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EAST%LOS%ANGELES%COLLEGE%%%%CHEMISTRY%DEPARTMENT%

Chemistry%102%%%%%%%%%%%Laboratory%%Manual%

Version%5%"

Edited"8/21/12"

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Table&of&Contents&

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Introduction to the Phytoremediation Semester Project: ...........................................................................3"

Set up of test pots ......................................................................................................................................4"

Spectroscopy: Beer’s Law .........................................................................................................................5"

A Discovery-Based Experiment illustrating how Iron Metal is used to remediate Contaminated Groundwater ........................................................................................................................................... 10"

Chemical Kinetics - the Iodine Clock Reaction...................................................................................... 17"

Is"solubility"a"periodic"property? ............................................................................................................ 24"

An"Environmental"Project"for"First"Year"Chemistry"Students:"Analysis"of"Natural"Buffer"Systems"and"the"

Impact"of"Acid"Rain................................................................................................................................. 30"

Electrochemistry and the Nernst Equation ............................................................................................. 34"

The analytical determination of aqueous [Cu2+] ..................................................................................... 40"

Equilibrium: Measuring the soil/water distribution coefficient: the effect of decreasing pH ................. 41"

Equilibrium: Measuring the soil/water distribution coefficient: the effect of increasing T .................... 45"

Phytoremediation Analysis ..................................................................................................................... 47"

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Introduction to the Phytoremediation Semester Project: Goal: To become familiar with phytoremiadition: What it is, how it works, why it is needed, and what is its advantages to the environment 1. Read the phytoremediation paper by M.M. Lasat (J Environ Qual 31:109-120 (2002)) posted on ACE, then define the following terms:

Phytoremediation Phytoextraction Bioavailability Hyperaccumulator

2. Read section 1 of www.inchem.org/documents/ehc/ehc/ehc200.htm, and then answer the following:

a. Describe the variety of forms that copper takes in the environment. b. Describe 2 methods of copper determination. c. Describe 2 sources of human/environmental exposure to copper.

d. What is the LD50 range among animal species? (LD50 = Lethal dose 50% = the amount of a substance required to kill off half of a population, in mg Cu per kg body weight). Briefly comment on the range; is it expected? Does it seem high, low, or unexpected?

3. Explain why complete removal of Cu from soils would be equally as undesirable as excess Cu.

4. What molarities are 10, 100, and 1000 ppm Cu2+

in water? (ppm = parts per million = mg/kg)?

5. What total mass of CuSO4·5H2O (in mg) do you need to make Cu2+

solutions which are 10, 100, and 1000 ppm?

These Questions will be collected at the end of the lab period.

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Set up of test pots Objectives:

This experiment demonstrates the nature of science by having students experience experimental design, forming research objectives and hypothesis formation and testing. Students will assess the effectiveness of phytoremediation. Students will explore the chemistry and biology of Cu

2+ in water and soil in some

depth. Students will learn to perform careful quantitative analytical procedures. Materials: Soil, pots, water, balances, CuSO4·5H2O (CAS#: 7758-98-7) Precautions: CuSO4 is mildly toxic; wear gloves, eye protection and appropriate apparel

Procedures:

1. Students should be grouped in 4 and each student will be responsible for a plant. 2. Label 4 different pots with (Each student will be responsible for their pot):

No Cu2+ (label ‘control’) 10 ppm Cu2+ (label appropriately) 100 ppm Cu2+ 1000 ppm Cu2+

3. Each student will transfer a starter tomato plant into their labeled pot, and measure the amount (mass) of soil being added to each pot. Make sure to record this data in your laboratory notebook. WITHIN EACH GROUP, BE SURE TO USE THE SAME SIZED POT AND THE SAME AMOUNT OF POTTING SOIL. 4. There will be (1) large carboy each of 10 ppm, 100 ppm, and 1000 ppm Cu2+ in the back of the lab. Use these solutions for regular watering of plants. The control should be deionized water. 5. Each student will “water” their plant with their designated Cu2+ solution WEEKLY, so that each plant within a group is receiving the same amount of liquid. Each student should be maintaining a log in their lab notebooks of the weekly “watering” routine, including date, time, amount and observations. Make sure to communicate with all members of the group, so that each plant is treated equally. 6. Ask your instructor where plants will be located. Plants that are placed outside in the planter circle should be monitored with care…pay attention to extreme hot or cold days. It may be best to place outside plants in partial shade. WHEN WATERING PLANTS, PLEASE MAKE SURE THEY DRAIN COMPLETELY OVER A SINK.

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Spectroscopy: Beer’s Law INTRODUCTION

A useful analytical tool for determining the concentration of colored material in solution is absorbance spectrophotometry. Colored substances absorb light in the visible electromagnetic spectrum. The amount of light absorbed by a substance in solution is easily measured by a photocell detector and is proportional to the solute concentration.

If Io is the intensity of light entering a solution and It is the intensity of light exiting the solution, then the transmittance, T, of the solution is given as It/Io. Transmittance is also expressed as a percentage, (It/Io)×(100%). Because it relates more directly to solute concentration, the absorbance A is more often used to express the amount of light a solution absorbs. A = -log(T) or A = log(Io/It). At a given wavelength of light, absorbance depends on the absorptivity of the solute, the length of solution the light passes through and the solute concentration. This relationship, known as Beer's law, is expressed as A = εbc. Here “ε” is a proportionality constant (molar absorptivity for molar concentration units), “b” is the path length through the solution, and “c” is the solute molar concentration.

In Beer’s Law experiments, it is necessary to find a wavelength at which light is absorbed strongly by the solute being studied, without interference by other absorbing solutes and where the absorbance does not change rapidly with changing wavelength. Then it necessary to measure the absorbance of several solutions of known concentration so a calibration curve can be drawn.

In this experiment a curve of absorbance versus wavelength will be obtained for a cobalt(II) nitrate hexahydrate Co(NO3)2·6H2O solution to give the optimum wavelength for measuring the salt absorbance. At that wavelength, the absorbance values of four solutions of known concentration will be measured to construct a calibration curve from which the molar concentration of an unknown Co(NO3)2·6H2O solution will be determined.

PreLab questions

1. Calculate the mass of cupric sulfate pentahydrate needed to prepare 50.00 mL of 0.0800M CuSO4·5H2O. 2. Describe how to prepare 50.00 mL of 0.0400M CuSO4·5H2O from the above cupric sulfate solution. 3. The following data were obtained for CuSO4·5H2O solutions at a wavelength of 650 nm.

[CuSO4·5H2O] (M) Absorbance 0.000 0.000 0.0100 0.110 0.0200 0.200 0.0400 0.450 0.0800 0.920

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a) Plot a graph of absorbance vs. concentration using the above data. Fit to the best straight line with the linear regression tool on excel. Attach the graph. b) By interpolation, determine the molarity of a copper sulfate pentahydrate solution whose absorbance is 0.170 at 650 nm.

Turn in Pre Lab questions to your instructor before performing experimental procedure.

EXPERIMENTAL

Solution Preparation

To prepare a 0.0500 M solution of the salt, pipet 5.00 mL of the 0.100M Co(NO3)2·6H2O stock solution into another 10 mL volumetric flask, dilute to the mark with distilled water and mix. Pipette 5.00 mL of the 0.0500M solution into a 10 mL volumetric flask and dilute with distilled water to the mark to give 0.0250M. Prepare a 0.0125 M solution by pipetting 5.00 mL of the 0.0250 M solution into a 10 mL volumetric flask and diluting to the mark. This will give a total of four standard solutions for groups of two students to be used for the calibration curve.

Absorbance Spectrum Determination

For this part only, work in groups of two. Record your data directly from the spectrophotometer into your lab notebook. Write your partners’ name in your notebook. Set up Ocean Optics USB4000 Spectrometer as described below.

To take an absorbance measurement using SpectraSuite, follow the steps below:

1. Open SpectraSuite and close any open graphs (i.e. Graph A). 2. Place your reference or blank cuvette in the cell compartment. 3. Access the Absorbance wizard by going to the main menu under File/New/Absorbance Measurement. This will open the Select spectral source wizard. You only have one spectrometer with one channel, it will be pre-selected for you. Click the “Next” button. 4. Set Acquisition Parameters wizard as shown in snapshot below.

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" 5. You MUST select the Strobe/Lamp Enable checkbox before you begin adjusting the integration time. If you forget to check the box, it will be necessary to start the wizard (and possible restart SpectraSuite) before you can continue. This is because SpectraSuite gets stuck on a very long integration time and cannot quit out of it. Usually, if you close the wizard and wait a couple of minutes, the acquisition (seen in the lower right corner) will stop. Then, you can start a new Absorbance wizard. 6. Set Acquisition Parameters wizard. Click the “Set Automatically” button or adjust the integration time manually to get the spectrum peak close to the “Recommended peak value.” 7. Set Scans to Average and Boxcar Width. These settings are optional. Scans to Average increases the signal to noise ratio. Boxcar Width smoothes the curve, but lowers optical resolution. To start, we can try 10 Scans to Average and a Boxcar Width of 2. Click the “Next” button. 8. Store Reference Spectrum wizard. Click the yellow light bulb button to store the reference spectrum, as illustrated in snapshot below.

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8. Store Dark Spectrum wizard. Uncheck the “Strobe/Lamp Enable” checkbox and click the dark light bulb button to store the dark spectrum. The Strobe/Lamp enable checkbox does not turn off the source, but only closes the shutter; therefore, you need not worry about having to warm up the source again. Check the “Strobe/Lamp Enable” button again, and then click the “Finish” button as shown in snapshot.

9. Absorbance Mode. You are now in Absorbance mode. With your reference or blank still in the cuvette holder, you should be reading absorbance of 0 OD. (Note to look at specific wavelengths, left click mouse in spectrum region to get the cursor.)

10. Fill another cuvette with 0.100 M cobalt(II) nitrate solution, put it into the cell compartment and measure its absorbance spectrum. From this data determine the optimum absorbance wavelength for the salt by moving the vertical cursor. Left-click on the spectrum for the vertical cursor to appear.

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Calibration Curve

This can be done with a partner.

Now measure the absorbance of the 0.100 M, 0.0500M, 0.0250 M, and 0.0125 M standard cobalt(II) solutions at the optimum wavelength you determined previously by placing cursor on that wavelength and record the values in your lab notebook.

Unknown Solution

Each student is expected to have their own unknown.

Without changing any of the settings, immediately measure the absorbance of the unknown solution issued to you. Don’t forget to record the unknown number.

Standard and Unknowns solution absorbances must be measured consecutively or else there will instrument error in your measurements.

Analysis

1. Calibration Curve- Using Excel plot a graph of absorbance vs. concentration for the standard cobalt solutions. Fit the best straight line to the data and obtain the computer-generated value for m, the slope of the straight line. The y-intercept, b, should be close to zero. Attach the graph to this report.

2. Unknown Solution Concentration: From the calibration curve, determine the molarity of your unknown solution, not by interpolation on the graph but algebraically from the computer-generated m value in the equation for a straight line, y = mx + b. If b equals zero, then y = mx, where m is the slope of the line, x is the solution concentraton and y is the measured absorbance. Since the molar absorptivity ε and the path length are constants, this equation is consistent with Beer’s law A = εbc which becomes A = mc.

3. Conclusion: State whether or not you believe Beer’s law was validated by this experiment. Give a full explanation for your conclusion.

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A Discovery-Based Experiment illustrating how Iron Metal is used to remediate Contaminated Groundwater Barbara A. Balko

1* and Paul G. Tratnyek

2

1Department of Chemistry, Lewis & Clark College 0615 SW Palatine Hill Road Portland, OR 97219

2Department of Environmental Science and Engineering Oregon Graduate Institute of Science & Technology P.O. Box

91000, Portland, OR 97291-1000 *Corresponding Author, email: balko@lclark.edu 1. Instructions for Students 1.1 Overview In this experiment, you will investigate the chemistry behind iron permeable reactive barriers (iron PRBs), a new technology that is widely used to remediate contaminated groundwater. Contaminant remediation involving iron PRBs is a redox process: the iron metal undergoes oxidative dissolution while the contaminant is reduced. The reaction is complicated, however, by the fact that it involves a surface that changes due to the development of a layer of rust (iron oxide) on the iron. You will examine the iron PRB-contaminant reaction by measuring the pseudo first-order rate constant for the degradation of a dye (the model contaminant) in the presence of granular iron under various experimental conditions to investigate some aspect of the degradation reaction that is of interest to your group. 1.2. Background 1.2.1. Use of Permeable Reactive Barriers It has long been known that the oxidation (i.e., corrosion) of metals such as iron, tin, and zinc can bring about the reduction of halogenated organics (1). In the late 1980’s, researchers at the University of Waterloo rediscovered these redox reactions while investigating groundwater contaminated with halogenated solvents. The researchers recognized that these reducing metals could be used for remediating contaminated groundwater by constructing permeable reactive barriers (PRBs) composed of one of the reducing metals (2-4). A PRB is a zone comprised of granular metal that extends below the water table and intercepts the flow of contaminated groundwater; as the contaminants pass through the PRB, they are reduced to non-toxic compounds and, ideally, the groundwater that emerges is free of hazardous substances. Figure 1 shows a schematic of a PRB. Iron is the current metal of choice for PRBs because it is readily available, inexpensive, nontoxic, and a good reducing agent (3,5).

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Figure 1: Water infiltrating soil contaminated with chemical waste forms a plume of contaminated groundwater. This plume travels in the direction of groundwater (GW) flow and spreads through the water table if left untreated. When contaminated groundwater flows through an iron permeable reactive barrier, the contaminants are reduced by the oxidation of the iron. The plume can then be considered treated. (From http://www.powellassociates.com/sciserv/3dflow.html with permission.) Iron PRBs, colloquially known as iron walls, have many advantages over more traditional groundwater remediation technologies. First, iron walls are able to remediate many contaminants in addition to halogenated solvents, such as pesticides (6-8), munitions (9), nitrate (10,11), and heavy metals (12,13) (e.g., chromium (14) and uranium (15,16)). Second, iron walls are a “passive way” to remove contaminants from groundwater in that they require little maintenance after installation (4). Because of these advantages, as well as the prevalence of contaminants that iron walls are capable of removing, iron PRBs are now widely used in North America and Europe to clean up contaminated sites. 1.2.2. Chemistry of Permeable Reactive Barriers The degradation of contaminants by iron metal, Fe

0, can be explained in terms of textbook redox

chemistry. Consider the degradation of carbon tetrachloride, CCl4, by Fe0. The anodic half-reaction is

the oxidative dissolution of Fe0 (Equation 1) and the cathodic half-reaction of interest, assuming a

hydrogenolysis reaction mechanism (17, 18), is the reduction of CCl4 to CHCl3 (Equation 2). The Eo for

the next cell reaction at pH 7 is 1.11 V. Fe

0 → Fe

2+ + 2e

- E° = 0.44 V (19) (1)

CCl4 + 2e-+ H

+ → CHCl3 + Cl

- E° = 0.67 V (20) (2)

CCl4 + H+ + Fe

0 → CHCl3 + Cl

-+ Fe

2+ E° = 1.11 V (3)

It should be noted that the product of the degradation reaction, CHCl3, is a hazardous and regulated substance. While CHCl3 is not readily degraded by Fe

0, it is much more biodegradable than CCl4 so

groundwater contaminated with CCl4 should be cleaned by an iron PRB. It is possible; however, that iron is not directly responsible for contaminant reduction by iron walls (17). Fe

2+ and/or H2 will be

present in the system due to the reduction of dissolved oxygen and water by Fe0:

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2Fe

0 + O2 + H2O → 2Fe

2+ + 4OH

- (4) Fe

0 + 2H2O → Fe

2+ + H2 + 2OH

- (5) Thermodynamically, both Fe

2+ and H2 are capable of reducing some contaminants. The degradation of a

typical halogenated organic contaminant, RX, by Fe0 typically obeys the following rate law (21):

Rate = -d[RX]/dt = -k[Fe active sites][RX] (6). This equation can be simplified into a first-order kinetic equation, since in most experimental systems it has been found that the concentration of iron active sites does not vary significantly during the course of the degradation reaction.

-d[RX]/dt = -kobs[RX] (6)

Thus, the kinetics of degradation of groundwater contaminants by iron metal generally is pseudo first-order. It should be emphasized that kobs is not a true first-order rate constant and will depend on the concentration of active sites on the iron metal. Thus, kobs should be proportional to the surface area of the iron particles as well as the mass of the iron present. The following relationship exists between kobs, the specific surface area of the iron, as, the mass concentration of iron, ρm, and the specific reaction rate constant, kSA (21).

kobs = kSAasρm (7)

Pre-Lab Assignment

1. Consider"the"following"data"for""2ClO"(g)""""Cl2O2"(g),"which"is"second"order"with"respect"to"the"

concentration"of"ClO."

Time"(s)"" [ClO]"(M)"

0" " 2.60*1011"

1" " 1.08*1011"2" " 6.83*1010"3" " 4.99*1010"4" " 3.93*1010"

"" Solve"for"the"rate"constant."

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2. For"the"reaction""""2"HI(g)""H2(g)"+"I2(g)"at"700"K","the"following"data"was"collected:"

Time(s)"" [HI]"(M)"0 10"1000" " 4.4"2000" " 2.8"3000" " 2.1"4000" " 1.6"5000" " 1.3"

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A"plot"of"ln"[HI]"as"a"function"of"time"yields"a"straight"line"with"a"slope"equal"to"^0.5127."a) Write"the"complete"rate"law"for"the"reaction."b) How"much"time"is"required"for"HI"to"decrease"to"25%"of"its"initial"concentration?"

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1.3. Lab 1.3.1. Summary In this lab, you simulate the remediation of contaminated groundwater with an iron wall by mixing iron particles in an indigo carmine dye solution. The rate of degradation of the dye is measured by following the change in the solution absorbance; assuming Beer’s Law is followed, the solution absorbance will be proportional to the dye concentration. The slope of a plot of ln(At/A0) versus time (where At is the dye absorbance at time t and A0 is the dye absorbance at the start of the experiment) is –kobs. Initially, each group of 2 students measures the rate of degradation under the same conditions. To give you a sense of what scientific research involves, detailed laboratory instructions are kept to a minimum. For example, you are instructed to measure the absorbance of the dye but not what wavelength to monitor. It is up to your group to determine the best wavelength to use. Next, each group designs their own experiment to learn more of the details of contaminant degradation using iron PRBs. Once your group has thought of an experiment, be sure to discuss your plans with your instructor before proceeding.

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1.3.2. Experimental Procedures Part 1- Determine the wavelength of maximum absorption (λmax) of indigo carmine dye. (Hint: The visible spectrum goes from 400-700 nm.) Part 2- Make measurement of Absorbance at various indigo carmine dye concentrations in order to construct a Beer’s Law plot to confirm that the dye concentration is proportional to the absorbance in the concentration range used. You will be given a 20 ppm stock solution that you should use to construct a Beer’s law plot. You need to dilute the solution enough such that the absorbance reading is not saturated. IF THE SPECTRUM IS SATURATED, DILUTE DOWN TO A LOWER CONCENTRATION. You will perform serial dilutions to make the other needed solutions for the calibration curve. Make sure to use the appropriate volumetric pipettes and volumetric flasks. Part 3- (1) Put 0.25 g of coarse iron into a UV-Vis polystyrene cuvette (2) Prepare 10 ppm indigo carmine dye solution. Fill a cuvette completely with the 10 ppm indigo carmine dye solution. It is important to minimize trapped air in the cuvette because the oxygen will oxidize the reduced form of the dye. (3) Cap the cuvette to minimize the amount of oxygen that enters during the experiment. (4) Record the initial absorbance of the 10 ppm indigo carmine dye. (5) Set the cuvette up on the rotator or begin to shake the cuvette by hand (if you use the manual method, think of some way to keep the shaking rate as constant as possible). Be sure to record the rotation or shaking rate. (6) Record the dye absorbance as a function of time. Monitor dye concentration at 1 minute intervals (be sure to take an initial reading) for a 10 – 15 minute period. Be sure to shake the cuvette between intervals. Record at least 6 time points; you should record over at least the length of time required for the dye concentration to halve (i.e., the reaction half-life). Initially, you may want to run through the experiment quickly just to get a sense for how long it takes to degrade the dye and then repeat the experiment more carefully. Part 4 Run two more tests in which you’ve changed a single experimental variable. You will have available:

1. Iron particles washed in 0.1M HCl 2. Iron particles washed in 5% H2O2 solution 3. A more concentrated dye solution-try 20 ppm

1.3.3. Controls to Consider There are certain experimental problems that must be considered. Plan your experiments so that these do not affect your experiments or your conclusions. Some possible problems are as follows—the list is by no means complete: (1) Degradation of the dye in room light (indigo carmine is “photosensitive”) (2) Degradation of the dye during UV-Vis analysis (3) Irreversible adsorption of the dye on the iron particles and cuvettes (4) Oxidation of the reduced form of indigo carmine by oxygen (5) Scattering of the spectrophotometer analysis beam by particulate matter in the iron

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1.3.4. Data Analysis: Determination of the Degradation Rate Constant, kobs (1) Assuming that the dye concentration is proportional to the absorbance, plot ln(At/A0) versus time, where At is the dye absorbance at time t and A0 is the initial dye absorbance. (2) The slope of the line will be -kobs if the degradation follows first-order kinetics. Estimate the uncertainty in the value of the rate constant. Using a spreadsheet fit the data, the standard deviation of the slope is one measure of uncertainty. (3) It is possible that a plot of ln(At/A0) versus time will not be linear. Does the data suggest zero or second order kinetics rather than first (i.e., is a plot of (At/A0) or (A0/At) versus time linear)? Can you justify omitting any of the data points? 1.3.5. Lab Report In addition to your abstract, introduction, procedure, and data, be sure to address the following: Results/Analysis:

(a) Beer’s Law plot: Does your data show that dye absorbance is proportional to concentration? (b) Determination of Rate Constant for each Experiment:

(1) Plot of ln(At/A0) versus time: Does your data show that the degradation follows first order kinetics? (2) Calculation of rate constants: Be sure to include an estimate of the uncertainty in these values.

(c) Sample Calculations (d) Summary of results in table form and well-labeled graphs

Discussion: (a) What were the effects of the variables on the rate constant? (b) Why is it important to prevent the re-oxidation of the dye? In what direction will your value of k be shifted if the dye is re-oxidized? (c) What were some of the difficulties you encountered? How did you try to overcome them or what would you change to avoid them if you were to repeat the experiment? (d) What do your results tell you about the contaminant degradation using iron PRBs? Discuss how your results can be used to improve the efficiency of this remediation technique.

Conclusions (a) Summarize the results. What did you learn about the degradation using iron PRBs? How can the results be used to improve the efficiency of iron walls? (b) Research the “pump and treat” method of groundwater remediation and compare this technique with the iron PRB method.

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1.4. Literature Cited 1. Hudlický, M. Reductions in Organic Chemistry; Halsted: New York, 1984. 2. Gillham, R. W.; O'Hannesin, S. F. Ground Water 1994, 32, 958-67. 3. Tratnyek, P. G. Chem. Ind. (London) 1996, 499-503. 4. Powell, R. M.; Powell, P. D. Iron metal for subsurface remediation. In Encyclopedia of Environmental Analysis and Remediation; Myers, R. A., Ed.; Wiley: New York, 1998; Vol. 8; pp 4729-4761. 5. Scherer, M. M.; Richter, S.; Valentine, R. L.; Alvarez, P. J. J. Critical Reviews in Environmental Science and Technology 2000, 30, 363-411. 6. Eykholt, G. R.; Davenport, D. T. Environ. Sci. Technol. 1998, 32, 1482-1487. 7. Sayles, G. D.; You, G.; Wang, M.; Kupferle, M. J. Environ. Sci. Technol. 1997, 31, 34483454. 8. Singh, J.; Shea, P. J.; Hundal, L. S.; Comfort, S. D.; Zhang, T. D.; Hage, D. S. Weed Science 1998, 46, 381-388. 9. Hundal, L. S.; Singh, J.; Bier, E. L.; Shea, P. J.; Comfort, S. D.; Powers, W. L. Environ. Pol. 1997, 97, 55-64. 10. Huang, C.-P.; Wang, H.-W.; Chiu, P.-C. Water Research 1998, 32, 2257-2264. 11. Cheng, I. F.; Muftikian, R.; Fernando, Q.; Korte, N. Chemosphere 1997, 35, 2689-2695. 12. Cantrell, K. J.; Kaplan, D. I.; Wietsma, T. W. J. Haz. Mat. 1995, 42, 201-212. 13. Blowes, D. W.; Ptacek, C. J.; Benner, S. G.; McRae, C. W. T.; Bennett, T. A.; Puls, R. W. Journal of Contaminant Hydrology 2000, 45, 123-137. 14. Blowes, D. W.; Ptacek, C. J.; Jambor, J. L. Environ. Sci. Technol. 1997, 31, 3348-3357. 15. Gu, B.; Liang, L.; Dickey, M. J.; Yin, X.; Dai, S. Environ. Sci. Technol. 1998, 32, 33663373. 16. Fiedor, J. N.; Bostick, W. D.; Jarabek, R. J.; Farrell, J. Environ. Sci. Technol. 1998, 32, 1466-1473. 17. Matheson, L. J.; Tratnyek, P. G. Environ. Sci. Technol. 1994, 28, 2045-2053. 18. Criddle, C. S.; McCarty, P. L. Environ. Sci. Technol. 1991, 25, 973-978. 19. Bratsch, S. G. J. Phys. Chem. Ref. Data 1989, 18, 1-21. 20. Vogel, T. M.; Criddle, C. S.; McCarty, P. L. Environ. Sci. Technol. 1987, 21, 722-736. 21. Johnson, T. L.; Scherer, M. M.; Tratnyek, P. G. Environ. Sci. Technol. 1996, 30, 2634-2640.

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Chemical Kinetics - the Iodine Clock Reaction Purpose To measure the rate of a reaction and determine the rate law, rate constant, and Ea Equipment Needed timer, hot plate, thermometer Chemicals Needed 0.050 M KI, 0.050 M Na2S2O3, starch solution, pH 4.7 buffer, 0.30 M HC2H3O2, 0.80 M H2O2 Discussion In this experiment, you will investigate the rate of the reaction -

3 I-(aq) + H2O2 (aq) + 2 H

+(aq) I3

- (aq) + 2 H2O (l)

The rate law for this reaction is given by

where the exponents x, y, and z and the rate constant k are to be determined. The reaction rate will be monitored by a clock reaction. A fixed amount of thiosulfate ion will be added to each reaction. Thiosulfate reacts rapidly with I3

-

2 S2O3-2 (aq) + I3

- (aq) 3 I-(aq) + S4O6

-2 (aq)

When all the thiosulfate is consumed, the I3 formed is free to react with added starch indicator and the solution will turn dark blue. The reaction rate

is determined by the stoichiometric relationship between S2O3-2 , I3

- , and I-

By measuring the rate for a series of reactions in which only one reactant concentration at a time is varied, the order of the reaction for each reactant can be determined. Consider the reaction

2 NO + 2 H2 N2 + 2 H2O

"

In"experiments"1"and"2,"doubling"[H2]"results"in"a"doubling"of"the"reaction"rate;"therefore"the"reaction"must"be"first"order"in"hydrogen."In"experiments"1"and"3,"doubling"[NO]"results"in"a"quadrupling"

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(double^squared)"of"the"reaction"rate;"therefore"the"reaction"is"second"order"in"NO."Now"that"the"rate"

law"is"known"to"be"rate"="k"[NO]"2"[H2]""

the value of k can be determined by substituting known values for [NO], [H2] and rate into the rate equation.

Experiment [NO] (M) [H2] (M) Initial Rate (M/s)

1 0.10 0.10 1.23 x 10-3

2 0.10 0.20 2.46 x 10-3

3 0.20 0.10 4.92 x 10-3

By measuring the reaction rate at varying temperature, we can determine the activation energy for the reaction. According to collision theory, the rate of a reaction is controlled by the rate at which reactants collide to form the activated complex, a temporary species formed by collision prior to formation of products. The activation energy (Ea) is the minimum amount of energy required to form the activated complex.""

A + B activated complex C + D

The rate of formation of activated complex controls the rate of the reaction.

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The Arrhenius Equation relates the reaction rate constant to the activation energy and the temperature.

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so a plot of ln k versus 1/T gives a straight line with a slope = -Ea/R and an intercept ln A. If k is known at two different temperatures, the following equation applies.

Pre-Lab Assignment:

1. The"chemistry"of"smog"formation"includes"NO3"as"an"intermediate"in"several"reactions."

a. If"Δ[NO3]/Δt"is"^2.2*105"M/s"in"the"following"reaction,"what"is"the"rate"of"formation"of"NO2?"

NO3(g)""+"NO(g)""""2NO2(g)"

"

b. What"is"the"rate"of"change"of"[NO2]"in"the"following"reaction"if"Δ[NO3]/Δt"is"^2.3"M/s?"

2NO3(g)""""2NO2(g)""+""O2(g)"

"

"

"

Activated Complex Activated Complex

E

C + D

exothermic reaction endothermic reaction

where Ea = activation energy in kJ/mol -Ea/RT k = Ae R = gas law constant, 8.314 J/K mol T = temperature in K A = frequency factor

"C"+"D"

%"

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"

2. Predict"the"experimental"rate"orders"for"each"reactant"given"the"following"information:"

NO(g)""+""Br2(g)""""NOBr2(g)"

" It"is"observed"that:"

" " The"rate"doubles"when"[NO]"is"doubled"and"[Br2]"remains"constant."

" " The"rate"is"cut"by"one^quarter"when"[NO]"remains"constant"and"[Br2]"is"reduced"by"one^half."

" Based"on"your"rate"orders,"what"is"the"molecularity"of"the"overall"reaction?"

"

3. For"the"reaction""""NO2(g)""+"O3(g)""NO3(g)"+"O2(g)","the"following"data"was"collected:"

""""""""Experiment"#" """"""[NO2]o(M)" """""""[O3]o(M)" """Initial"rate"(M/s)""""""""""""""""""""""""""""""""""""""""""""""""1" " """"0.21" " """0.70" " """"6.3"

2" " 0.21" " 1.39" " 12.5"3" " 0.38" " 0.70" " 11.4"

"

Write"the"complete"rate"law"for"the"reaction."Show"all"of"your"work."

Procedure

A. Determination of the rate law

Obtain 4 clean 150 mL beakers. Label them 1 through 4. To each beaker, add the reactants as specified in Table 1. Measure the temperatures of each solution and record these on the data sheet. If the temperatures are not within ± 0.5° C of each other, place the beakers close to each other and gently stir with magnetic stir bar until the temperatures are within ± 0.5° C. Make sure to use the same stir rate for each solution. Use a pH meter to measure the pH of each solution and record this data in your notebook.

In a separate beaker, measure 2.0 mL of 0.80 M H2O2 solution for reactions 1, 2, and 4, and 4.0 mL for reaction 3.

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Table 1: Initial volumes of reactants

In 150 mL Beaker In Separate Beaker

exp volume of H2O (mL)

volume of 0.050 M KI (mL)

volume of 0.050 M Na2S2O3 (mL)

volume of starch (mL)

volume of pH 4.7 buffer (mL)

volume of 0.30 M HC2H3O2 (mL)

volume of 0.80 M H2O2 (mL)

1 25 5.0 1.0 1.0 6.0 0 2.0

2 20 10.0 1.0 1.0 6.0 0 2.0

3 23 5.0 1.0 1.0 6.0 0 4.0

4 20 5.0 1.0 1.0 6.0 5.0 2.0

Perform each reaction separately. Use a timer. Quickly add the H2O2 solution to the beaker with stir bar and note the time required for the appearance of the intense blue color. Record this time (in seconds) in your lab notebook.

B. Determination of the activation energy

Obtain a clean 150 mL beaker and label it 5. Add the reactants as specified in Table 2. Use a pH meter to measure the pH and record this data in lab notebook.

Table 2: Initial volumes of reactants

In 150 mL Beaker In Separate Beaker

exp volume of H2O (mL)

volume of 0.050 M KI (mL)

volume of 0.050 M Na2S2O3 (mL)

volume of starch (mL)

volume of pH 4.7 buffer (mL)

volume of 0.30 M HC2H3O2 (mL)

volume of 0.80 M H2O2 (mL)

5 25 .0 5.0 1.0 1.0 6.0 0 2.0

"

"

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"

Prepare a hot water bath by heating about 250 mL of water in a 400 mL beaker to a temperature of 40 -45° C. Remove the water bath from the hot plate and put beaker 5 in the water bath. When the temperature of solution 5 is the same as the water bath temperature, record the temperature in your lab notebook and add the 2.0 mL of H2O2 solution. Record the time (in seconds) required for the appearance of the intense blue color in your lab notebook.

Calculations

A. Determination of the reactant concentrations

Use the dilution formula M1V1 = M2V2 to determine the initial concentration of each reactant after being diluted from the original concentration to the new final volume of 40 mL. Show this calculation in your sample calculation section and record calculated result in a results table.

[H+] is determined from the pH value. [H

+] = 10

-pH

Show this calculation in your sample calculation section and record calculated result in a results table.

B. Determination of the rate of the reaction

The rate of the reaction is determined by the time required for the formation of the blue starch complex.

This is set by the amount of S2O3

2-present

Calculate the rate for each reaction and record these values in the results table. Remember to show sample calculations in report. C. Determination of the order of the reaction

For a reaction with a rate law in the form rate = k[A]x [B]

y[C]

z, the order of the reaction for each

reactant can be found by comparing two experiments in which the concentrations of only one reactant (say, A) changes. Then, "

"

"

From this, x can be determined by taking the log of both sides;

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Calculate the reaction order of I-, H2O2, and H+ from reactions performed.

Calculate and record the order of the reaction for each reactant to 2 significant figures, then rounded to the nearest integer value. The overall order of the reaction is the sum of the orders in each reactant.

D. Determination of the rate constant

Once the orders of the reaction x, y, and z are determined, the experimental rate constant for each run is calculated by substituting known values for rate and concentration into the rate expression

Calculate the experimental rate constant for each reaction and record these values in the results table. Calculate the average experimental rate constant for reactions 1 – 4 and the standard deviation. Make sure to include sample calculation.

E. Determination of the activation energy

The activation energy is calculated from the Arrhenius equation once values of k are known at two different temperatures:

where R = gas law constant, 8.314 J/mol K

T = temperature in K (= °C + 273.15)

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Is&solubility&a&periodic&property?&

Lab Documentation Supplementary Material for Online Publication, Journal of Chemical Education K.L. Cacciatore, J. Amado, J. J. Evans, H. Sevian

Connecting solubility, equilibrium, and periodicity in a green inquiry experiment for the general chemistry laboratory

In this experiment you will try to determine how soluble three ionic solids are in water, the hydroxides of magnesium, calcium, and strontium ions. We will know discuss one of these solids, magnesium hydroxide, Mg(OH)2, in detail. The cation in magnesium hydroxide is magnesium, Mg2+, and the anion is hydroxide, OH-. Symbolically we can represent what happens when magnesium hydroxide dissolves by writing the reaction below:

Mg(OH)2 (s) Mg2+ (aq) + 2OH- (aq)

The double arrow between the right and left sides of the reaction indicates that the process is

reversible. In other words, the solid dissolves into aqueous ions (forward reaction) and at the same time the ions form the solid and come out of solution (reverse reaction). If a system is originally set up at non-equilibrium conditions, the forward and reverse reactions will proceed at different rates (speeds) until equilibrium is established. At that point, the concentrations of the chemicals involved will no longer change, but both the forward and reverse reactions will continue to happen at equal rates.

If the system has been set up in such a way that the maximum amount of the compound could dissolve (you would know this if there is some excess solid at the bottom), then eventually it will reach an equilibrium point where the maximum amount of the compound is dissolved in the water. At that point, the concentrations of the ions coming from the compound will no longer change, even though the dissolving recrystallization equilibrium is still occurring. Once equilibrium has been established, and if the maximum amount of the compound is dissolved, then the equilibrium constant at that temperature can be calculated from the measurements of the ion concentrations.

The concentrations of magnesium ions and hydroxide ions in the saturated solution can be determined by any technique which allows one to measure concentration of an ion in solution. Some techniques are direct measurements and other techniques involve indirect measurements which are then mathematically converted into the ion concentration. One technique for measuring ion concentrations is titration. This technique is explained below. The experimental results can then be used to figure out the solubility – the maximum amount of solid that will dissolve in a given amount of water -- of magnesium hydroxide. One way to express the solubility of any compound is in terms of the equilibrium constant for the dissolving process. This equilibrium constant is also sometimes called the solubility product constant, and it has the symbol Ksp. For the magnesium hydroxide dissolving equilibrium,

Mg(OH)2 (s) Mg2+ (aq) + 2OH- (aq),

the Ksp expression is Ksp = [Mg2+][OH-]2.

Recall that the square brackets refer to the molar concentration of the ion in solution, so, for example, [Mg2+] is the molarity of magnesium ions. Also, note that the hydroxide ion concentration is squared because the stoichiometric coefficient of hydroxide ion in the reaction is two, and the magnesium hydroxide does not appear in the Ksp expression because it is a solid.

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Calculations with Ksp and ion concentrations

Example #1: Finding Ksp from ion concentration The concentration of silver in a saturated solution of silver sulfate is 0.0288M. What is the Ksp of silver sulfate? Solution The dissolving reaction is shown below:

Ag2SO4 (s) 2 Ag+ (aq) + SO42- (aq)

The ICE table (“ICE” stands for Initial, Change, Equilibrium) below indicates the process used to solve the problem. The concentration of silver is given information (A. in the table). Based on the stoichiometry of the dissolving equilibrium, there must be twice as many silver ions as sulfate ions (B.), and therefore the concentration of sulfate is half the concentration of silver (C.).

Ag2SO4 (s) 2 Ag+ (aq) + SO42- (aq)

Initial

Change B. 2x B. x

Equilibrium A. 0.0288M C. (0.0288M/2)=0.0144M

Now that the concentrations of both ions are known, the Ksp expression can be written and the ion concentrations used to calculate the value of Ksp. Ksp = [ Ag+]2[SO4

2-] = (0.0288)2(0.0144) = 1.19 × 10-5

NOTE: This calculation is only valid if silver and sulfate are the only ions present in the solution in significant amounts.

Acid-Base Titration

Solubility and Ksp can be studied experimentally through the technique of titration. In a titration, a sample of the unknown solution (analyte) is analyzed by slowly adding measurable amounts of a known solution (titrant). The analyte and titrant react with each other as the titrant is added. The titration ends when exactly the right amount of titrant has been added to react with all of the analyte solution. This is called the equivalence point of the titration, because at this point the number of moles of titrant added is exactly the same as the number of moles of analyte in the sample. At this point the amount of titrant added can be measured and used to figure out the concentration of the analyte, which is the goal of the titration.

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One type of titration is an acid-base titration. In an acid-base titration the analyte is either an

acid or base. If the analyte is an acid then the titrant is a base. Conversely, if the analyte is a base then the titrant must be an acid. An acid-base titration can be used to determine the Ksp of magnesium hydroxide. Magnesium hydroxide is a base, because it yields hydroxide (OH-) ions in solution. Because Mg(OH)2 is a base, a saturated solution of magnesium hydroxide can be titrated with an acid such as hydrochloric acid. The acid-base reaction between the titrant HCl and the analyte Mg(OH)2 is:

2HCl (aq) + Mg(OH)2 (aq) 2H2O (l)+ MgCl2 (aq).

The net ionic equation for this acid-base reaction is:

H+ (aq) + OH- (aq) H2O (l)

The indicator phenolphthalein is added to the magnesium hydroxide solution before the titration begins, which makes the solution bright pink. Phenolphthalein is used as the indicator in this titration because it changes color near pH 7. Phenolphthalein is pink in basic solution and colorless in acidic solution. The magnesium hydroxide solution changes from bright pink to colorless as the titration happens. This color change occurs because the hydrochloric acid titrant that is being added is reacting with the magnesium hydroxide and so the solution is becoming less basic. When the solution is less basic, the phenolphthalein is less pink. The solution becomes completely colorless right when all of the magnesium hydroxide has reacted with hydrochloric acid. When the solution is colorless the titration is over. The amount of hydrochloric acid added is then measured on the buret. This data is then used to calculate the concentration of the hydroxide ions in the original magnesium hydroxide solution. Once the concentration of hydroxide is known, the Ksp of magnesium hydroxide can be found.

Left:&&The&titration&apparatus&includes&a&buret&filled&with&titrant&solution&clamped&over&a&flask&containing&a&few&drops&of&indicator&and&a&

precisely&measured&amount&of&the&analyte&solution.&Note&that&the&titrant&is&colored&yellow&

in&this&illustration&in&order&to&make&it&easier&to&see&–&the&actual&titrant&is&often&colorless).&&

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Periodicity and Periodic Trends When the elements are arranged in order of increasing atomic number, many of their physical

and chemical properties, including ionization energy, reactivity, and atomic size, display periodic trends. This phenomenon, which is often called simply periodicity, was long ago used to organize the elements into the periodic table you are familiar with. The table is arranged the way it is so that the elements that are most similar to each other are in the same column or group of the table.

If one examines the properties of the elements in a particular group, many trends can be observed, particularly in the case of the representative elements, which are in Groups 1, 2, and 13-18. Many of these within-group trends can be understood in terms of Coulomb’s Law, a basic principle of physics which states that the electrical force between two charged particles is inversely related to the square of the distance between the two particles and directly related to the magnitudes of the charges on those particles. In terms of an atom, the two charges being considered can be thought of as the negative charge on the outermost electron in the atom, whose distance from the nucleus is the radius of the atom, and the positive “effective nuclear charge” (Zeff), whose location is the nucleus. The size of Zeff is approximately determined by the number of protons in the atom and the number of inner-shell electrons in the atom, which shield the outermost electron from some of the electrical attraction from the protons. As one moves down a group in the periodic table, Zeff remains about the same and the atoms’ radii increase, so elements at the bottom of the group have atoms containing electrons which are much further away from the nucleus than are the outer electrons found in atoms of elements at the top of the group. Because of this difference in size, it makes sense that the outer electrons in the elements at the bottom of the group are more weakly attracted to their nuclei than are the outer electrons in the elements at the top of the group. This difference in attractions underlies the trends that we observe within a group.

The trend in covalent atomic radii, a measure of atomic size, of the Group 1 elements presents a good example of periodicity. Li has a radius of 123 pm (picometers, 10-12 m), Na of 154pm, K of

Above:&The&volume&of&titrant&in&the&buret&is&recorded&before&the&titration&and&at&the&endpoint,&and&the&total&volume&of&titrant&used&is&determined&by&subtraction.&&Ksp&can&then&be&determined&from&this&titration&data.&

&

&

&

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203pm, Rb of 216pm and Cs of 235pm.1 These data show that as one moves from the first element in the group, Li, to the fifth element in the group, Cs, the size of the atoms gradually increases. In general from top to bottom in a group the atomic size gradually increases. Other trends in measurable physical properties of elements include ionization energy and electron affinity, which both gradually decrease as one looks from the top to the bottom of any group.

Within-group periodic trends are also observed in the compounds formed by elements, also as a result of the Coloumb’s Law relationship. The chloride compounds of Group 2 metals, BeCl2, MgCl2, CaCl2, SrCl2. and BaCl2, present a good example. The melting points of these chloride compounds gradually increase from top to bottom in Group 2, from 415°C for BeCl2 to 962°C for BaCl2.2 The solubility of similar compounds is another property that often shows a trend within a group.

Green Chemistry Green chemistry is a philosophy of how to do chemistry so as to minimize harm to chemists, the larger human population, and the environment. The philosophy of green chemistry contends that the important goals of chemistry like producing medicines, finding ways to make alternative sources of energy usable, and educating future scientists can still be accomplished using materials and methods that are less dangerous or wasteful than what may have been used in the past. For example, this experiment on the solubility of the Group 2 metals was designed using principles of green chemistry, so that it is less potentially harmful than another similar experiment. At its most basic level green chemistry can be thought of as the application of the well-known slogan, “Reduce, Reuse, Recycle” to the world of chemistry. Pre-lab questions

1. Write the balanced dissolving equilibrium reactions for the following compounds. Be sure to indicate what phases they are in (i.e., aqueous, solid).

a. Ca(OH)2 b. CaSO4 c. Al2(SO4)3

2. Write the Ksp equilibrium constant expressions for each of the compounds in #1 above.

3. Suppose that you are going to do a titration experiment to determine the solubility of calcium

hydroxide using a hydrochloric acid solution. a. What chemical reaction is occurring during the titration? b. What is the analyte solution in the titration? c. What is the titrant solution in the titration? d. What else do you need to add to the analyte before you begin the titration? What

purpose does it serve? e. How will you know when the titration is over?

"""""""""""""""""""""""""""""""""""""""""""""""""""""""""""""1 Miessler, G. and Tarr, D., “Inorganic Chemistry,” 3rd ed., Prentice Hall, Saddle River, NJ 2004 p.45 2 CRC Handbook of Chemistry and Physics, 85th edition. David R. Lide, Editor. CRC Press, 2004.

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Is solubility a periodic property? Overview

Below is a description of the laboratory task you will do during today’s laboratory session, and guidelines and questions for the report you will turn in at the end of the laboratory. You will also need the information contained in your pre-lab handout to complete today’s experiment. Laboratory Task

During today’s laboratory period you will attempt to replicate the work of a student who did this experiment in this class last year. That student’s lab report is being provided instead of a standard laboratory procedure for you to use as you plan and carry out your experiment. Other students in your class will receive the lab reports of different students, and thus they may have different information than you do. In order to have a better understanding of the experiment, you and your classmates may discuss information with each other during the lab period, but you may not read each other’s lab reports.

Post-lab Questions

1. What sections of the student lab report you received were done well? What made those sections good (give detailed reasons!).

2. What sections of the student lab report you received were done poorly? How could those sections have been improved (give detailed suggestions!)?

3. Are your values of Ksp about the same as those found in the lab report? Please explain. 4. What periodic trend in Ksp values do your experimental findings show? Did you find the same

periodic trend in Ksp values as the trend given in the lab report? 5. Based on your results, which of the following values of Ksp do you think is most likely for

beryllium hydroxide, Be(OH)2? Explain your choice. • 3.0 × 10-8 • 7.0 × 10-22 • 2.0 × 10-3

6. Based on your results, which of the following values of Ksp do you think is most likely for barium hydroxide, Ba(OH)2? Explain your choice.

• 3.0 × 10-8 • 7.0 × 10-22 • 2.0 × 10-3

7. a. How do beryllium hydroxide and barium hydroxide each compare to the Group II metal

hydroxides you did titrate in the lab in terms of ionic size and effective nuclear charge (Zeff) on the ions?

b. Based on the comparison of size and Zeff, which of the three values of Ksp listed in question 5 would Coulomb’s Law predict for beryllium hydroxide’s Ksp?

c. Based on the comparison of size and Zeff, which of the three values of Ksp listed in question 6 would Coulomb’s Law predict for barium hydroxide’s Ksp?

8. Of the five Group 2 metal hydroxides, Be(OH)2, Mg(OH)2, Ca(OH)2, Sr(OH)2, and Ba(OH)2, beryllium hydroxide and barium hydroxide are definitely the most toxic. The goal of this experiment was to determine the periodic trend in the solubility of Group 2 hydroxide compounds. Explain how the experiment you did today is more aligned with the philosophy of green chemistry than a similar experiment that involved titrations of all five Group 2 metal hydroxides.

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An&Environmental&Project&for&First&Year&Chemistry&Students:&

Analysis&of&Natural&Buffer&Systems&and&the&Impact&of&Acid&Rain.&&

David C. Powers, Andrew T. Higgs, Matt L. Obley, Phyllis A. Leber, Kenneth R. Hess, and Claude H. Yoder Department of Chemistry Franklin and Marshall College Lancaster, PA 17604 I. Experimental Procedure Due to the possible hazards involved in working with acids and carbonate salts, safety glasses should be worn at all times. Determination of Relative Importance of Components of the Carbonate Buffer Since the carbonate buffer system involves multiple equilibria, this set of experiments has been designed to enable students to understand the relative importance of each of the species involved. These experiments allow the individual buffering contributions of Na2CO3, NaHCO3, CaCO3, MgCO3, and CO2 to be measured and compared. Sulfuric acid is used in this project because it is the primary acid that is present in acid rain. Before preparing the experimental solutions, students need to standardize an acid solution to use in the titrations. A solution of approximately 0.03 M sulfuric acid should be used in these experiments. This standardization is performed in two steps: potassium hydrogen phthalate (KHP) is used to standardize a sodium hydroxide solution, which is then used to standardize the sulfuric acid. First, prepare an approximately 0.03 M sulfuric acid solution by appropriate dilution of concentrated H2SO4. Next, dissolve approximately 0.3 g of sodium hydroxide in distilled water and dilute the solution to 1 L. Before preparing the KHP solution, dry the KHP for about an hour in an oven set to approximately 115°C. Measure approximately 0.5 g of KHP to the nearest 0.001 g with an analytical balance. Dissolve the KHP in distilled water and dilute the solution to 1 L (volumetric flask). Using 50 mL of KHP solution, standardize the sodium hydroxide solution 3 times using phenolphthalein as an indicator. Calculate the average molarity and standard deviation. Show these results to your instructor for approval. Using 5 mL of the 0.03 M sulfuric acid solution and the standardized sodium hydroxide solution, standardize the sulfuric acid solution 3 times. Calculate the average molarity and standard deviation. Show these results to your instructor for approval. Phenolphthalein should again be used as an indicator in this titration. For some of the titrations in this project, a 0.003 M sulfuric acid solution is needed. This solution can be made by diluting part of the original sulfuric acid solution by a factor of ten. The pH of the experimental solutions was determined using a pH meter. By measuring pH as a function of the amount of H

+ added, the buffer capacities of each of the solutions can be determined. For solution

1, weigh approximately 0.130 g of sodium carbonate into a 400 mL beaker and dissolve in about 300 mL distilled water. Stir the solution until all of the sodium carbonate has dissolved. For solution 2, dissolve approximately 0.100 g sodium bicarbonate in 300 mL distilled water. This solution should also be stirred until all the material has dissolved. The masses of sodium carbonate and sodium bicarbonate are different because equal moles of the salts are being used. This allows the buffering action of carbonate and bicarbonate ions to be directly compared because the same number of moles of each is present in its respective solution. Titration of solutions 3 and 4 are designed to highlight the effect of an insoluble carbonate source on the buffer capacity of the solution. Prepare these solutions by weighing approximately 0.130 g calcium carbonate (Solution 3) and 0.110 g magnesium carbonate (Solution 4) into separate beakers with

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approximately 300 mL distilled water. These quantities have been selected to approximate the same number of moles of carbonate as solutions in 1 and 2. These solutions should be allowed to stir for approximately 1 hour before being titrated. The titrations of these solutions should be done allowing approximately 5 minutes between the additions of aliquots of acid to permit the insoluble carbonate salts to react with the added acid and the pH of the solution to stabilize. Analysis First, rationalize the different initial pH values for each of the titrations. After the titrations are completed, prepare a graph for each of the titrations labeling the x-axis as volume of acid added and the y-axis as pH. Use the graphs to determine the buffer capacity of each solution and rationalize the difference in the values. After buffer capacities have been calculated, compare the different systems based on the values of buffer capacity and explain the relationships that are observed between the numbers. If the solutions examined in this project are imagined to represent natural water systems, decide which of the systems would be a good buffer against acid rain and which would be poor buffers. If additional investigations are pursued, explain why the buffer capacity of a system increases linearly as the concentration of the buffer solution is increased. Use these data to predict the buffer capacity of buffer solutions of different concentration. If systems of mixed sodium salts are examined, explain why the carbonate portion of the titration curve is shorter than the bicarbonate portion. Explain why every carbonate ion has twice the buffering potential of each bicarbonate ion. Example Calculation: Method 1 Midpoint pH – 1 (NaHCO3 titration):

Midpoint pH = 6.4; Midpoint pH – 1 = 5.4

∆pH = 1 unit; ∆Volume between pH 6.4 and 5.4 = 7.75 mL (7.75 mL) (0.0572 M H

+) = 0.44 mmol H

+

Buffer Capacity: 0.44 mmol H+ / 0.300 L buffer solution = 1.5 mmol H

+ / L of buffer per pH unit

Method 2 Midpoint ± 0.5 (NaHCO3 titration): Midpoint pH = 6.4; Midpoint + 0.5 = 6.9; Midpoint – 0.5 = 5.9

∆pH = 1 unit; ∆ Volume between pH 5.9 and 6.9 = 10.25 mL

(10.25 mL) (0.0572 M H+) = 0.59 mmol H

+

Buffer Capacity: 0.59 mmol H+ / 0.300 L buffer solution = 2.0 mmol H

+ / L of buffer per pH unit

II. Background Acid rain is generally defined as rain having a pH lower than 5.0 (1). The primary acidic components of acid rain are H2SO3, H2SO4, and HNO3. These acids are formed through the dissolution of gaseous sulfur oxides and nitrogen oxides into raindrops as they fall to the ground (1). These acids can affect the pH of rain enormously because both H2SO4 and HNO3 are strong acids – both completely dissociate in water. The oxides that are responsible for acid rain are the byproducts of important industrial processes. Historically, the most common sources of sulfur oxides were the burning and roasting of commercially

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valuable metal sulfide ores such as those of nickel (NiS), copper (Cu2S), zinc (ZnS), lead (PbS), and mercury (HgS). These processes have been responsible for SO2 production. Because this traditional source of SO2 has been mediated, today most SO2 pollution stems from the burning of coal. Nitrogen oxides are generated through combustion, an example of which is the operation of car engines (1). Buffer capacity is the most significant variable in determining how acid rain will affect a water system. In general, a buffer is formed any time an acid or base is in solution with its conjugate. Buffer capacity is a measure of a buffer’s ability to resist change in pH upon the addition of another acid or base (2). Unfortunately, there is no standard, quantitative definition for buffer capacity. For the purposes of this paper, buffer capacity will be defined as the number of millimoles of hydrogen ion required to change the pH of one liter of buffer solution by one pH unit. The most significant buffer system in natural water systems is that which is governed by the equilibrium between carbonic acid and carbonate ions (3). Carbonate ions dissolved in water are a fairly strong weak base (Kb = 2.1 x 10

-4) (4). Since carbonate readily picks up a proton, we assume that virtually all

of the carbonate ion in solution will be converted to bicarbonate. A comparison of the Kb (2.4 x 10-8

) and Ka (4.8 x 10

-11) values of bicarbonate reveals that bicarbonate will also act as a base in solution (4).

The atmosphere contains a significant amount of CO2, which dissolves in water to form carbonic acid. Because the carbonic acid produces enough hydrogen ions to convert a portion of the carbonate ions into bicarbonate ions, the major buffering action of natural water systems is affected by the equilibrium between bicarbonate ion and carbonic acid (3). Carbonic acid is formed through the solution of atmospheric carbon dioxide; carbonic acid is in equilibrium with dissolved carbon dioxide (3).

CO2(g) CO2(aq) (eqn. 1) CO2(aq) + H2O(l) H2CO3(aq) (eqn. 2)

Carbonic acid is also in equilibrium with bicarbonate ions, which are formed by the dissociation of carbonic acid.

H2CO3(aq) + H2O(l) H3O+(aq) + HCO3

-(aq) (eqn. 3)

In addition, the dissolution of carbonate salts produces carbonate ions, which accept a proton from water and also add bicarbonate ions to the system.

CaCO3(s) Ca2+

(aq) + CO32-

(aq) (eqn. 4) CO3

2-(aq) + H2O(l) HCO3

-(aq) + OH

-(aq) (eqn. 5)

The addition of equations 1 through 5 yields the overall equation that accounts for the formation of bicarbonate ions and the equilibrium between atmospheric carbon dioxide and aqueous ions necessary for acid rain buffering.

CO2(g) + CaCO3(s) + H2O(l) Ca+2

(aq) + 2HCO3-(aq)

Acidic precipitation disrupts the equilibria that naturally occur between minerals, water, and atmosphere by adding significant quantities of hydrogen ions via the dissolution of the strong acids like H2SO4 and HNO3. The excess hydrogen ions are tied up by carbonate and bicarbonate ions in the water system to form carbonic acid. Clearly, this buffer system is only effective when a source of solid carbonate is readily available. Water systems in areas with large amounts of carbonate can withstand the introduction of large amounts of acid from acid precipitation without undergoing a correspondingly large decrease in pH. Bodies of water that do not have insoluble forms of carbonate available are drastically affected by acid rain as their buffer capacity is quickly overwhelmed. Lakes in areas where the bedrock has large amounts of carbonate ion like chalk, marble, or shale typically have a significant buffer capacity. Lakes in areas where the bedrock does not supply a significant amount of carbonate are particularly susceptible to acidification due to acid rain. Lakes resting on igneous bedrock, like the lakes on the granite Canadian Shield, have very little carbonate ion

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available. As a result, such bodies of water have low buffer capacity and consequently their ability to neutralize acidic precipitation suffers. These lakes readily demonstrate the environmental consequences of acid rain (5). References 1. Bunce, N. Environmental Chemistry, 2nd ed.; Wuerz Publishing Ltd: Winnipeg Canada, 1994; pp. 371-374. 2. Urbansky, E.; Schock, M. “Understanding, Deriving, and Computing Buffer Capacity,” J. Chem Educ. 2000, 77, 1640-1644. 3. Baird, C. Environmental Chemistry; W.H. Freeman and Co.: New York, 1995; pp 325-334. 4. Yoder, C.; Retterer, O.; Thomsen, M.; Hess, K. Interactive Chemistry; Academy Artworks: York, PA, 1999; p 650. 5. Bunce, N. Environmental Chemistry, 2nd ed.; Wuerz Publishing Ltd: Winnipeg Canada, 1994; p 385.

34"

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Electrochemistry and the Nernst Equation LEARNING OBJECTIVES

The objectives of this experiment are to . . . • construct galvanic cells and develop an electrochemical series based on potential differences between half-cells. • understand the Nernst Equation.%"BACKGROUND Any chemical reaction involving the transfer of electrons from one substance to another is an oxidation-reduction (redox) reaction. The substance losing electrons is oxidized and the substance gaining electrons is reduced. Let us consider the following redox reaction:

Zn%(s)%+%Pb2+%(aq)%%%%%%% Zn2+%(aq)%+%Pb%(s) This redox reaction can be divided into an oxidation and a reduction half-reaction:

Zn (s) Zn2+ (aq) + 2e- oxidation half-reaction and

Pb2+ (aq) + 2 e- Pb (s) reduction half-reaction

%A%galvanic&cell&(Figure%1)%is%a%device%used%to%separate%a%redox%reaction%into%its%two%component%halfTreactions%in%such%a%way%that%the%electrons%are%transferred%through%an%external%circuit%rather%than%by%direct%contact%of%the%oxidizing%agent%and%the%reducing%agent.%This%transfer%of%electrons%through%an%external%circuit%is%electricity.%Each%side%of%the%galvanic%cell%is%known%as%a%halfFcell.%For%the%redox%reaction%above,%each%halfTcell%consists%of%an%electrode%(the%metal%of%the%halfTreaction)%and%a%solution%containing%the%corresponding%cation%of%the%halfTreaction.%The%electrodes%of%the%halfTcells%are%connected%by%a%wire%along%which%the%electrons%flow.%In%the%cell,%oxidation%takes%place%at%the%zinc%electrode,%liberating%electrons%to%the%external%circuit.%Reduction%takes%place%at%the%lead%electrode,%consuming%electrons%coming%from%the%external%circuit.%

The%electrode%at%which%oxidation%occurs%is%called%anode.%

35"

"

The%electrode%at%which%reduction%occurs%is%called%the%cathode.%Since%oxidation%releases%electrons%to%the%electrode,%it%is%designated%the%negative%electrode%in%the%galvanic%cell.%Reduction%removes%the%electrons%from%the%cathode;%it%is%the%positive%electrode.%As%zinc%atoms%are%oxidized,%the%excess%positive%charge%(Zn2+%ions)%accumulates%in%solution%around%the%zinc%anode.%Likewise,%excess%3%negative%charge%(NO3T)%accumulates%around%the%lead%cathode%as%Pb2+%ions%are%removed%from%solution%of%%Pb(NO3)2%by%reduction%to%lead%metal.%These%excess%charges%create%an%electric%field%that%causes%the%ions%to%migrate:%positive%ions%(cations)%migrate%toward%the%cathode%and%negative%ions%(anions)%migrate%toward%the%anode.%In%order%to%make%this%flow%of%ions%between%the%two%halfTcells%possible,%the%cells%are%connected%by%aporous%barrier%(or%salt%bridge)%through%which%the%ions%flow.%The%barrier%prevents%free%mixing%of%the%two%solutions%but%permits%limited%movement%of%ions%so%that%the%electrical%neutrality%is%maintained%in%each%halfTcell.%%Different%metals,%such%as%zinc%and%lead,%have%different%tendencies%to%oxidize;%similarly%their%ions%have%different%tendencies%to%undergo%reduction.%The%cell%potential%of%a%galvanic%cell%is%due%to%the%difference%in%tendencies%of%the%two%metals%to%oxidize%(lose%electrons)%or%their%ions%to%reduce%(gain%electrons).%Commonly,%a%reduction%potential,%which%is%a%tendency%to%gain%electrons,%is%used%to%represent%the%relative%tendency%for%a%given%metal%ion%to%undergo%reduction.%%The%voltage%measured%in%the%cell%is%the%result%of%the%two%halfTreactions,%and%the%magnitude%of%the%potential%depends%on%the%concentrations%of%the%ions,%the%temperature,%and%pressure%of%gases.%When%all%the%concentrations%in%the%zinc/lead%system%are%1%molar%and%the%temperature%is%25%°C,%the%cell%voltage%is%0.63%volts.%It%would%be%a%monumental%task%to%assemble%a%list%of%all%possible%cells%and%report%their%voltage.%Instead%we%use%the%potential%of%the%halfTreactions.%We%cannot%measure%any%halfTcell%potential%directly,%so%we%pick%one%half%reaction,%call%it%the%standard,%construct%a%cell,%measure%the%cell%voltage%and%report%the%potential%relative%to%the%standard.%The%standard%that%has%been%chosen%by%convention%is:% 2 H+ (aq) + 2 e- H2 (g) E ° = 0.00 V Here the notation E ° is called the standard electrode potential and is the assigned potential of the standard hydrogen%electrode%when%the%concentration%of%H+%is%1%M"and%the%pressure%of%the%hydrogen%gas%is%one%atmosphere.%The%measured%cell%voltage%using%the%standard%hydrogen%electrode%is%therefore%the%potential%of%the%other%half%reaction.%%Tables%of%standard%halfTreaction%potentials%have%been%computed.%The%reactions%by%convention%are%written%as%reductions%and%hence%the%tables%are%called%tables%of%standard%reduction%potentials.%A%brief%example%follows%below%in%an%excerpt%from%a%Standard%Reduction%Potentials%table.%%The%greater%the%tendency%of%the%ion%to%gain%electrons%and%undergo%reduction,%the%less%negative%(or%the%more%positive)%the%reduction%potential%of%the%ion.%In%the%zinc/lead%cell,%the%lead%has%a%greater%tendency%to%undergo%reduction%than%the%zinc.

&

%

36"

"

&

&

&

&

&

%%In%the%

zinc/lead%cell%the%measured%potential%of%0.63%volts%can%be%deduced%from%the%sum%of%the%potentials%of%the%two%halfTreactions.%%

Zn (s) Zn+2 (aq) + 2 e- 0.76 V Pb+2 (aq) + 2 e- Pb (s) - 0.13 V

Zn"(s)""+"Pb+2"(aq)"" "Zn+2"(aq)""+""Pb"(s)"""" " " 0.63V"

Note: The sign of the standard reduction potential for the zinc half reaction is reversed to give the potential of the oxidation half reaction. In Part I of this experiment, other metal/ion half-cell combinations will be tried. From the data, a table will be developed, listing various elements and ions in order of their tendency to gain or lose electrons. &

The&Nernst&Equation&Theoretical%predictions%of%tendency%to%gain%electrons%are%used%to%predict%the%voltage%difference%between%two%electrodes.%The%voltage%difference%between%electrodes,%the%cell%voltage,%is%also%called%the%electromotive&cell&force,&or&emf&(or&E%).%Under%standard%conditions%(25%°C,%1%M"solution%concentration,%1%atm%gas%pressure),%these%theoretically%predicted%voltages%are%known%as%standard&emfs&(&or%E°cell&).%In%reality,%standard%conditions%are%often%difficult,%if%not%impossible,%to%obtain.%The%Nernst&Equation&allows%cell%voltages%to%be%predicted%when%the%conditions%are%not%standard.%Walter%Nernst%developed%the%following%equation%in%the%late%1800's%while%studying%the%thermodynamics%of%electrolyte%solutions:%

In%equation%(1),%R%is%the%gas%constant%(8.314%J%moleT1%KT1),%T%is%the%temperature%(Kelvin),%F%is%Faraday's%constant%(96,485%coulombs/mole),%n%is%the%number%or%electrons%transferred%in%the%balanced%oxidation/reduction%reaction,%and%Q%is%the%reaction%quotient,%or%([products]/[reactants]).%If%the%reactions%are%carried%out%at%room%temperature%(25%°C),%the%Nernst%equation%becomes%

Note in equations (1) and (2) that if the reaction quotient is equal to 1, then Ecell = E°cell.

Some&Standard&Reduction&Potentials&at&25°C&%

&HalfFreaction&% % %Potential&(volts)&%

Cu 2+ + 2 e- Cu (s) +0.34

2 H+ + 2 e- H2 (g) 0.00 Pb2+ + 2 e- Pb (s) - 0.13 Zn2+ + 2 e- Zn (s) - 0.76 Mg2+ + 2 e- Mg (s) - 2.37

Li+ + e- Li (s) - 3.05 "

37"

"

In Part II of this experiment, voltages will be measured at various solution concentrations for the copper/zinc galvanic cell and compared with voltages calculated using the Nernst Equation. SAFETY PRECAUTIONS Safety goggles must be worn in the lab at all times. Any skin contacted by chemicals should be washed immediately. &

EXPERIMENTAL&PROCEDURE&Part&I:&Galvanic&cells&and&the&electrochemical&series&1.%Into%a%central%cup%of%the%12Twell%container,%carefully%pour%approximately%2%mL%of%0.1%M"KNO3%.%2.%Into%the%four%wells%up%(#1),%right%(#2),%bottom%(#3)%and%left%(#4)%of%the%central%well,%pour%about%2%mL%of%the%metallic%salt%solutions%listed%below:%

#1%0.1%M"Cu(NO3)2%#2%0.1%M"ZnSO4%#3%0.1%M"Pb(NO3)2%

3.%With%clean%tweezers%(do%not%use%your%fingers)%take,%one%by%one,%four%strips%of%filter%paper%and%dip%one%end%into%the%central%cup%(where%immersion%in%the%KNO3solution%will%hold%one%end),%the%other%end%into%a%different%one%of%each%of%the%four%outer%wells.%Any%two%strips%of%filter%paper%from%any%two%of%the%four%outer%wells,%together%with%the%KNO3%solution%in%the%central%well,%make%your%salt%bridge.%4.%Hold%a%copper%metal%strip%with%clean%tweezers%and,%on%top%of%a%piece%of%scratch%paper,%sand%the%strip%to%remove%any%oxide%coating.%(DO&NOT&SAND&THE&STRIP&ON&THE&LAB&BENCH&!&!)%One%end%(2%cm)%of%the%strip%should%be%bent%and%immersed%into%the%Cu(NO3)2%solution%in%its%half%cell%(#1).%The%rest%(3%cm)%should%extend%out%to%the%edge%of%the%cell%and%should%be%bent%over%the%rim.%There%the%electrical%leads%(alligator%clips)%from%the%interface%will%be%attached%later.%Repeat%the%same%procedure%with%the%zinc%metal%strip%and%place%it%in%cup%#2.%Insert%an%inert%electrode%(nichrome%wire)%into%the%#3%solution,%which%contains%Fe2+%and%Fe3+%at%equal%concentrations.%Insert%a%strip%of%tin%foil%into%cell%#4.%The%pieces%of%metals%and%the%Fe2+/Fe3+%solution%will%be%the%electrodes%of%the%galvanic%cells.%The%tin%will%be%the%reference%electrode;%that%is,%we%will%measure%all%cell%voltages%relative%to%the%reduction%of%tin:%

Sn%2+%(aq)%+%%2%eT%% % %Sn%(s)%%% E°cell&=%0.00%V%%5.%Starting%with%the%Sn/Cu%combination,%measure%the%voltage%produced%from%the%galvanic%cells.%Using%the%red%and%black%lead%wires%(probes)%attached%to%the%red%and%black%posts%at%V,%secure%the%tin%foil%and%copper%metal%to%the%digital%volt%meter%with%the%alligator%clips.%Do%not%allow%the%clips%to%come%in%contact%with%the%solutions.%Position"the"probes"so"that"you"read"a"positive"voltage"for"the"Sn/Cu"system."Record%the%voltage%(note:&the&interface&displays&voltages&in&volts&when&using&the&voltage&leads).%Leave"the"probes"attached"only"long"enough"to"get"a"voltage"reading,"then"disconnect"to"minimize"

38"

"

chemical"changes"by"the"current"flow."The%red%probe%is%the%positive%terminal;%the%black%is%the%negative%terminal.%The%measured%voltage%will%have%a%positive%sign%if%the%black%probe%is%on%the%anode%and%the%red%probe%is%on%the%cathode.%Identify%and%record%which%metal%serves%as%the%anode%and%which%as%the%cathode.%6.%Measure%the%voltage%from%the%Sn/Zn%system%and%the%Sn/Fe%system,%always%having%the%same%probe%on%the%Sn%as%you%did%for%the%Sn/Cu%combination.%7.%Measure%the%voltage%of%the%Cu/Zn%cell,%Cu/Fe%and%Fe/Zn%cells.%Be%sure%to%keep%the%polarities%correct%as%you%do%this.%Part&II:&The&Nernst&Equation&In%this%part%of%the%experiment,%you%will%examine%the%effect%of%solution%concentration%on%the%cell%voltage%for%the%reaction:%%% Cu2+%(aq)%+%Zn(s)%%% % %Cu%(s)%+%Zn2+%(aq)% % (1)%The%Nernst%Equation%allows%you%to%calculate%E°cell&as%a%function%of%the%reactant%and%product%concentrations.%For%the%above%reaction%at%25%°C,%the%Nernst%Equation%becomes:%

Ecell%=%E°cell%T%(0.0591/2)%log%{[Zn2+]%/%[Cu2+]}%% % (2)%Remember,%solids%and%pure%liquids%are%not%included%in%the%Q%expression.%Theoretically,%E°cell&for%the%above%reaction%is%1.10%V.%Thus,%the%value%for%Ecell&can%be%calculated,%knowing%[Zn2+]%and%[Cu2+].%%1.%Set%up%five%zinc/copper%cells%using%the%following%Zn2+%and%Cu2+%solution%concentrations%(in%mole%per%%liter).%The%cells%should%be%assembled%in%cups%1%through%5,%with%Zn/ZnSO4%in%the%central%cup.%%%%%%%%%2.%Using%the%tweezers,%dip%five%filter%paper%strips%in%0.1%M"KNO3%in%a%small%beaker.%Use%approximately%5%mL%of%KNO3.%Insert%the%strips%with%one%end%in%the%center%cup,%the%other%end%into%one%of%the%outer%cups%just%before%you%are%to%measure%the%potential%of%that%combination.%3.%You%should%now%have%the%five%numbered%cups%each%containing%a%different%copper%solution%with%a%copper%strip%attached%as%in%Part%I.%The%central%cup%contains%1%M%zinc%solution%with%a%strip%of%zinc%placed%in%the%solution.%The%filter%paper%strips,%moistened%with%KNO3%act%as%salt%bridges%for%each%cell.%4.%Measure"the%voltage%(%Ecell&)%from%each%of%the%above%halfTcell%combinations.%Record%this%voltage%in%the%Data%Table%in%your%notebook.%%5.%Determine%Ecell&for%an%unknown%Cu2+%concentration.%%&DATA&ANALYSIS&Part&I:&Galvanic&cells&and&the&electrochemical&series&Make%a%table%that%indicates%the%relative%position%of%the%reduction%reactions%you%observed%with%respect&

Cell%#% [Cu2+]%(M)% [Zn2+]%(M)%1% 1.0% 1.0%2% 0.10% 1.0%3% 0.010% 1.0%4% 0.0010% 1.0%5% 0.00010% 1.0%

39"

"

to%the%tin%halfTreaction.%Place%the%reaction%for%your%most%positive%voltage%on%top,%and%your%reaction%for%your%least%positive%(or%most%negative)%voltage%on%the%bottom.%Part&II:&The&Nernst&Equation&Calculate"Ecell%%at%25%°C%for%each%of%the%above%halfTcell%combinations.%%Graph%Ecell&(observed)%versus%log[Cu2+].%%From%your%plot%of%Ecell&(observed)%versus%log[Cu2+],%determine%the%unknown%copper%ion%concentration%in%the%solution.%

40"

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The analytical determination of aqueous [Cu2+] Objectives:

• Utilize the redox chemistry of Cu2+

and the chelation chemistry of BCA to develop a method of copper determination in aqueous samples

• Understand both calibration curve and standard addition quantitative analytical techniques • Gain further experience in analytical spectrometry

Materials -1000 ppm Cu

2+ in water (3921 mg CuSO4·5H2O/L)

-100 g/L glucose solution (our reducing agent) (CAS#: 50-99-7) -bicinchocinic acid (BCA; CAS#: 1245-13-2) solution (Pierce or Sigma BCA protein determination reagent A). It contains BCA, sodium carbonate, sodium tartrate and sodium bicarbonate in 0.1M NaOH (pH 11.25)

Summary of reactions:

Cu2+

+ mild reducing agent (i.e. glucose) → Cu+ + oxidized agent (i.e. gluconate)

Cu+

+ 2 BCA → Cu+

-2BCA complex Procedure:

1. Gather 20 clean test tubes into a rack, in a 4x5 matrix. 2. Make serial dilutions of the glucose solution. You should have three sugar solutions of 1, 0.1 and 0.01g/L. 3. Make serial dilution of the copper solution. You should have four copper solutions of 1000, 100, 10, and 1 ppm. 4. Put 1 ml of BCA solution in each test tube 5. Put 1 ml of 1.0 g/L glucose in five tubes 1 ml of 0.1 g/L glucose in another five tubes 1 ml of 0.01 g/L glucose in another five tubes 1 ml of deionized water in the last set of five tubes 6. Put 0.1 ml of 1000ppm Cu

2+

in four tubes (one of each sugar concentration) 0.1 ml of 100ppm Cu

2+

in another four tubes 0.1 ml of 10ppm Cu

2+

in another four tubes 0.1 ml of 1ppm Cu

2+

in another four tubes 0.1 ml of deionized water in the last set of four tubes

Note: Since glucose can reduce Cu+ to Cu

o, it might be important to consistently add the copper solution to

the premixed glucose and BCA solutions. 7. Incubate tubes at 60

o

C for 15 minutes. Observe mixtures occasionally, and note which tubes develop color the most rapidly. BE SURE TO PARAFILM YOUR TUBES. 8. Take absorbance readings every 20nm from 700 to 350nm of your most colorful tube. (Dilute if necessary, and scan again).

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"

Pre-lab questions

1. Describe in detail the exact steps that you will take in lab to determine the concentration of an

unknown aqueous Cu2+

sample using glucose and BCA solutions, by calibration curve (measure at least 7 calibration points) AND use standard addition methodologies. (Your unknown will have a concentration greater than 1 and less than 200 ppm.)

2. Describe what circumstances might make a standard addition methodology more desirable than a calibration curve method.

3. Briefly explain how you would alter the calibration curve technique to determine the concentration of glucose in a urine sample.

4. Briefly explain how you would alter the calibration curve technique to determine the concentrations

of both Cu+ and Cu

2+ in a mixed sample.

5. Balance the following redox equation; note that it only occurs in basic solutions of water (i.e. you

might have to include OH-and H2O as reactants and/or products to get it to balance):

Glucose (C6H12O6) + Cu2+ → Gluconate

-(C6H11O7

-) + Cu

+

Procedure:

a. You will be given a solution of unknown [Cu2+

], between 1 and 200ppm. b. You will find the concentration using the method that you develop in prelab question 1. I will only provide the glucose, copper sulfate and BCA solutions listed above, and spectrophotometers. c. You will repeat the measurement three times (including repeating all of the calibration measurements), and then find a mean and standard deviation approximation of the measured values. d. At the end of the period, the unknown concentration will be revealed.

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Equilibrium: Measuring the soil/water distribution coefficient: the effect of decreasing pH Background:

In this lab exercise you will measure equilibrium constants for Cu2+

binding to soil at varying pH’s to demonstrate Le Châtelier’s principle and to determine whether soil pH will affect our phyto-

remediation projects. In this exercise, we focus on the water/soil adsorption equilibrium of Cu2+

. To mathematically describe this equilibrium, environmental chemists have defined the soil-water distribution coefficient, Kdist, which is the ratio of the equilibrium concentration of the pollutant in the solid (hydrated colloid) soil-bound phase to that in the unbound aqueous phase:

Kdist = Csoil/Caqueous where Csoil is reported in units of mg pollutant per kg soil, and Caqueous is reported as mg pollutant per L solution. Kdist is therefore useful in determining the concentrations of pollutants in each phase, and is used extensively in modeling their fate, transport, and bioavailability in aquatic systems and ground waters.

You will be adding a Cu2+

solution of known concentration to some soil and then determining the

concentration of Cu2+

in the water left over after you filter it. The amount of Cu bound to the soil can then be calculated, along with the Kdist. When the pH is varied, you can measure changes in the equilibrium constant. When, instead, the temperature is varied, you can find the ΔH and ΔS for binding Cu to soil.

Pre-lab questions:

1. Suppose that you started with 5 ml of 100 ppm Cu2+

solution and added it to 0.1g of soil. After mixing, you filter it and find that the concentration of the filtered solution is 5 ppm Cu

2+. From

the above scenario, calculate the mass of Cu2+

bound to the soil. 2. Calculate the Kdist.

Objectives

• To measure an equilibrium constant • To test the effects of altering pH on that equilibrium constant • To determine the effects of pH changes on soil chemistry and apply that knowledge to our

phytoremediation projects.

Materials

• Soil, balances, pipettors, hot water bath @60oC, pH sensitive paper indicators

• 1000ppm Cu2+

, BCA and 100g/L Glucose stock solutions • M solutions of HCl and NaOH • Filter papers and syringe filters

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"

Precautions: CuSO4 is mildly toxic; HCl and NaOH are caustic; wear gloves, eye protection and appropriate apparel.

Procedure (This lab exercise is based on that presented by Dunnivant and Kettle, J. Chem. Ed. 2002, 79, 715-717)

1. In three different test tubes, measure out 0.1g of potting soil (use homogeneous soils, i.e. filter out vermiculite and larger pebbles or pieces of organic materials.) 2. Add 5 ml of 100 ppm Cu to each tube. 3. Add 1.0 ml of the following three different solutions to the three differently labeled tubes: Water

0.1 M HCl 0.1 M NaOH

4. Filter a. First through folded paper in a funnel. (You only need a milliliter or so.) b. Then through a 0.2µm syringe filter. 5. Measure the pH in each of the three filtered solutions (use papers to estimate pH)

6. Measure Cu2+

in each of the three filtered solutions by standard addition to the BCA assay, like this: a. Set up 9 different clean test tubes as in the table below: b. Add 1 ml of BCA solution to all 9 tubes c. Add 1 ml of 100 g/L glucose to all 9 tubes

d. Put tubes in the 60oC water bath for 5-15 minutes.

e. Measure Absorbance at 560nm for each tube

f. Calculate the concentration of Cu2+

in each sample by graphing a Beer’s Law plot.

Tube # ml of filtered sample ml of standard 100 ppm Cu2+

Increase in [Cu] (ppm)

[Cu] (ppm) in tube

Water 1 0.1 0 0 [Cu] Water 2 0.1 0.01 10 [Cu]+ 10 Water 3 0.1 0.02 20 [Cu]+ 20 HCl 1 0.1 0 0 [Cu] HCl 2 0.1 0.01 10 [Cu]+ 10 HCl 3 0.1 0.02 20 [Cu]+ 20

NaOH 1 0.1 0 0 [Cu] NaOH 2 0.1 0.01 10 [Cu]+ 10 NaOH 2 0.1 0.02 20 [Cu]+ 20

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"

Cu2+ (ppm) 7. Calculate

a. the mass (in mg) of Cu2+

bound to the dirt in all three samples b. the Kdist in all three samples

Things to include in the report: a. Make a graph of Kdist vs. pH b. Explain the trend (if any) using chemical equations and Le Châtlier’s principle. c. What do the results mean for our phytoremediation projects (if anything)?

45"

"

Equilibrium: Measuring the soil/water distribution coefficient: the effect of increasing T

Background: In this exercise you will focus on the water/soil adsorption equilibrium of Cu

2+ as it changes with

temperature. Like you did last week, you can find Kd experimentally. The difference this week is that we will look at three different temperatures, and use the data to calculate thermodynamic parameters, including the free energy, enthalpy and entropy of the adsorption of Cu

2+ to soil. Kd can be treated as a

chemical equilibrium constant (however, see the note below). As such, Kd can be used in the thermodynamic equation:

ΔGo = -RT lnKd (2)

Where ΔGo, the standard state spontaneity criterion for chemical reactions, is composed of both

enthalpic and entropic considerations: ΔG

o = ΔH

o - TΔS

o (3)

Combining equations (2) and (3) by substitution and solving for lnKd, we arrive at the Van’t Hoff equation:

lnKd = -ΔHo/(RT) + ΔS

o/R (4)

The Van’t Hoff equation has the structure of y = mx+b, and thus, when Kd is measured at several temperatures, a plot of lnKd vs. 1/T is linear, and yields a slope of -ΔH

o/R, and a y-intercept of ΔS

o/R.

Prelab:

1. How would you determine the ΔS and ΔH for Cu2+

binding to soil given Kdist values at two different temperatures? 2. Based on your understanding of enthalpy, what do you predict is the sign of ΔH for the binding of Cu

2+ to soil? Briefly explain your answer.

3. Based on your understanding of entropy, what do you predict is the sign of ΔS for the same reaction? Briefly explain your answer. 4. Based on your understanding of free energy, what do you predict is the sign of ΔG for the binding of Cu

2+ to soil? Briefly explain your answer.

5. Based on the experiment that we performed last week, the prelab questions above, and the materials noted below, write a brief procedure for finding the sign on ΔG, ΔS and ΔH for the binding of Cu

2+ to

soil (Cu2+

(aq) + soiln-

Cu2+

soil n-

). Materials

• Soil, balances, pipettors, hot water bath @60oC

• 1000ppm Cu2+

, BCA and 100g/L Glucose stock solutions • Filter papers and syringe filters

46"

"

Objectives

• Apply thermochemistry (ΔG, ΔS and ΔH) principles to an environmental example • Gain a deeper understanding of enthalpy and entropy and their relationships to equilibrium

Questions to answer in the report

1. Is the reaction endothermic or exothermic? Explain how you know. 2. Is the reaction favorable or unfavorable? Explain your answer. 3. Do the signs on ΔS and ΔH make sense for this reaction? 4. Do your results match your pre-lab predictions? If not, propose a model to explain your

findings. 5. What do the results mean for our phytoremediation projects (if anything)?

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"

Phytoremediation Analysis In our pots that we contaminated with Copper (II) sulfate back in February, we have a complex mixture of Cu

2+ in solution, Cu

2+ bound up as mineral salts and/or bound to colloidal soil particles, Cu

+ in

solution and bound (microbial activity in the soil can reduce Cu2+

to Cu+), and copper ions absorbed

into the plant tissues. In order to determine whether or not our plants have been cleaning up the soil for us, we need to measure the copper content in the soil and/or in the plant tissues, and compare those concentrations to controls. In order to measure copper in soil, the trick is to ‘digest’ it with concentrated acid so that all (or at least most) of the copper is leached off of the soil (recall the pH dependence of the Kdist), and released from ionic compounds such as Cu(OH)2, and then filter it prior to readjusting the pH to neutral so that we can measure [Cu

2+] and [Cu

+].

You might also have noted that the plants may or may not have thrived in copper. This might be evident in your experiments, based on reduced growth, browning (loss of chlorophyll), or some other visible affect. We can determine the chlorophyll content through UV-Vis spectroscopy. Prelab assignment: A. Which eight samples will you analyze for Cu using the above procedures? Soil from which pots? B. Will you measure Cu total in eight different samples, or find both [Cu

2+] and [Cu

+] in four samples?

Will you measure chlorophyll in your plant samples in addition to (or instead of?) copper? Will you measure plant size (mass or length)? C. Give a rationale for your selection of your eight samples and the test(s) that you hope to perform on them: what do you hope to learn from those eight? D. What are your expected results? Explain. Procedure for finding [Cu

2+/+] in Soil:

1. Remove a small representative portion of soil and place it in a drying oven for at least one hour (It may take much longer for some larger samples. You are welcome to come in at any time to start plant tissues drying.) 2. Weigh out 0.1-0.15g of dried soil into a small beaker. 3. In the fume hood, add 4-6 ml (40 times the mass) concentrated HCl (use goggles and gloves and take extreme caution!) to each soil sample. 4. Digest for at least 20 minutes with frequent agitation. 5. Filter through paper, so that you have at least 1mL of filtered material 6. Determine [Cu

2+] + [Cu

+] in each sample by standard addition (remember how you do this?) You

might need to add a few drops of NaOH if the solutions look cloudy and yellow, instead of light purple. If NaOH is required, add this BEFORE incubation. If you would like to determine how much [Cu

2+] you have relative to [Cu

+], use another portion to

perform the BCA assay without the addition of glucose (the reducing agent), to measure [Cu+]

48"

"

Procedure for finding [Cu2+/+

] plant tissue: Copper in plant tissue will be bound up with anionic charges in proteins and polysaccharides. Therefore, in order to measure [Cu] in plant tissue, you should first dry the material completely, then grind it to small particles, and finally digest it with acid just like you would with soil. Remove a piece of plant tissue (you may use leaves, stems or roots) and place it in a drying oven for at least 60 minutes, maybe as much as 4 hours. (It will take a lot of plant material to get to 0.1 grams dried weight!) Grind the dried plant material with a mortar and pestle to a fine powder. Treat the plant material as you would soil starting at step 2 above. Procedure for finding [Chorophyll a] in plant tissue.

a. Wash leaves with water. Dry well with paper towels. b. Weigh 0.2-0.4 g of leaves into small bottle (record the mass for each sample). c. Fill bottle with 2-4 ml of methanol (10 ml methanol /g leaf tissue). d. Let samples sit at room temperature on the rocker shaker overnight. e. Decant into UV-Vis cuvette and measure absorbance at 418 nm with a UV-Vis spectrometer. f. Use an extinction coefficient of 110700 cm

-1M

-1 and the Beer-Lambert equation, to estimate

[chlorophyll a] in each sample. Assume that the leaves have a density of 1.00 g/ml. (Note: did you account for the dilution factor in making your determination?) http://omlc.ogi.edu/spectra/PhotochemCAD/abs_html/chlorophyll-a(MeOH).html