green

Methyl Violet

Methyl Orange

Bromocresol Green

Methyl Red

Bromothymol Blue

Cresol Red

Thymol Blue

Phenolphthalein

Alizarin Yellow R

Indigo Carmine

Y B V C

Y V

R Y

Y B

R Y

Y B

R Y V

R Y B

R

O

C

Y

B Y

pH Color Chart

B = Blue Y = Yellow V = Violet C = Clear O = Orange R = Red

Exam IRate Laws

Activation EnergiesMechanisms

Radioactive Decay

Kinetics and Activation Energy

1) Rate information was obtained for the following reaction at 25C and 33C;

Cr(H2O)63+ + SCN- ---> Cr(H2O)5NCS2+ + H2O(l)

Initial Rate [Cr(H2O)63+] [SCN-]

2.0x10-11 1.0x10-4 0.102.0x10-10 1.0x10-3 0.109.0x10-10 2.0x10-3 0.152.4x10-9 3.0x10-3 0.20

@ 33C 2.0x10-11 1.0x10-4 0.20

a) Write a rate law consistent with the experimental data.

b) What is the value of the rate constant at 25C?

c) What is the value of the rate constant at 33C?

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. Cl2(g) <---> 2 Cl(g)slow COCl2(g) + Cl(g) ---> COCl(g) + Cl2(g)fast eq. 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 [ClO2]/M Rate/(M/sec)

0.10 0.010 1.2x10-3

0.10 0.040 4.8x10-3

0.20 0.010 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+ + Br-

The rate of disappearance of bromine was measure for several different concentrations ofacetone, bromine and hydrogen ions;

Rate [Acetone] [Br2] [H+]6x10-5 0.30 0.050 0.0506x10-5 0.30 0.100 0.0501.2x10-4 0.30 0.050 0.1003.2x10-4 0.40 0.050 0.2008x10-5 0.40 0.050 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 1.50 0.80 3.601.20 3.00 0.80 7.201.20 0.75 0.40 0.902.40 1.50 0.40 0.90

Write a rate law in agreement with the data.

Which of the following mechanisms is consistent with this ratelaw? (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 slowO + NO ---> NO2 fast

D) None of the above

6) Given the following rate data, calculate the rate law at 25C.

Rate/10-4 [I-]/M [OCl-]/M

6.10 0.20 0.05012.2 0.40 0.05024.4 0.20 0.100

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 [ClO2]/M Rate/(M/sec)

0.10 0.010 1.2x10-3

0.10 0.040 4.8x10-3

0.20 0.010 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 [NO]/10-2 M [H2]/10-2 M

0.60 0.50 0.202.40 1.00 0.200.30 0.25 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 toaccount for the presence of the catalyst?

f) Would platinum be a homogeneous or heterogeneous catalyst for this reaction?

If the temperature is increased by 6.8C the reaction rate increases 1.65 times. What is theactivation energy of the reaction?

9)The following experimental data were obtained for the reaction at 25C,

2 NO + H2 ---> N2O + H2O

Initial [NO] Initial [H2] Initial Rate(x103 M) (x103 M) (M/s x 105)

6.40 2.20 2.6012.8 1.10 5.206.40 4.40 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)


CO(g) + NO2(g) ---> CO2(g) + NO(g)

is;Rate = k [NO2]

2

Which of the following mechanisms is consistent with this ratelaw? (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 slowN2O4 <----> 2 NO + O2 fast equil.O2 + CO2 ---> 2 CO2 fast

11) Given the following rate data, calculate the rate law at 25C.

Rate/10-4 [I-]/M [OCl-]/M@ 25C

6.10 0.20 0.05012.2 0.40 0.05036.6 0.60 0.100

@ 33C14.4 0.20 0.050

What is the rate law?

Rate =

What is the value of the rate constant at 25C and at 33C?

What is the activation energy for this reaction?

What is the half-life of this reaction at 45C if [I-] = [OCl-] = 0.25 M?

12) Thiosulfate (S2O32-) can react with triiodide (I3

-) according to the following reaction at 25C,

2 S2O32- + I3

- <----> S4O62- + 3 I- Keq = 3.75x105

Experimentally the forward rate law can be determined from the following data,

Rxn Rate/Msec-1 [S2O32-]/M [I3

-]/M

#1 2.56x10-4 0.040 0.12#2 1.28x10-4 0.020 0.12#3 1.92x10-4 0.060 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 35C the rate increases to 4.8x10-4 M/sec. What is the activationenergy 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 100C in an open sauce pan and only12 minutes to boil them in a pressure cooker at 110C. Estimate the activation energy forcooking potatoes, which involves the conversion of cellulose into starch. Remember that thereis 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 leftafter 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 anEgyptian mummy. The cloth was found to have a decay rate of 9.1 dpm. If the decay rate is 15.3dpm 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. Apiece of zirconium was dated using this process and it was found that this rock contained3.2x10-3 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 half-life of 35 days. Compute the number of days required to forget 90% of the material that youlearned in preparation for this exam.

18) Under acidic conditions sucrose (table sugar) can be broken down into its individual sugars,glucose and fructose. At 27C it takes 54.5 minutes to convert half the sucrose to glucose andfructose and at 37C it takes 13.7 minutes. Estimate the activation energy for the breakdown ofsucrose.

18b) The above reaction is known to second order which means that the half-life is dependent onthe 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 usingthe Arrhenius equation?)

19) A 10 gram sample of 131I was sent from a pharmaceutical company to a hospital for use in thetreatment of hyperthyroidism. If the half life of 131I is 8.07 days, how much of the sample wouldbe left after a 2 day mail delivery?

20) The denaturation of the virus that causes the rabbit disease Myxomatosis can be followed byheating the virus under a microscope. It is observed that the reaction is first-order and that ittakes 22.35 minutes at 50C and 0.35 minutes at 60C for the virus to denaturate. Estimate theactivation energy for the denaturation of the Myxomatosis virus.

Exam IIChemical Equlibria

LeChatlier’s PrincipleAcid DissociationBase HydrolysisTitration Curves

Solubility ProductpH Controlled Solubility

Amphoterism

Equilibria and LeChatliers Principle

1) For each of the following sets of compounds write the equilibrium reaction that would occurwhen the compounds are mixed together.

a) HNO3 and NaBenzb) NH4Cl and KOHc) NaCN and NaOH

2a) If H2 and Cl2 are added to a container, both at 2 atm, what will the pressure of HCl be afterthe 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 beenadded to the container originally?

3) If 100 grams of NaF and 70 grams of KOH are added to 250 mL water what is the equilibriumconcentration 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 MKOH? Ka = 6.46x10-5

6) What is the pH of a solution where 100 mL of 0.50 M NH4OH is completely neutralized by 65mL 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 (totalpressure) 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 aceticacid? Ka = 1.8x10-5

11) In class we noted that a 0.10 M solution of calcium ions can be precipitated with 0.10 MNaOH but not in 0.10 M NH4OH. The solubility product for Ca(OH)2 is 7.88x10-6. Pleaseexplain 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 HClare 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 and0.25 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 MNaOH 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 NaOHare added to 160 mL of 3 M dl Aspartic acid?

H2Asp <---> H+ + HAsp- 1.38x10-4

HAsp- <---> H+ + Asp2- 1.51x10-10

0 mL200 mL400 mL700 mL800 mL900 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 bufferof pH 4.5?

16) What is the pH of the solution when 75 mL of 0.8 M dl-Histidine is titrated with thefollowing volumes of 1.20 M NaOH?

H3His <---> H+ + H2His- pKa1 = 2.40H2His- <---> H+ + HHis2- pKa2 = 6.04HHis2- <---> H+ + His3- pKa3 = 9.33

0 mL30 mL50 mL75 mL100 mL125 mL135 mL150 mL175 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 makea 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 bufferof 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 acid basic neutral can't tellCr(NO3)3 acid basic neutral can't tellKCl acid basic neutral can't tellNH4CN acid basic neutral can't tell

22) Are the following compounds soluble or insoluble in water?

NaIO3 Soluble Insoluble FeSO4 Soluble InsolubleCr(OH)3 Soluble Insoluble ScCl3 Soluble InsolublePbSO4 Soluble Insoluble Na2O Soluble InsolubleCaCl2 Soluble Insoluble PbSO4 Soluble 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 equilibriumconcentration 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 forPbCl2, 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+ + S2- Keq = 1.1x10-20

34) What is the solubility (silver ion concentration) of silver sulfide (Ag2S) in a saturatedsolution of H2S = 0.10 M at pH = 3?

Ag2S <--> 2 Ag+ + S2- Ksp = 1.6x10-49

H2S <--> 2H+ + S2- Keq = 1.1x10-20

35) How much of each precipitate will form and what will the concentration of lead be if0.75 mole sodium oxalate (Na2C2O4 = Na2Ox) is added to 100 mL of solution containing 0.30 MMg(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+ + CO32- Ksp = 8.1x10-12

H2CO3 <--> 2 H+ + CO32- Ka = 2.02x10-17

Chemistry 121 NameSecond Exam 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 aremixed. 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 of0.6 M NaOH

H2CO3 <---> H+ + HCO3- Ka1= 4.3x10-7

HCO3- <---> H+ + CO3

2- 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 solublein acid solutions and FeS is soluble. Draw and approximate ion separation curve for CuS andFeS.

5b)Define amphoterism and give an example.

Exam IIIThermodynamics

1st, 2nd, and 3rd LawsEnthalpy, Entropy, and Gibbs Energy

Heats of FormationHess’ LawSpontaniety

Gibbs EquationElectrochemical Cells

Balancing Redox ReactionsNernst Equation

BatteriesElectroplating and Hydrolysis

Thermodynamics

1) A 100 gram block of copper was heated with a bunsen burner and then thrown into a cup ofwater (250 mL) initially at 25C. When the system reached thermal equilibrium the finaltemperature of the water and the block was 61C. What was the initial temperature of the blockof 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/molC + O2 -----> CO2 ΔH = -393.5 kJ/molH2 + ½ O2 -----> H2O ΔH = -241.8 kJ/mol

3) When ethanol (C2H5OH) burns in oxygen it forms CO2 and water. If the heat of combustionis -1237.7 kJ/mol, the heat of formation of water is -241.8 kJ/mol, and the heat of formation ofcarbon dioxide is -393.5 kJ/mol, what is the heat of formation of C2H5OH?

4) Given a block of ice at 0C how much of the ice will melt if you put a 100 gram block ofaluminum on it that is initially at 100C? Cp(Al)=24.1 J/mol-K, Cp(H2O)=75.2 J/mol-K, Heatof 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 excessoxygen. 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 8.6C. What is the heat of formationof the diethyl ether? Cp(H2O) = 75.2 J/mol-K, ΔHf(CO2) = -393.5 kJ/mol, ΔHf(H2O) = -241.8kJ/mol, FWT (Diethyl ether) = 74 g/mol, FWT (water) = 18 g/mol.

6) If a 25 gram block of copper at 45C is added to 50 grams of liquid ammonia at -60C,calculate the final temperature of the system. For NH3: FWT(NH3) = 17 g/mol, ΔHvap = 23.4kJ/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 theminto the cup and 8 grams of the ice melts, how hot were the nickels? Each nickel weighs 5 gramsand has a heat capacity of 24.3 J/mol-K. Assume that nickels are pure nickel with an atomicmass 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 25C? Theenthalpy for this reaction is - 210 kJ/mol.

9) Calculate the ΔG for the making of 6M NaOH from solid NaOH. Saturated solutions ofNaOH have a concentration of 19.25 M.

10) Mathematically we write the first law of thermodynamics as U = q + w. The work, w, ispressure-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 25C?

BaBr2(aq) + CuSO4(aq) ----> BaSO4(s) + CuBr2(s)

ΔH S

BaBr2 -186.47 kJ/mol 42.0 J/mol KCuSO4 -201.51 kJ/mol 19.5 J/mol KBaSO4 -345.57 kJ/mol 7.0 J/mol KCuBr2 -25.1 kJ/mol 21.9 J/mol K

Circle all that apply about this reaction,

spontaneous not spontaneousexothermic endothermicentropy increases entropy decreasesthe reaction occurs quickly 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 forthe following reaction at 25C (Enthalpies are in kJ/mol and entropies are in J/mol-k)

CoCl2(s) <----> Co2+ + 2 Cl-

ΔH ΔSCoCl2 -325.5 106.3Co2+ -67.36 -155.2Cl- -167.4 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- + S2O3

2- —> I- + S4O62-

H2S + Cr2O72- —> Cr3+ + S

Fe2+ —> Fe3+ + Fe

NiOOH + OH- —> Ni(OH)2 + O2

15) Consider the following electrochemical reaction,

PbI2 + 2 e- ---> Pb + 2 I-

Pb ---> Pb2+ + 2 e-

PbI2 ----> Pb2+ + 2 I-

How would the voltage change if you,

a) Add KI(s) to the oxidation cell Inc. N.C. Dec.b) Add AgNO3 to the reduction cell Inc. N.C. Dec.c) Add heat Inc. N.C. Dec.d) Add water to the reduction cell? Inc. N.C. Dec.e) Add NaNO3(s) to the oxidation cell Inc. N.C. Dec.f) Enlarge the oxidation electrode? Inc. N.C. Dec.

16) Consider the following electrochemical reaction,

Au ---> Au3+ + 3e-

AuCl4- + 3e- ---> Au + 4 Cl-

------------------------------AuCl4

- ---> Au3+ + 4 Cl-

How 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 thecell and lable the anode and cathode. You may use a carbon electrode for the I- solution.

Fe3+ + e- ----> Fe2+ +0.771 eVI2 ----> 2 I- +0.535 eV

18) What is the value of the half reaction,

Cu2+ + e- ----> Cu+

given that,

Cu2+ + 2e- ----> Cu E = +0.3402Cu+ + e- ------> Cu E = +0.522

19) Given the following half-reactions draw an electrochemical cell that would work. Calculatethe voltage of the cell and label the anode and cathode, tell which electrode is positive and whichis negative, and where the oxidation and reduction reactions are occuring. In addition indicatethe direction of electron flow, and the concentration of any ionic species in solution.

Co3+ + e- ---> Co2+ +1.842 eVPb2+ + 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 allionic 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) isreasonable or not.

21) Balance the following half-reactions in a BASE and then use these reactions to set up aWORKING electrochemical cell.

Bi2O3(s) ---> Bi(s) -0.46 VoltsHg2O(s) ---> Hg(l) -0.123 Volts

Label the anode, the cathode, the direction of electron flow, and indicate the concentrations of allionic species. Indicate the composition of the electrodes.


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 10amps? 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 aconstant current of 25 amps? MW (Pd) = 106.4

Chemistry 121 Name____________________Third Exam May 15, 2007

CLOSED BOOK EXAM - No books or notes allowed. ALL work must be shown for fullcredit. You may use a calculator.

Question Credit

1 ( 18)

2 (40)

3(12)

4(15)

5(15)

Total

R = 8.314 J/mol-K = 96486 C/mol e- C = amp x sec = - 0.0592/n log Q ΔG = ΔG + RT ln Q ΔG= -RTlnKeqΔG = -n PV = nRT 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 + 2Cl- = 0.27 V2 H+ + 2 e- —> H2 = 0.000 V



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 25C. You add a 20 gram ice cube at 0C and 50 mL ofboiling water at 100C. 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 25C? The ΔH =- 210 kJ/mol.for this reaction.

Final ExamComprehensive 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 NameFinal Exam May 27, 1998


Question Credit Question Credit

1(15) 5(20)

2(20) 6(20)

3(20) 7(25)

4(40) 8(40)

Total Total

TOTAL POINTS =

I want my final to count 300 pts.______________________________(This will drop your lowest exam) Signature

R = 8.314 J/mol-K = 96486 C/mol e- C = amp x secΔE = ΔE - 0.0592/n log Q ΔG = ΔG + RT ln Q ΔG= -RTlnKeqΔG = -nΔE PV = nRT R = 0.08205 L-atm/mol-Kln(Ai/Af) = kΔt 1/Af - 1/Ai = kΔt 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 notcold and didn't taste good so I put a couple of ice cubes in it. The ice cubes melted and the watertasted better. If my cup had 250 mL of water at 22C, and I put two ice cubes into the water thatweighed 20 grams each, what was the final temperature of the water if the ice cubes were initiallyat 0C? 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 3minutes to half-cook (soft boil) an egg at 100C. If the activation energy of the cooking of anegg is 75 kJ/mole, how long would it take to soft boil an egg at the top of Mt. Everest wherewater boils at 80C?

4a) Calculate the pH at each of the following points for the titration of 100 mL of 2.00 MPhthalic acid with 5 M NaOH

H2Phthal ---> H+ + HPhthal- pKA1 = 2.89HPhthal- ---> H+ + Phthal2- 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 abuffer of pH = 3.5?


CO(g) + NO2(g) ---> CO2(g) + NO(g)

can be determined by the following data;

Initial Rate/10-4 [NO]/10-3 M [CO]/10-3 M [NO2]/10-3 M

3.60 1.20 1.50 0.807.20 1.20 3.00 0.800.90 1.20 0.75 0.400.90 2.40 1.50 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 circlethe 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 slowO + 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.

OHCH3

H2NC C C

OH OH

C C C

O

C

C C

O C C C C

C C

CCl

CH3

C

Cl

7d) If the Ksp for this cell is 1.8x10-10 then what is the cell voltage when all the ionic species are0.01 M?

8) Please name or draw the following compounds.

butanone 3,3-dichloro butanoic acid

trans 1,3-cyclopentadiol o-aminotoluene

Chemistry 121 NameFinal Exam May 26, 2000


Question Credit Question Credit

1(15) 6(20)

2(15) 7(15)

3(15) 8(30)

4(15) 9(20)

5(25) 10(30 )

Total Total

TOTAL POINTS =

I want my final to count 300 pts._________________________________(This will drop your lowest exam) Signature

R = 8.314 J/mol-K = 96486 C/mol e- C = amp x secΔE = ΔE - 0.0592/n log Q ΔG = ΔG + RT ln Q ΔG= -RTlnKeqΔG = -nΔE PV = nRT R = 0.08205 L-atm/mol-Kln(Ai/Af) = kΔt 1/Af - 1/Ai = kΔt

1) Which of the following compounds is the most soluble in water? Show your work.

BaSO4 Ksp = 1.98x10-10

Fe(OH)2 Ksp = 1.64x10-14

Ag3PO4 Ksp = 1.2x10-15

2) Please write the equilibrium reaction that will occur when the following compounds aremixed. 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+ + Br-

The rate of disappearance of bromine was measured for several different concentrations ofacetone, bromine and hydrogen ions;

Rate [Acetone] [Br2] [H+]6x10-5 0.30 0.050 0.0502.4x10-5 0.30 0.100 0.0501.2x10-4 0.30 0.050 0.1003.2x10-4 0.40 0.050 0.2008x10-5 0.40 0.050 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+ + S2- Keq = 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-



5c) If the = 1.056 volts, what is the voltage of the cell when concentration of the ionic speciesin 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 3minutes to half-cook (soft boil) an egg at 100C. If the activation energy of the cooking of anegg is 75 kJ/mole, how long would it take to soft boil an egg at the top of Mt. Everest wherewater boils at 80C?

6b) If the rate of a reaction doubles with a 15C change in temperature, what is it's activationenergy? State your assumptions.

7) A block of copper intially at 500C was placed on top of a 100 gram block of ice at 0C. Allof the ice melted and turned into water. The water and the block of copper ended up at 20C.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 Phthalicacid with 5 M NaOH

H2Phthal ---> H+ + HPhthal- pKA1 = 2.89HPhthal- ---> 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 -5C to turn it into 60 grams of water and 40grams of steam, both at 100C? 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

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