General Chemistry Atoms First Plus MasteringChemistry 2E John E McMurry Solution Manual

Type:

Solution Manual

Resource:

General Chemistry Atoms First Plus MasteringChemistry

Edition:

2nd Edition

Author(s):

John E. McMurry

To the Instructor: When looking through the instructor versions of the other laboratory manuals currently available, my first thought was that I wouldn’t have time to wade through even 10 percent of the material contained within them. The instructor versions often contained detailed discussions of concepts I didn’t need (I have a degree in chemistry too) and redundant materials already available in the student manual. The instructor versions also seemed to lack the kind of information that I would really like to have, e.g., potential problems that can occur when the students performed the experiments or notes on points to stress that I could give to my teaching assistants. So when designing this instructor manual, I tried to include all of the information and materials I would like to have and left out those materials that seemed unnecessary.

For each experiment in the lab manual you will find the following in the instructor version:     

A list of chemicals, equipment, and supplies Preparation instructions for chemicals and reagents as needed A discussion of points to stress and potential problems to be aware of when performing the experiments A key for the pre-laboratory assignments A key for the post-laboratory assignments

The keys provided do not include a grading scale as I believe that different instructors will want to emphasize different portions of the assignments and reports. The reports in the manual are already delineated as to what the students should be including and will be based on their experimental data so there is no key provided for them beyond the identity of unknowns etc.

It is my hope that you find this instructor manual concise, to the point, and above all, useful. If I have been remiss in including anything that you feel would help, please let me know. If you need any further assistance or have suggestions or comments regarding the use of the manual, the performance of the experiments, or the use of the keys please do not hesitate to contact me.

Stephanie R. Dillon Coordinator General Chemistry Laboratories Florida State University sdillon@chem.fsu.edu

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Table of Contents To The Instructor 0)

Part A. Measurement and Expression of Experimental Data Part B. Graphing and Statistical Analysis A math review, including practice with significant figures, rounding rules, graphing techniques, and statistical analysis.

1)

Conservation of Matter A reaction of copper nitrate and zinc is observed to determine if mass is conserved.

2)

Atomic Spectra Several unknown light sources are identified by use of spectrometry.

3)

Halogen Reactions Hexane and water are used to determine the solubility of the halogens and halides.

4)

Paper Chromatography: An M&M’s® True Colors Paper chromatography is used to analyze the food colorings used in M&M’s candy shells.

5)

Molecular Models Lewis structures are presented in model form and investigated.

6)

Limiting and Excess Reagents Copper nitrate and potassium iodide are reacted at different concentrations to determine which is the limiting reagent and which is in excess.

7)

Redox Reactions in Voltaic Cells: Construction of a Potential Series Micro-scale unit cells of several metals are created and their potentials are measured in order to create a potential series.

8)

Reactions in Aqueous Solutions: Strong Acids and Bases The pH and concentration relationships for the reaction of HCl and NaOH are investigated.

9)

Calorimetry & Hess’s Law The enthalpies of reaction are determined for both magnesium metal and magnesium oxide in water. These enthalpies are then used in conjuction with Hess's Law to determine the enthalpy of reaction for magnesium burning in oxygen. .

10)

Gas Laws and Air Bags Ideal gas law and stoichiometry are used to construct a working mock-up of an automobile airbag.

11)

Intermolecular Forces and the Triple Point of CO2 Intermolecular forces are investigated by observing the temperature depression created when different solvents evaporate. The triple point of CO2 is observed and measured.

12)

The Purification of Water Fresh water samples are tested for phosphate before and after purification.

13)

Kinetics: The Iodine Clock The Iodine Clock Reaction is measured at room, high, and low temperatures, and its rate constant is determined.

14)

Photometric Determination of an Equilibrium Constant The equilibrium constant for the reaction of iron(III) thiocyanate is determined by photometric methods.

15)

Titration of 7-Up® Monoprotic and polyprotic acids are titrated using NaOH.

16)

Hydrogen Phosphate Buffer Systems Students create and then test the buffering capacity of a hydrogen phosphate buffer.

17)

Entropy, Free Energy, and Chemical Equilibrium The thermodynamics of the solubility of Ca(OH)2 are determined using titration.

18)

Electrochemistry: The Nernst Equation Unit cells are built from various metals and then tested at several temperatures. The free energy is determined using the Nernst equation.

19)

Analysis of a Hydrate The formula and percentage of water is determined in an unknown hydrate.

20)

Qualitative Analysis: Cations, Anions, and Complex Ions A single-day, microscale qualitative analysis is used to determine the cations and anions in an unknown mixture.

.

21)

Isomerism in Coordination Chemistry A study of the three-dimensional structures of octahedral complexes is performed.

22)

Light Emitting Diode Students will identify the components of a complex ion and predict if that complex ion has a large or small crystal-field splitting energy while constructing an organic diode.

23)

Amino Acid Analysis by Ninhydrin Students identify the amino acids in an unknown mixture using a combination of ninhydrin and TLC.

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Experiment 1: Measurement and Expression of the Experimental Data Instructor Notes and Lab Preparation: Each Fall and Spring semester at Florida State University starts one week before a national holiday, Veterans Day and Martin Luther King Day, respectively. This means that we manage to check in the labs, but then have an off week before the students are able to start performing experiments. Rather than allow the students to waste a full week of lab time, I always assign the following math and graphing review as the first experiment of the semester. Part A is designed for the first semester, and Part B is designed for the second semester of general chemistry.

As the check in process is generally a short one, I have each of my teaching assistants give a short lecture on the math involved at the end of the check in day.

Particular points to stress: Scientific or exponential notation is used so often in scientific literature and calculation that I think we often forget how confusing it can be to someone dealing with it for the first time. I always stress that students should look at the values they calculate to make sure they make sense.

For example, one of the most common mistakes a student will make when dealing with exponents is when they must change an exponent in order to add or subtract two exponential values. The process requires the student to “move the decimal” in the correct direction to “match up” the exponents. The most common mistake is the movement of the decimal in the wrong direction, so ask the students to check and make sure that the number resulting after the move still makes sense, (i.e., has the same value it did before the move). Simply plugging it into a calculator should do the trick.

Another sticky point with my students is an overall lack of graphing skills. I have placed a brief Excel® tutorial at the end of the manual to encourage students to use this program, or one similar to it to complete the graphing exercises. If you choose to have your students complete the graphs by hand, I understand the reasoning, but I would encourage you to consider asking for the graphs in both styles. The reality of the world today is that computers are everywhere. Skillful use of technical programs like Excel is a must for scientific analysis.

The most common mistakes that should be discussed with respect to graphing are 1) attempts to fit straight lines to data that does not fit y = mx + b; and 2) irregular spacing of values on the axis. I often see graphs whose axes are marked at regular intervals with the values given as part of the problem. Markings of this sort always generate perfectly straight lines, which thrill the students, but give me migraines.

Last but not least, I have found that our discussion of logarithms and the proper use of significant figures in logarithms is generally the first exposure the students will have to the subject. Special attention should be paid to the introduction of this material. .

Lab Report Key Exercises 1A Key 1. For each of the following numbers: a) Indicate the number of significant figures in the number; b) Round the number to 2 significant figures; and c) Place the number into scientific notation: 3.0 x 10-3 a. 0.0030405 1.2 x 106 b. 1,237,888 1.0 x 10-1 c. 0.10101 1.5 x 101 d. 15.0005 e. 0.000009995 1.0 x 10-6 2. Complete the following calculations using the correct number of significant figures and order of operations: o a. Convert 432 K to oF. 318 F  %T  b. Convert 1.67 %T to Absorbance (A) using A   log   • A = 1.77  100   P  HVap  1 1  c. Solve for P2 using ln  1      , where P1 = 1.5 atm, T1 = 273K, T2 R  T2 T1   P2  = 400K, HVap = 40.7 kJ/mol, and R = 8.3145 J/mol • K P2= 445 atm d. Calculate the total mass in grams for: 4.445 x 10-3 g + 15.3 mg + 9.77561 x 10-2 mg = 19.8 mg e. Calculate the total pressure in kPa for: 2.34 atm + 788 Torr – 650 mmHg = 256 kPa 3. Calculate the average and standard deviation for the following set of calibration data: 66.788, 66.898, 66.999, 65.998, 65.778, 67.434, 67.311, 66.009

Average = 66.652

Standard Deviation = 0.638

4. The table of data below was collected using a helium lamp to calibrate a spectroscope used in this general chemistry lab. Use the data to create a calibration graph and then use the graph to calibrate the experimental data collected for other light sources: Color

Wavelength (nm)

Violet Blue Green

388.9 468.6 501.6 .

Spectroscope Reading (AU) 3.78 4.75 4.93

Yellow Orange Red

587.6 667.8 706.5

5.82 6.55 6.83 a. Sodium line spectroscope reading = 5.80, corrected wavelength =

__563_______ b. Hydrogen line spectroscope reading = 6.43, corrected wavelength =

_____626____ c. Mercury line spectroscope reading = 5.69, corrected wavelength =

__552_______ d. and line spectroscope reading = 4.65, corrected wavelength =

____448_____

5. Create a graph of the Absorbance (Optical Density) versus Molar Concentration and answer the questions about the resulting line.

Concentration (M)

Optical Density

0.19

0.15

0.34

0.34

0.49

0.52

0.64

0.68

0.79

0.81

0.94

0.92

1.09

1.02

.

1.24

1.09

1.39

1.14

1.54

1.17

NOTE: The Absorbance, or Optical Density, of a solution for a particular wavelength of light is given by log(Io/I), where Io is the incident light on a sample, and I is the transmitted light. Beer's law states that the Optical Density of a substance in solution is directly proportional to its concentration (i.e., has a linear relationship). Not all substances follow Beer's law, but a calibration graph of Optical Density versus concentration can still be used to measure concentrations of unknown solutions. a. Does the solution in your graph follow Beer's law? Not

entirely, curved line at higher concentrations. b. c. What is the concentration of a solution that has an Optical Density of 0.98? 1.0 M d. e. What is the concentration of a solution that has an Optical Density of 0.71? 0.68M

6. Robert Boyle used an open-ended barometer (similar to that in the figure) for the experiments that led to the discovery of Boyle's law, describing mathematically the effect of pressure on the volume of a gas. He determined the nature of the law by graphing his data. The law states that PV = constant, where P is the total pressure, which is the sum of Patm + PHg. The pressure can be measured in Torr, which corresponds to the height of the Hg column in mm.

The following table contains some data in an experiment similar to that of Boyle's, though the open-end manometer is considerably larger. The volume of the gas in the closed end of the tube is measured in mL, and the pressure recorded is PHg, the height of the column, not the total pressure. .

Pressure (torr)

Volume (mL)

58

16.5

162

13.9

239

12.4

316

11.1

385

10.1

468

9

536

8.3

620

7.4

733

6.5

You are to analyze this data by graphing it, discovering that a direct graph (V vs. P) gives a curved line. Scientists often try manipulating data in a way to give a straight line, and when the data is re-graphed as (P vs. 1/V), you will see that a straight line is produced, revealing the nature of Boyle's law (that pressure and volume are reciprocally related). Further analysis of this straight line will also reveal the atmospheric pressure at which the measurement is made. Since Ptotal *V = constant Ptotal = constant(1/V) Patm + PHg = constant(1/V) PHg = constant(1/V) - Patm the graph is of the form y = mx + b, where the slope, m, gives the value of the constant, and the P intercept is the negative of the atmospheric pressure. a.

What is the slope of the line (i.e., the constant in the Boyle's law equation)? Give units in your answer. y = 7202.5x - 347.77 R² = 0.9925

b. What is Patm (i.e., the negative of the P intercept)? Give units in your answer. 348 Torr 7.

Using the same data as above, calculate the logarithm of the volume. Graph the Pressure versus the log10 Volume and answer the questions about the resulting line. a. If this graph is a straight line then it is of the form P = mLog10V + b, in which m is the slope and b is the y-intercept. Is the graph a straight line or a curved line (nonlinear variation)? Straight line b. If possible, calculate the equation of the line for the graph. Show your work.

.

y = -1669.9x + 2071.7 R² = 0.9972

c. Using the resulting equation from question 2, what would be the volume of a system that had a pressure of 1.9 atm?

1444 mmHg = -1669.9x + 2071.7; -2297.6 mL

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Exercises 1B Key 1. Consider the following experiment in which you wish to determine the molarity of the concentrated HCl in your new bottle of reagent. You carefully measure out 100.0 mL of the concentrated acid and dilute it to 1.000 L (in a one-liter volumetric flask). You then use a 5.00 mL volumetric pipette to remove 9 aliquots for titration with 0.500 M sodium hydroxide solution. The following are the volumes of base needed for each sample.

14.75 mL 14.70 mL 14.81 mL 14.73 mL 14.73 mL 14.76 mL 14.48 mL 14.55 mL 14.69 mL Calculate Qexp for both the maximum and minimum values in this data. (a) Qexp for 14.81 mL = 0.1818 (b) Qexp for 14.48 mL = 0.2121 (c) Can you reject 14.81 mL with 90% confidence? No (d) Can you reject 14.81 mL with 95% confidence? No (e) Can you reject 14.81 mL with 99% confidence? No (f) Can you reject 14.48 mL with 90% confidence? No (g) Can you reject 14.48 mL with 95% confidence? No (h) Can you reject 14.48 mL with 99% confidence? No .

Average your measurements (throwing out the outlier(s), if justified, with 90% confidence), and calculate the molarity of the original concentrated HCl solution.

Average = 14.69 mL

[HCl] = 0.073 M

2. You are asked to calibrate a 10-mL volumetric pipette by weighing to the nearest 0.1 mg the mass of water delivered by the pipette. You weigh five samples of water delivered by the pipette and convert the mass of each to volume by dividing by the density of water at 25 oC (0.997048 g/mL). Following are your measurements: 9.9755 mL

10.0029 mL

9.9899 mL

10.0302 mL

9.9963 mL

Calculate the following statistical measures for this data: Mean (x) = _____9.9990__mL____________

Standard Deviation(s)= ____0.02019 mL__

Variance (s2) = ___0.000408__________Standard Error of the Mean (smean) = ____0.000182___ 90% Confidence Interval = 9.9899 ± _0.000388_________ (The "true value" of the volume will lie within this interval 90% of the time.)

95% Confidence Interval = 9.9899 ± ___0.000507_____________ (The "true value" of the volume will lie within this interval 95% of the time.)

99% Confidence Interval = 9.9899 ± ____0.000839____________ (The "true value" of the volume will lie within this interval 99% of the time.)

3. The table below shows some data in which the peak height of triplicate samples of a substance is listed as a function of the quantity of material injected. First you would like to know how good the precision of the method is by determining the standard deviations and confidence intervals of each set of triplicate samples. Then, you would like to know whether the means of the triplicate samples show a linear relationship with sample size, and the equation for this relationship.

Gas Chromatograph

Calibration

Sample Mass μg

Peak Height (cm) Trial 1

Trial 2

Trial 3

15

1.4

1.2

1.3

25

2.6

3.1

3.7

35

5

5.4

4.8

45

7.8

7.6

7.3

55

10.1

9.8

10.3

65

12.5

11.9

12.6

.

a.

Using Excel or another similar spreadsheet program, graph the average peak height versus Sample Mass. Apply a best fit line to the resulting graph.

b. Again using the spreadsheet program, calculate sample mean, variance, standard deviation, standard error of the mean, and confidence intervals for the data above.

Gas Chromatograph Calibration Sample Peak Height (cm) Mass 99% 95% 90% μg Trial Trial Trial Mean Variance Standard Standard 1 2 3 Deviation Error in confidence confidence confidence Interval Interval Interval the Mean 15

1.4

1.2

1.3

25 35

2.6

3.1

3.7

5

5.4

4.8

45

7.8

7.6

7.3

55

10.1

9.8

10.3

65

12.5

11.9

12.6

1.30 3.13 5.07 7.57 10.07 12.33

0.01 0.30 0.09 0.06 0.06 0.14

0.10 0.55 0.31 0.25 0.25 0.38

0.06 0.32 0.18 0.15 0.15 0.22

0.17 0.93 0.52 0.42 0.42 0.64

0.25 1.37 0.76 0.62 0.62 0.94

c. Answer the following questions: i.

For the 45 μg sample, what is the sample standard deviation and 90% confidence interval of the three peak height measurements? See above. ii.

Excel gives the best straight line fit to the data as an equation y = mx + b.

What are the slope and intercepts of this line? (Be sure to include units.)

.

0.57 3.15 1.75 1.44 1.44 2.17

Slope = 0.2242 cm/g; Intercept = - 2.3898 cm 4. During this semester, you will encounter several sections that discuss kinetics. Kinetics is the study of the rate at which a reaction proceeds from reactants to products. These rates can only be determined experimentally by direct observation of the appearance or disappearance of products or reactants, with respect to time. Catalysts are compounds that increase the rate at which a reaction proceeds, and inhibitors are compounds that slow down or stop reactions.

Enzymes are biological catalysts. In discussing the properties of an enzyme, certain values, or parameters, are determined experimentally under steady state conditions. The values are determined through kinetics studies and include: VMax : The so-called maximal rate of the catalyzed reaction. The enzyme's active site is saturated.

KM : The Michaelis constant is the substrate concentration at which the reaction rate is one half its maximum value. Also known as the turnover number.

These values are determined experimentally by recording the progress of an enzyme-catalyzed reaction using fixed amounts of enzyme and a series of different substrate concentrations.

A typical data set looks like Table 1 where Vo is the initial reaction velocity, and [S]o is the substrate concentration.

VMax and KM can be determined from linear regression analysis of a plot of 1/Vo vs. 1/[S]o, a socalled Lineweaver-Burk plot.

K  1  1 1 The Lineweaver-Burke Equation: has the familiar form y = mx + b,  M    V0 VMax  S 0  VMax where m =

KM 1 and b = .

VMax VMax

a. For the data below, create a Lineweaver-Burk plot of the data, and calculate VMax and KM. So (mol/L)

Vo (mol/L.min)

0.07

0.155

0.075

0.185

0.09

0.285

0.15

0.555

0.25

0.795

0.55

1.135

1.05

1.375

.

Vmax = 6.61 mol/L• min; Km = 2.67 mol

5. In biological systems, the enzyme-catalyzed reaction rates are often affected by substances that inhibit or interfere with the enzyme’s interaction with the substrate(s). There are three basic types of inhibition: noncompetitive, uncompetitive, and competitive inhibition.

The type of inhibition can be determined from graphical analysis of experimental data with Lineweaver-Burk plots. Inhibition experiments involve a series of experiments with fixed amount of inhibitor added to varying amounts of substrate. Lineweaver-Burk plots are constructed showing multiple lines for the various inhibitor concentrations.

The inhibition types and their graphical characteristics are summarized below, and graphically in example graph 1.

Reference Table 1: Definition and Graphical Characteristics of Enzyme Inhibition Inhibition type Definition Graphical characteristics

Noncompetitive

Competitive

Uncompetitive

The inhibitor binds the enzyme at a different site than the substrate causing a conformational change.

The conformational change affects the rate of catalysis, but the overall turnoff remains constant.

The inhibitor competes for the enzyme’s binding site with the substrate. The proportion of substrate molecules bound by the inhibitor reduces the rate of catalysis.

The inhibitor binds the active site after the substrate. Binding of the inhibitor can stimulate the binding, but produces a nonproductive complex.

.

The slopes of Lineweaver-Burk plots are different, as are the yintercepts.

Yet, the x-intercepts remain constant.

The slopes of Lineweaver-Burk plots are different, yet the yintercepts are the same.

The slopes of Lineweaver-Burk plots are the same and both intercepts are different.

Example Graph 1: Representation of the possible types of enzyme inhibition

You have created two inhibitors (A and B) and done experiments to compare their effectiveness.

For each set of data collected: a. Create a Lineweaver-Burk plot of the data.

.

b. Calculate VMax and KM of the uninhibited and inhibited line.

A: inhibited Vmax = 5000 mol/L•min; inhibited Km = 0.01 mol ; Non-inhibited Vmax = 54.0 mol/L•min; non-inhibited Km = 2.16e − 5 mol B: inhibited Vmax = 33.7 mol/L•min; inhibited Km = 1.67e − 4 mol ; Non-inhibited Vmax = 54.0 mol/L•min; non-inhibited Km = 2.16e − 5 mol

c. Determine the type of Inhibition A = Competitive; B = Non-competitive d. Which is the better inhibitor? Why? B is the better inhibitor because the Vmax is reduced and the Km is also lowered.

Typical Inhibitor Data: Inhibitor A Without Inhibitor So (mol/L) Vo (mol/L•min)

With Inhibitor Vo (mol/L•min)

0.000005

9.1

2.8

0.000007

13.2

5.1

0.000012

21.2

10

0.000022

26.8

15.1

0.000032

32.5

21.3

0.000062

37.3

28.9

0.000092

39.2

32.5

0.000122

39.8

32.78

.

Typical Inhibitor Data: Inhibitor B Without Inhibitor So (mol/L) Vo (mol/L.•min)

With Inhibitor Vo (mol/L•min)

0.000005

9.1

0.8

0.000007

13.2

1.6

0.000012

21.2

3.2

0.000022

26.8

4

0.000032

32.5

5.5

0.000062

37.3

6.4

0.000092

39.2

6.8

0.000122

39.8

7.2

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Experiment 1: Conservation of Matter Instructor Notes and Lab Preparation Chemicals and Equipment: 3.0 M HCl in 1-Liter bottles with pour spout Bottle of anhydrous copper (II) sulfate Weighing boats Bottle of zinc metal (mossy) 50 mL volumetric flasks with stopper 125 mL flasks with side arm #4 black rubber stoppers *** DO NOT USE CORK *** 1 bag of 7-inch balloons 1 box of calipers Long tweezers Disposable pipets Parafilm 1 box of baking soda for acid spills Setup:    

Make sure there are Kimwipes near the balance(s).

Set out equipment and chemicals.

Place 1 bottle of HCl on each lab bench.

Place 50 mL volumetric flasks on each lab bench.

Preparation: 3.0 M HCl – Use 247.93 mL of concentrated HCl per 1 L of solution.

Prepare a large carboy: To prepare 12 L of 3.0 M HCl: Fill a carboy with approximately 5 L of water. Add 2.98 L of concentrated HCl (in the hood). Then dilute to 12L with D.I. water. (Supply for 20 classes sections @ 24 students) Mix well.

To prepare 6 L of 3.0 M HCl: Fill a carboy with approximately 3 L of water. Add 1.49 L of concentrated HCl (in the hood). Then dilute to 6 L with D.I. water. (Supply for 10 classes sections @ 24 students) Mix well.

.

Instruction Notes: The Conservation of Matter experiment was originally developed for our non-majors course. For that course we prepare the copper (II) sulfate solution for the students. When we adapted it for the majors we decided to have them prepare their own solution and introduce the use of volumetric glassware. If you feel your students are not ready to mix their own solution, this is an easy change that both shortens and simplifies the experiment. The experiment is very visual as both the color of the solution changes dramatically in a short period of time and the balloon is inflated to a large size. This leaves no doubt that a reaction has taken place.

Points to stress: Students are working with 3.0 M HCl. This is strong acid and it should be stressed repeatedly that all safety precautions must be taken to avoid burns. All spills of any size should be cleaned up by lab personnel, not the students. I provide my teaching assistants with baking soda and gloves to handle the clean-up.

When the students are preparing the solution, they should weigh the CuSO4 in a weigh boat and then use a small amount of HCl to assist in the transfer of the salt to the volumetric flask. CuSO4 is a “sticky, wet” salt and thus clumps, making dry transfer difficult (the neck of the 50 mL volumetric flask is small). You can also use paper funnels but the “stickiness” of the salt may cause a loss of mass there as well. We have found that using the HCl to aid in the transfer reduces the loss of mass best and also keeps the lab benches much cleaner.

Potential Problems: As mentioned above, the necks of the volumetric flasks are fairly small so students should be shown the proper way to transfer the salt so that they don’t make a huge mess on the bench tops and, more importantly, don’t lose mass of salt.

The only other problem that sometimes occurs is a hole in a balloon or a balloon that is not securely fastened to the side-arm of the flask. Have your students blow the balloon up once before attaching it to the side-arm. This will allow them to determine if the balloon is air tight and it also stretches the balloon so that the resistance is lowered when the reaction starts to fill it with hydrogen gas.

You might note tweezers in the list of items needed for this lab. Because our lab balances are only capable of weighing up to ~100g of material, at the end of the lab we have to remove the pellets from the solution and weigh them separately using a tared weigh boat. If you have the same difficulty, tweezers work well as long as the students are gentle. The removal and separate measurement also reinforces the notion that the total mass of the products is what is important; not that they all remain in the same state or container.

.

Pre-Lab Key

Experiment 1 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Based on your reading of the laboratory Introduction and Background, prepare a hypothesis regarding what results you expect from your experiment. Be detailed.

This answer will vary based on the student’s interpretation of the reason for the experiment but should contain a reference to both the concepts of conservation of matter and also the techniques that will be employed to explore the concepts.

2) Calculate the number of moles or grams in each of the following quantities of salt.

5.233 g of Cu(NO3)2 · 6H2O =

________0.01770 mol____

7.77 g of zinc nitrate trihydrate =

________0.0319 mol_____

6.215×10-2 mol of Na2SO4 · 10H2O =

________20.02 g________

2.0×10-2 mol of zinc nitrate hexahydrate =

________6.0 g_________

3) You would like to make up 100 mL of a 0.23 M solution of Na2SO4 · 10H2O. How many moles of the salt do you need? How many grams of the salt do you need?

0.023 mol .

You have a 2.5 M solution of Cu(NO3)2 · 6H2O. What volume of the solution (in mL) must you measure in order to have 0.50 moles of the salt?

200 mL You have a 0.75 M solution of ZnSO4 · 7H2O. What volume of the solution (in mL) must you measure in order to have 32.5 grams of the salt?

Mol amount = 0.11 mol; Vol = 151 mL 4) In your laboratory you will be required to make a solution of 0.25 M CuSO4 · 5H2O in a 25.00 mL volumetric flask. This exercise will walk you through the steps involved. After completion, write down the values for your answers in your notebook and take them to the laboratory to use in preparing your solution. a. What is the formula weight of CuSO4 · 5H2O? 249.684 g/mol b. How many moles of CuSO4 · 5H2O are required to make 25.00 mL of a 0.25 M solution?

0.010 mol c. What mass of CuSO4 · 5H2O will be needed to make this solution?

1.56 g d. You will add the salt to a weighing boat, which you determine has a mass of 1.9785 g. What will be the reading on the balance when you have put sufficient CuSO4·5H2O in the boat?

3.54 g

.

Post-Lab Key

Experiment 1 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Based on your experimental results, does the Law of Conservation of Mass apply to this experiment? Explain.

The answer here will depend on the student’s results. If they have a high value of recovery they should indicate that the Conservation of Mass does apply. If they get a lower value of recovery it will most likely lead them to assume that the conservation of mass is not upheld.

2) If the balloon had expanded more how would the mass of H2 and the mass percent recovered have changed?

Both the mass obtained and the percent recovery would increase.

3) How would increasing the molar concentration of CuSO4 have affected this reaction?

Assuming that there was an excess of zinc, the increase in copper sulfate concentration would lead to a higher product amount but the conservation of mass should stay relatively constant.

4) Zinc metal reacts with acid solution to produce hydrogen gas and zinc ion as follows: Zn(s) + 2H+(aq)  Zn2+(aq) + H2(g) You add 6.825 grams of Zn metal to 29.0 mL of 1.70 M HCl and set up the flask to

.

catch the hydrogen in a balloon as in the following picture: How many moles of Zn have you added to the flask?

0.1044 mol How many moles of H+ are in the solution?

0.0493 mol How many moles of Zn will have reacted with the H+ when the reaction is complete?

0.02465 mol How many moles of Zn will remain unreacted?

0.07975 mol What mass of Zn will remain unreacted?

5.214 g How many moles of H2 gas will be produced?

0.02465 mol What mass of H2 gas will be produced?

0.04969 g Assuming that one mole of H2 gas would occupy a volume of 22 L at the temperature and pressure in the balloon, what would be the volume of the balloon?

0.5423 L

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Experiment 10: Gas Laws and Air Bags Instructor Notes and Lab Preparation Chemicals and Equipment: 1 bottle of 6 M acetic acid 1 bottle of solid sodium bicarbonate 1 box each of 3 different types of zipper-type plastic bags: sandwich, freezer, storage Small weighing boats

Preparation: 6 M acetic acid—Fill the large carboy labeled 6 M acetic acid partially with DI water. In the hood, add 6.90 L of concentrated acetic acid. Dilute to the 20 L mark with DI water and mix thoroughly.

Instruction Notes: The concept of this lab is very simple and yet it combines gas law, stoichiometry, and limiting and excess reagents. The lab is short, normally taking no more than 1 hour to complete. A successful airbag should have a full pillow-like appearance and there should be no particulate residue from the bicarbonate left unreacted. Although we require that the bags be tested in the hood, the reaction is not really fast enough to “explode.”

The heavier zipper-style bags with a double zipper work better than the lighter single zipper bags. They are slightly more expensive but are far less likely to leak. To save on expense, the bags can be used more than once as long as they are well rinsed and dried in between classes.

Points to Stress: The comparison of the “airbag” the students are building in the laboratory to the ones found in cars is a very valid one. The requirement for an exact calculation of the materials needed to completely fill but not overfill the bag is an important point to make.

Potential Problems: The only problem in the experiment occurs when the bag leaks. The reason we started determining the volume of the bags by filling them with water instead of simply converting the measurements given on the box is to test the bag for leaks, i.e., if the water leaks so will the gas.

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Pre-Lab Key

Experiment 10 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Airbags in cars are devices that inflate very rapidly (in about 30 milliseconds) when there is a collision, creating a cushion that deflates slowly and cushions the impact of the collision on the driver or the passenger. The bag is inflated by a gas which is generated by the rapid heating of sodium azide (NaN3) to about 300 °C, causing it to decompose into sodium metal and nitrogen gas.

Complete and balance the equation for this reaction, indicating the states of the reactant (NaN3(s)) and products (Na(s) and N2(g)).

2NaN3(s) 2Na(s) + 3N2(g) Enough nitrogen must be generated in the bag to create a total pressure of 3.50 atm, which then drops as the bag slowly deflates. Assuming the volume of the bag is 65.0 L and the temperature in the car is 28 °C, calculate the molar quantity of N2 that must be generated.

9.21 mol What molar quantity of sodium azide is required to generate this much nitrogen?

6.14 mol What mass of sodium azide is required?

400 g

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The sodium metal produced in the reaction is highly reactive and potentially explosive, so it is removed in a reaction with excess KNO3(s) that produces K2O(s) and Na2O(s), and additional nitrogen. Write a balanced equation for this process, showing the states of the reactants and products.

10 Na(s) + 2 KNO3(s)  K2O(s) + 5 Na2O(s) + N2(g) 2) In this laboratory experiment you are to generate sufficient carbon dioxide to inflate a plastic ziploc bag. You first must carry out some calculations to determine the quantity of reagents needed.

Carbon dioxide will be generated by reacting solid sodium bicarbonate (NaHCO3) with a solution of acetic acid (CH3COOH) as follows: NaHCO3(s) + CH3COOH(aq)  H2CO3(aq) + CH3COO−(aq) + Na+(aq) H2CO3(aq)  H2O(l) + CO2(g) You measure the plastic ziploc bag dimensions to be 6.90 inches by 7.15 inches.

Assuming the average thickness of the bag when inflated will be 1.45 inches, what do you calculate the volume of the bag to be?

71.5 in3 = 1172 mL Of course, this volume is only approximate as the width of the inflated bag is not uniform. To get a better value for the volume, you fill the bag with water, then measure the volume of the water to be 1174.5 mL. Now you measure room temperature to be 24 °C and the barometer on the wall gives a reading of 753 torr for the atmospheric pressure. Using the ideal gas law, calculate the number of moles of CO2 that will be required to fill the bag at this temperature and pressure.

0.0477 mol Now inspect the reactions above generating CO2 and calculate the molar quantity of NaHCO3 and CH3COOH needed.

Quantity of NaHCO3: 0.0477 mol Quantity of CH3COOH: 0.0477 mol What mass of NaHCO3 will be needed? 0.401 g You have a solution of 5.00 M acetic acid. What volume of this solution will be needed?

9.54 mL

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Post-Lab Key

Experiment 10 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Is the gas produced in this experiment really an “ideal” gas? Explain.

An ideal gas is one that completely obeys the ideal gas law. Most gases deviate in some fashion from the law in its absolute form due to intermolecular forces, and thus the gas we are working with is most likely not an “ideal” gas.

2) The gas production in this experiment is fairly slow, but the production in a car airbag is quite rapid. What is the reason for the difference in the two reactions?

The sodium azide in an airbag is ignited and thus a combustion reaction takes place. The input of energy creates a much faster reaction than the simple mixing that takes place during our experiment.

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Experiment 11: Intermolecular Forces and the Triple Point of CO2 Instructor Notes and Lab Preparation Chemicals and Equipment: Liquids: 1 liter bottles of each, labeled only as unknowns A Methanol B Acetone C Water D 2-propanol E Hexane (1) box of small (4.25”) 1 m filter paper (1) bag of small rubber bands Buret or ring stands and clamps Thermometers glass, preferably non-Hg (not digital temperature probes) Clear plastic cups Thick-stemmed pipettes Thin-stemmed pipettes Pliers Bottle of blue food coloring Metric rulers Hot glue gun with glue sticks Bucket of dry ice ~ 1–1.5 lbs for a class of 28 students String

Preparation: Place equipment and chemicals on front lab bench.

Place 5 waste containers in the hood, labeled A, B, C, D, and E Waste in jars C and D can be disposed of in the sink drain with running water. Waste from jars A, B, and E goes into an organic waste jar.

Instruction Notes: This is currently the first wet lab (emphasis on wet) of our second semester of general chemistry at Florida State University. It is a good lab to start with because it requires a minimal amount of .

lab skills (no new techniques) and the “exploding” cups of water refresh the students’ memories regarding the need for goggles in the lab.

The first part of the experiment is about as exciting as watching paint dry, but does reinforce the concepts of intermolecular forces and their effect on physical properties quite effectively. The second part of the experiment dealing with the triple point of CO2 is a student favorite (mostly because of the exploding cups of water).

The setup for this experiment is easy since there are no solutions to make and the equipment is pretty standard. I have found that the buret stands and clamps are best for supporting the thermometers. It also allows simultaneous determination of the temperature change for two solvents at one time. Students can work as groups or benches to determine the T values for the solvent unknowns. This saves on chemical cost. It is advisable to have the two parts of the experiment running at opposite ends of the bench or table due to the possible splash of water.

Also, make sure the students keep their lab notebooks and other personal items clear of the “impact” area. I generally have students work in teams of four with two students working on the first part of the lab while the other two complete part two. At some point, the pairs switch to complete all of the experiment. This also cuts the need for thermometers and pliers in half.

The dry ice needed for the experiment is not as great as it might seem. One or two ~2 cm pellets are sufficient to fill the pipette bulb 1/3 full. We use about 30 pounds during a week for 20 sections of ~28 students. We collect a large amount (~10 lbs at a time) and store it all day in a sealed cooler for easy access.

I suggest purchasing glue guns and pliers from a store such as Walmart or K-mart. The mini glue guns and glue sticks are generally less than $5 and you only need 1 or 2 for a class of 28 students. We actually purchased the pliers at a local flea market for $1 per pair.

Points to stress: The correlation between the strength of the intermolecular forces and the magnitude of the observed change in temperature is a difficult concept for the students to grasp. I have done the best I can to explain it in the background, but a discussion in lab is probably wise.

Students often rush through the observations of the CO2 phase change, but it should be mentioned that very few people get the opportunity to observe liquid CO2.

Potential Problems: There are no known difficulties with either part of this experiment.

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Pre-Lab Key

Experiment 11 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

Copy the information gathered in part 1 and 2 of this pre-lab into your lab notebook for later use.

1) For the following liquids: methanol, isopropanol, n-hexane, acetone and water, (a) write the molecular formula; (b) draw a Lewis structure for each compound; and (c) calculate the molecular weight of each compound.

methanol : CH3OH

MW = 32 g/mol H H

C

O

H

H

2-propanol : (CH3)2CHOH

MW = 60 g/mol H H

C

H

H

C

O

H

C

H

H

.

H

n-hexane : C6H14 MW = 86 g/mol H

H

H

H

H

H

H

C

C

C

C

C

C

H

H

H

H

H

H

H

acetone : C3H4O MW = 58 g/mol O H

H

C C

C

H

water : H2O

H

H

H

MW = 18 g/mol O H

H

2) The relative strength of the intermolecular forces depends on the size (i.e., MW), and structure of the molecules, and include in rough order of strength: ion-ion, ion-dipole, hydrogen bonding, dipole-dipole (polar), dipole-induced dipole, and induced dipole-induced dipole (dispersion) forces. The latter increases with molecular weight and can sometimes become the most predominant force. There are no ions in these liquids, and all the molecules are alike, leaving primarily hydrogen bonding, dipole-dipole, and disperson forces to consider.

What are the strongest intermolecular forces in acetone? Dipole-Dipole What are the strongest intermolecular forces in methanol? H-Bonding What are the strongest intermolecular forces in 2-propanol? H-Bonding What are the strongest intermolecular forces in water? H-Bonding What are the strongest intermolecular forces in hexane? Dispersion 3) The figure at the right shows your experimental setup for estimating the triple point of CO2. Carefully mark off your capillary tube in ten 1-cm segments, and note the position of the drop of dye solution when the capillary is inserted in the open plastic tube. When the top of the plastic tube is sealed, the pressure inside will increase as the CO2 sublimes, and the drop of dye will be forced up the capillary, compressing the air inside it. Note the position of the drop of dye when liquid CO2 starts to form. The pressure inside the capillary should then correspond to the pressure at the triple point of CO2.

You can calculate the pressure from a modification of Boyle's law, P1V1=P2V2.

The volume of the capillary V = hA where h is the height and A is the cross.

sectional area, which is constant. Boyle's law therefore reduces to P1h1=P2h2 where P1 is the pressure in the open tube, h1 is the distance of the drop from the top of the capillary in the open tube, h2 is the distance of the drop from the top of the capillary in the closed tube, and P2 is the pressure in the closed tube—the pressure at the triple point.

What are the heights of the air columns in the two capillaries? (Measure from the top of the dye droplet.)

Height of column in the open tube (h1): 10 cm Height of column in the closed tube (h2): 2 cm Calculate the pressure in the closed tube, assuming the pressure in the open tube is 1.0 atm.

Pressure in the closed tube: 4.5 atm

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Post-Lab Key

Experiment 11 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) How is the change in temperature we observe as the various solvents evaporate related to their Hvap? Explain this relationship.

The Hvap is a physical quantity that measures the amount of energy required to vaporize the liquid. Evaporation is an endothermic process that results in a drop in temperature of the surroundings. The compound requiring the least amount of energy (smallest Hvap) and thus the weakest intermolecular forces will exhibit the largest temperature drop. This occurs because as the molecules evaporate they carry heat away from the thermometer, which registers each loss as a drop in temperature. Molecules with stronger intermolecular forces require a larger amount of energy to accumulate before they can evaporate so their removal of heat from the thermometer is a much slower process. The length of time between evaporation events allows the thermometer to recover heat from the surrounding air and thus we do not observe as great a drop in temperature for these molecules.

2) In our use of the microgauge to determine the triple point pressure of CO2, we only measure the height change (distance the dye moves) in the gauge, not the actual volume. Why does this approximation not ruin our results?

From the rules of multiplication and division, we know that any equivalent value found on two sides of an equality can be ignored in the calculation as it cancels out. Since the diameter of the gauge is a constant, it is an equivalent value that appears on both sides of the P1V1 = P2V2 equation and thus it cancels out and can be left out of the overall calculation. .

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Experiment 12: The Purification of Water Instructor Notes and Lab Preparation Chemicals and Equipment: pH probes and meters (students work individually, but share a pH meter) pH 4 buffer pH 7 buffer dropping bottles of 3 M HCl 1 L bottles of 10 ppm phosphate stock solution dropping bottles of ammonium molybdate reagent dropping bottle of stannous chloride dropping bottles of sodium hypochlorite reagent 1 jar of lime powder (CaO) 1 jar of aluminum sulfate powder (Al2(SO4)3) 1 jar of sand 1 jar of activated charcoal (fish tank rocks) 1 box of 0.1 m filter paper Buchner funnels 250 mL filtering flasks vacuum tubing ring stands clamps spectrometers cuvettes 10 mL glass pipettes pipette pumps or bulbs 25 mL volumetric flasks

.

Preparation: Ammonium molybdate reagent: Prep for 1 L: In a 1 L volumetric flask, dilute 280 mL of concentrated sulfuric acid by adding it to 400 mL of DI. water. Meanwhile, dissolve 25 g of ammonium molybdate (NH4)6Mo7O24.4H2O in about 150mL of D.I. water. When the acid mixture has cooled down, combine it with the molybdate and dilute the entire solution to 1 L with DI water.

For 100 mL: 2.5 g of ammonium molybdate and 28 mL of concentrated sulfuric acid; prepared as stated.

0.1 M stannous chloride: In a 1 L beaker, dissolve 12.0 g of tin(II) chloride dihydrate in 500 mL of glycerol.

Warm the solution on a hotplate to speed up the process.

10 ppm stock phosphate solution: Dissolve 0.07 g of K2HPO4 in 7 L of DI water.

Sodium hypochlorite reagent: Dissolve 250 mL of non-scented commercial grade (5.25%) bleach in a 1 L volumetric flask and dilute with DI to the mark.

Phenolphthalein indicator: Dissolve 0.5 g of phenolphthalein in 50 mL ethanol and 50 mL of DI water.

Setup: On every bench place: A Spectronic 20® (or similar spectrometer), a 25 mL volumetric flask, (2) pH 4 buffers, (2) pH 7 buffers, 1 box of Kimwipes®.

Place an aqueous metals waste container in the hood. *** Paper towels and Kimwipes need to be replenished often*** ***pH buffers need to be replaced after every two classes***

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Instruction Notes: This is a lab that was created to emulate a typical water quality laboratory and purification plant, as described in Sec 14.13, 4th edition of General Chemistry by McMurry and Fay. It also serves to introduce the spectrophotometer for a purpose, rather than just as an exercise. This is a fulllength lab intended to run for approximately 2 ½ to 3 hours. Students should work in groups for the creation of the standard solutions and when using the spectrophotometer. This will reduce the amount of standard solution and the number of spectrophotometers needed for a class. They may also share a pH meter since its use is limited. pH paper can be used instead if you prefer.

Because this lab has several parts and is longer than other labs, I generally set time goals for completion of each section to minimize student “dawdling.” If they can complete the serial dilutions and initial absorbance measurements by the end of one hour, they will be able to complete the purification process and re-test the phosphate concentration within the 3 hour limit.

Points to stress: The first part of this lab requires the students to go into the field and collect their own water samples. Even though it is stated in the manual, you should re-emphasize not using detergents to clean the sample bottle. We ask that students bring the collection bottles they will be using into the lab a week early in order to use a mild acid wash to clean the bottle(s). This will help them remember the need for sample collection the week before the lab is run. This need for advance preparation which is so typical of actual research is an important point to stress.

There are many different techniques used in this lab that should be carefully presented to the students: 1) the calibration and use of a pH meter; 2) the proper use of a spectrophotometer; 3) the setup and use of a filter funnel system.

True serial dilution is a process that can be invaluable in an actual research lab due to the everrising cost of chemicals, but was very seldom taught in laboratories I had seen previously. As a public university, we are also trying to lower the cost per student for our labs. When we changed the process from a simple creation of solutions to a serial dilution process, we reduced the amount of stock solution needed by 65%. This point should me made to the students as well.

The steps of the purification process are based on the typical process used by water treatment plants.

There are three principal stages in water purification: 1. Primary treatment–Collecting and screening including pumping from rivers and initial storage.

2. Secondary treatment–Removal of fine solids and the majority of contaminants using filters, coagulation, flocculation, and membranes.

3. Tertiary treatment–pH adjustment, carbon treatment to remove taste and smells, disinfection, and temporary storage to allow the disinfecting agent to work.

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The EPA has a nice website: http://www.epa.gov/safewater/kids/watertreatmentplant/ that shows the typical process used to treat waste water, and will demonstrate the similarities between the real process and the procedure the students are completing in the lab.

Potential Problems: We have had some problems with our distilled water blank in combination with the coloring agent, probably due to some contamination in the water.

Students often bring in water from “duck ponds.” Unfortunately, some municipalities treat these duck ponds with antibacterial agents and other chemicals to reduce odor and clear the water.

These chemicals can produce unknown reactions when processed in the lab. (We had one water sample that foamed like “elephant toothpaste” for 20 minutes when we added the CaO and Al powder). Just try to be aware of any water sources in your area that may be “quality” controlled and warn your students away from them.

The only other consistent source of error for the students is in the filtration process. They tend to get in a hurry and swamp the filter paper. It is important to show the students the proper way to use the filtration system.

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Pre-Lab Key

Experiment 12 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Calculate the [PO43-] concentrations corresponding to the following ppm values: Concentration = 0.41 ppm; Concentration = 275 ppm; Concentration = 21 ppm;

[PO43-] = 4.3e-6 M [PO43-] = 2.9e-3 M [PO43-] = 2.2e-4 M

Calculate the ppm values corresponding to the following [PO43-] concentrations: [PO43-] = 4.56 × 10-6 M; Concentration = 0.433 ppm [PO43-] = 3.88 × 10-3 M; Concentration = 369 ppm [PO43-] = 7.66 × 10-5 M; Concentration = 7.28 ppm 2) Use the following table of standard concentrations and absorbances for the phosphate ion to create a calibration graph. What are the slope and intercept of the calibration curve?

Phosphate Analysis What is the concentration of phosphate in a solution that has an absorbance of 0.708 AU? 6.54 ppm One of your samples shows a reading of 1.207 AU, which is beyond the range of the calibration curve. The lab is over and it is too late to redo the calibration over a broader range, or to dilute the sample and determine it a second time. Assume that Beer's Law still holds up to the value of the sample, and calculate the concentration using the equation for the calibration line. 12.0 ppm (Path length = 1.0 cm)

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Concentration (ppm)

Absorbance (AU)

0

0.109

2

0.288

4

0.476

6

0.651

8 10

0.847 0.997

3) Phosphate standards used to generate the calibration curve are usually made by first preparing a stock solution of known concentration, then diluting it to the required concentration.

First prepare a stock solution that is 1.00 x 10-2 M in phosphate ion. How many g of Na3PO4 would you add to a 250 mL volumetric flask to to prepare this solution?

Mass Na3PO4: 0.410 g What is this concentration in ppm of phosphate? (That is as PO43−, not Na3PO4).

PO43− concentration: 950. ppm What volume of this ppm solution would you dilute with water into a 1.000 L volumetric flask to make a solution of 10.0 ppm? 10.5 mL This much of the work will be done for you by the laboratory assistant. You will be given a stock solution of 10.0 ppm phosphate to use in calibrating your phosphate assay. To preserve the ammonium molybdate reagent, you will create the calibrating standards by doing a serial dilution, using an aliquot of each diluted solution to create the next dilution.

After adding 1 mL of ammonium molybdate reagent and 2 drops of stannous chloride to 25 mL of the 10.0 ppm solution to develop the color, use a 20 mL aliquot to dilute with 5 mL of water to make 25 mL of a solution of 8.0 ppm.

Using the dilution equation: V1 × Conc1 = V2 × Conc2 and rearranging: 20mL  25mL

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8 ppm 10 ppm

To create 25 mL of a 6.0 ppm solution, what volume of the 8.0 ppm solution and what volume of water should be mixed? 18.75 mL solution and 6.25 mL water To create 25 mL of a 4.0 ppm solution, what volume of the 6.0 ppm solution and what volume of water should be mixed? 16.67 mL solution and 8.33 mL water To create 25 mL of a 2.0 ppm solution, what volume of the 4.0 ppm solution and what volume of water should be mixed? 12.50 mL solution and 12.50 mL water

4) What is the molar concentration of each of your calibration solutions?

Concentration = 10 ppm; [PO43-] = 1.05E-04 M Concentration = 8 ppm; [PO43-] = 8.42E-05 M Concentration = 6 ppm; [PO43-] = 6.32E-05 M Concentration = 4 ppm; [PO43-] = 4.21E-05 M Concentration = 2 ppm; [PO43-] = 2.11E-05 M

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Post-Lab Key

Experiment 12 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Why should we care about the amount of PO43- found in a natural water source?

High concentrations of PO43- are indicative of human effects on an ecosystem (i.e., pollution). Generally, phosphorous is a limiting reagent in an ecosystem.

In this way, it regulates the amount of growth various species of plants and animals can achieve. Addition of excess phosphorous into the ecosystem can lead to explosive growth (especially in plants), and this overgrowth can endanger the balance of the ecosystem.

1) If you really wanted to be sure your water was drinkable, what further purification step(s) would you take?

Distillation predominantly, but allow for other examples.

2) Blue light has a wavelength between 450–495 nm. Why don't we measure the absorbance of a blue solution within this wavelength range?

The solution appears blue because it transmits light in the blue wavelength region, but it absorbs light at other wavelengths. Therefore very little absorption would be seen within the region of blue light.

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Experiment 13: Kinetics: The Iodine Clock Instructor Notes and Lab Preparation Chemicals and Equipment: 1 bottle of 0.2 M KI 1 bottle of 0.05 M (NH4)2S2O8 1 bottle of 0.00175 M Na2S2O3 Dropping bottles of 3% starch solution Hotplates Cooler of ice 1 box of salt Waste container in hood

Preparation: For an average section of up to 28 students, prepare: 0.2 M KI – 66.40 g of KI in a 2 L volumetric flask, dilute to the mark with DI water 0.05M (NH4)2S2O8 – add 22.82 g in a 2 L volumetric flask, dilute to the mark with DI water 0.00175M Na2S2O3 – add 0.553 g of anhydrous Na2S2O3 to a 2 L volumetric flask, dilute to the mark with DI water 3% Starch – in a 2 L beaker bring 1000 mL of DI water to a boil; weigh 30 g of starch in a large weighing boat. Dissolve in the weighing boat with DI water. Once the water is boiling, SLOWLY add the liquid starch, while stirring. Place a large watch glass on top of beaker.

Continue heating and stirring to dissolve. Be careful not to over boil. If not prepared fresh, keep refrigerated. Check during the week simply by combining a few drops of each reagent and observing the dark color.

Setup: Place a hotplate on each bench.

Place the chemicals on the front counter.

Place the waste container in the hood. ***Chemicals will need to be replenished after every lab*** ***Waste will need to be disposed of after every lab***

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Instruction Notes: The iodine clock experiment has been used off and on at Florida State University for the past 20 years. We go away from it every semester or so to try something new, but we always seem to return to this experiment because of its simplicity and its coverage of both simple kinetics and the Arrhenius Equation.

The experiment itself can be completed in about 1–1 ½ hours. There are no special techniques needed to complete the experiment so instructional time at the front end of lab is minimal. The time that remains available at the end of the lab should be used to help students understand how to process their data. I have always advised my teaching assistants to spend the first few minutes of lab on technique and wait until the end of lab to deal with concepts. Students generally find it easier to understand how to analyze their data once they have collected it.

Points to Stress: The final run of the experiment changes the concentration of the limiting reagent, S2O32-.

Students often assume that this will change the rate of the reaction. It should be pointed out that for this experiment, only variations in temperature change the rate.

Potential Problems: The reaction procedure is very simple and thus we haven’t seen very many problems occur.

Difficulties can arise, however, if the reagents are not fresh enough. Do not make substantial amounts of reagent in advance of the class. Small amounts as described above should be produced instead.

Results can be damaged by contamination of the reagents as well. Make sure students do not pour excess reagent back into the lab supply. I say this repeatedly but I am always catching students who accidentally take too much reagent trying to return it to the reagent bottle. On occasion we have had an entire class’s results become “skewed” and deduced that this type of contamination was the culprit.

Finally, the students need to be careful to use different graduated cylinders to measure each reagent. If they cross-contaminate their cylinders, the reaction may start early and completely destroy the timing of the reaction.

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Pre-Lab Key

Experiment 13 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) For KI, Na2S2O3, and (NH4)2S2O8, calculate the initial concentration (M) based on the table of solution concentrations and volumes given in the procedure section. Copy the completed table into your lab notebook for later use. The shaded columns will be completed in lab.

Run Number

[KI] [Na2S2O3] [(NH4)2S2O8] Time Rate of Temp.

K (M) (M) (M) (sec) Reaction (K) (determine (M/sec) units)

1 2 3 4 5 6 7 8

0.075 0.050 0.025 0.075 0.075 0.075 0.075 0.075

0.00044 0.00044 0.00044 0.00044 0.00044 0.00044 0.00044 0.00022

0.019 0.019 0.019 0.012 0.0062 0.019 0.019 0.019

2) A first order chemical reaction has a rate constant of 9.1 × 101 s-1 at 12 °C and a rate constant of 4.1 × 102 s-1 at 36 °C. Calculate the activation energy, Ea, for the reaction.

4.6E-01 kJ/mol

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3) In the iodine clock reaction, you are measuring the rate of the reaction of iodide ion (I−) with peroxydisulfate ion (S2O82−): S2O82− + 2I−  2SO42− + I2

(Reaction 1)

The rate of the reaction is measured by measuring the amount of I2 formed in a particular time, or Δ[I2]/Δt. The I2 is measured by “titrating” it with a fixed amount of thiosulfate ion (S2O32−) which rapidly converts the I2 formed in Reaction 1 back to iodide ion: I2 + 2S2O32−  2I− + S4O62−

(Reaction 2)

(While Reaction 2 is said to be "fast" with respect to Reaction 1, it actually can proceed no faster than the rate at which I2 is produced, so its actual rate is controlled by the rate of Reaction 1.)

When the S2O32− is completely used up, then I2 begins to accumulate. It combines with starch to form a dark blue complex, so the “end-point” of the “titration” occurs when the solution turns blue. The time it takes for the blue color to appear is called the period of the clock, or Δt, the time it takes for Reaction 2 to go to completion. The rate of S2O32− disappearance is therefore given by its initial concentration divided by the clock period, or [S2O32−]/Δt. Because the rate of disappearance of I2 in Reaction 2 is controlled by its rate of appearance in Reaction 1, the rate of both reactions can be determined from the value of ½ × [S2O32−]/Δt. (Two S2O32− ions are consumed for every I2 consumed.)

(Actually, the I2 in solution combines with iodide ion to form triiodide ion, I3−, which is the species that forms the blue color with starch. That step is not shown to keep the equations simple, but adding it would not change the conclusions.)

The following graphs illustrate what is happening to the concentration of each of the species involved in these two reactions.

Show you understand the reaction by identifying the species associated with each of the lines in the graphs.

Initial conditions: [S2O82−] = [I−] = 0.055 M; [S2O32−] = 0.0015 M

[SO42−] =

____F_______

[I−] =

____E_______

[I2] =

____D_______

[S2O32−] = ____B_______ [S4O62−] = ____A_______ [S2O82−] = ____C_______

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Post-Lab Key

Experiment 13 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Which of the eight values of k should be the same (assuming no experimental error)? 1-5 and 8 2) Explain what factors affect the rate law constant.

Temperature, concentrations of reactants, catalysts, surface area of a solid reactant, pressure of gaseous reactants, or products all effect the rate of a reaction. The rate law constant is only constant at a particular set of conditions.

3) Which reaction proceeds faster, one with a large Ea or a small Ea? Explain why.

A small Ea means less energy is required for the reaction to start so the process can move faster.

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Experiment 14: Photometric Determination of an Equilibrium Constant Instructor Notes and Lab Preparation Chemicals and Equipment: 0.200 M Fe(NO3)3 acidified (0.5 M nitric acid) 0.002 M Fe(NO3)3 acidified (0.5 M nitric acid) 0.002 M KSCN 0.5 M nitric acid 10 mL volumetric pipettes Bulbs for transfer pipettes Cuvettes Spectrophotometers Kimwipes Setup:   

Place one Spectrophotometer on each lab bench Place a waste container in the hood Place remaining chemicals and supplies on the front counter

*** Make sure to replenish Kimwipes often*** Preparation: 0.200 M Fe(NO3)3 – To prepare 7 liters in amber carboy: Dissolve 565.60 g of Fe(NO3)3.H2O in 0.5 M HNO3 and dilute to the mark with 0.5 M nitric acid. Fill 1 liter amber bottles with the reagent.

0.002 M Fe(NO3)3 – To prepare 7 liters in amber carboy: Place 70 mL of 0.200 M Fe(NO3)3 and dilute to the mark with 0.5 M nitric acid. Fill 1 liter amber bottles with the reagent. *** LIGHT SENSITIVE*** 0.002 M KSCN – To prepare 7 liters, dissolve 1.36 g of KSCN in DI water and dilute to the mark.

Instruction Notes: This experiment is extremely simple to execute and takes approximately 1 ½ hours to complete.

The time left at the end of the lab can be put to good use going over the calculations required to complete the lab report. The experiment again uses a spectrophotometer and requires the production of a graph to determine the extinction coefficient. This process has been done several

.

times in previous experiments, so students should require minimal instruction to complete this portion of the report. The calculation of the equilibrium constant, however, is best completed using a spreadsheet program such as Excel. The calculations are not difficult but their repetitive nature means they can be tedious.

Points to stress: Students will be using Mohr pipettes to make accurate concentrations of solution. This is the only technique being introduced in the experiment. Whether using bulbs or pumps, the proper technique should be demonstrated. The correct determination of the equilibrium constant depends on the accuracy of the solution concentrations.

Potential Problems: For obvious reasons, contamination of the solutions is one of the most serious sources of error.

Luckily, the solution changes color if contamination occurs and can therefore be corrected.

Another problem we have had is students not rinsing the cuvette completely between measurements. This can lead to skewed results.

The solutions themselves are light/heat sensitive. If the iron(III) nitrate solution loses its color, the results are no longer reliable. The solution should be kept in an amber bottle and made fresh as often as possible or practical.

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Pre-Lab Key

Experiment 14 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) You would like to determine the equilibrium constant for the binding of a ligand to a metal ion: M3+ + L− <=> ML2+ where K = [([ML2+])/([M3+][L-])]. The ML2+ complex has a unique absorbance in the visible region, so its concentration at equilibrium can be determined by measuring the absorbance of the solution at an appropriate wavelength. The concentrations of M3+ and L− can be determined from their initial concentrations and the concentration change in going to equilibrium using an ICE table. The following experiment is carried out first in order to determine the extinction coefficient of the complex at the wavelength chosen.

Calculate the complex concentration in each solution, and then its extinction coefficient.

0.200 M Solution Metal Ion (mL)

0.1 M HNO3 (mL)

0.0011 M Ligand Ion (mL)

Solution Concentration (M)

Absorbance

Extinction Coefficient (M-1cm-1)

A

5.0

4.0

1.0

0.00011

0.113

1027

B

5.0

3.0

2.0

0.00022

0.227

1032

C

5.0

2.0

3.0

0.00033

0.337

1021

D

5.0

1.0

4.0

0.00044

0.431

980

E

5.0

0.0

5.0

0.00055

0.562

1022

Plot the absorbance versus the solution concentration and determine the slope and intercept of the resulting line.

.

2) Now you will determine the equilibrium constant, K, at several starting concentrations of M3+ and L−. The following table summarizes your measurements.

0.0023 M Solution Metal Ion (mL)

[M3+]init

0.0011 M Ligand Ion (mL)

[L-]init Absorbance

F

3.0

0.00069

7.0

0.00077

0.286

G

4.0

0.00092

6.0

0.00066

0.312

H

5.0

0.00115

5.0

0.00055

0.308

I

6.0

0.00138

4.0

0.00044

0.276

J

7.0

0.00161

3.0

0.00033

0.221

Using an ICE table, calculate the equilibrium concentrations of ML2+, M3+ , L− and K for each of the five solutions.

.

Solution

[ML2+]=x

[M3+]equil

[L-]equil

F

0.000285481

0.00040

0.000485

1456.556

G

0.000311434 0.0006086

0.000349

1468.157

H

0.000307441 0.0008426

0.000243

1504.334

I

0.000275499 0.0011045

0.000165

1516.302

J

0.000220599 0.0013894

0.000109

1451.29

What value of K do you get for Solution H? 1504.334 What average value of K do you get? 1479.328

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K

Post-Lab Key

Experiment 14 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) How constant were your Kc values at room temperature? Explain any variation.

The answer to this question will depend on the student’s data. Their explanation should include references to any errors committed in measurements or calculation.

2) See if you can find literature and/or Internet references for the equilibrium constant for this equilibrium. Please cite the reference(s). How does your result compare with the one(s) you cited? Perhaps you will even find a similar experiment, with sample results, posted out on the Web.

The literature value most often found is around 138.

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Experiment 15: Titration of 7-Up© Instructor Notes and Lab Preparation Chemicals and Equipment: pH probes and meters Bottles of pH 4 buffer Bottles of pH 7 buffer Ring stands Clamps 50 mL burets 1 L bottle of 0.01 M NaOH 1 L bottle of 0.01 M HCL 2 L bottle of 7-Up 1 L bottle of 0.01 M phosphoric acid Dropping bottles of phenolphthalein Preparation: 0.05 g of phenolphthalein in 50 mL of ethanol and 50 mL of DI water.

0.01 M HCL and 0.01 M ***NaOH*** are purchased pre-made. The solutions are stored in sealed plastic carboys.

7-Up – pour into a 2 L beaker and boil to remove carbonation. Allow to cool and return to container.

0.01 M ***Phosphoric Acid*** – to make 7 liters in the large carboy: fill carboy with some DI water, add 4.73 mL of concentrated phosphoric acid, dilute to the mark. ***Do NOT label with Molarity***

Instruction Notes: This is a simple titration experiment which gives students experience titrating weak and strong, mono- and poly-protic acids, and analyzing their plots. We used to spend weeks drying and titrating KHP, but the students seemed to lose focus regarding why we were completing the experiment around week two. By combining the standardization of NaOH along with the titration of different acids as unknowns, there is an obvious and practical reason for completing each and every part of the experiment. The experiment requires approximately 2 ½ hours to

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complete all four titrations. If students add NaOH in no more than 1 mL increments at a time, the graphs are very nice. Each titration will require no more than ~ 45 mL to complete, so if the students refill the buret before beginning each titration, they will avoid the need to refill in the middle of a titration.

Points to stress: During the first titration, the students are asked to move through the process fairly quickly.

Before starting the second titration, they should develop a strategy as to where they will “slow down” to produce the most accurate plot of the equivalence point.

Once the concentration of the NaOH is determined during the first two titrations, the students should be able to determine the maximum pH that can be reached (about 12). This will prevent students from futilely continuing a titration trying to reach a higher pH. Potential Errors: We have not had any significant problems with this experiment.

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Pre-Lab Key

Experiment 15 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1)

To standardize your NaOH solution, you titrate 25.0 mL of 0.0822 M HCl solution, measuring the pH after each small addition of NaOH. Following is the graph you obtain.

The equivalence point occurs where there is a steep rise in pH.

What molar quantity of HCl is being titrated?

0.002 mol What molar quantity of NaOH has been added at the equivalence point? 0.002

mol What volume of NaOH solution has been added at the equivalence point? 23.9 mL Expanded view of the equivalence point: What is the concentration of the NaOH solution?

0.0084M

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2) A triprotic acid requires three moles of base to neutralize it, and the protons are removed one at a time as follows: H3A + OH− H2A− + H2O H2A− + OH−  HA2− + H2O HA2− + OH−  A3− + H2O A titration curve, which plots the pH of the acid solution versus quantity of base that is added, can reveal several things about the nature of the acid. First of all, some or all of the equivalence points can be observed in the graph as places where there is a maximum in the slope of the curve.

As an example, you have a 40.0 mL solution of a triprotic acid, H3A, with a concentration of 0.0588 M. You titrate it with a 0.350 M solution of NaOH. How many millimoles of H3A are you titrating? (Give units.) 2.35 mmol How many millimoles of NaOH are required to complete the titration? (Give units.)

7.056 mmol What volume of NaOH will be needed to completely titrate the acid? 20.16 mL What volume of NaOH will be needed to reach the first equivalence point? 6.72 mL What volume of NaOH will be needed to reach the second equivalence point? 13.44 mL 3) Citric acid is a weak organic acid found in citrus fruits and some soft drinks. It is an important intermediate in the citric acid cycle and therefore occurs in the metabolism of almost all living things. It is a triprotic acid, requiring three moles of OH− to remove all the acidic protons.

Titration of citric acid, as with all triprotic acids, occurs in a stepwise fashion in which one proton is removed at a time. The stepwise dissociation can therefore be represented by three equations, each with an acid dissociation equilibrium constant.

H3C6H5O7  H2C6H5O6− + H+

(Ka1)

H2C6H5O7−  HC6H5O62− + H+

(Ka2)

HC6H5O72−  C6H5O63− + H+

(Ka3)

The following graph shows how the concentration of each species changes when NaOH is added to a solution containing 100 mmoles of citric acid.

.

What is the molecular weight of citric acid? 192.124 g/mol What is the mass of citric acid used in this titration? 19.2124 g After addition of how much base does [H2C6H5O7−] = [C6H5O73−]? 200 mmol After addition of how much base does [H2C6H5O7−] = [HC6H5O72−]? 150 mmol After addition of how much base does [H3C6H5O7] = [HC6H5O72−]? 100 mmol

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Post-Lab Key

Experiment 15 Post-Laboratory Question Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Are you able to determine all of the Ka values for citric acid? Why or why not?

Because the first two Ka values of citric acid are so closely spaced they are seldom seen as separate equivalence points. Normally only the first and the third are resolved.

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Experiment 16: Hydrogen Phosphate Buffer Systems Instructor Notes and Lab Preparation Chemicals and Equipment: pH probes and meters Bottles of buffer pH 4 Bottles of buffer pH 7 1 L bottle of 0.5 M NaOH 1 L bottle of 0.5 M HCl 1 bottle of KH2PO4 1 bottle of NaH2PO4*H2O 1 bottle of Na2HPO4*7H2O 1 bottle of Na2HPO4*12H2O Small weighing boats

Preparation: No preparation needed. 0.5 M standard solutions of NaOH and HCl can be purchased.

Instruction Notes: The buffer lab is designed to do three things: 1) give the students more experience preparing a solution, 2) give them more experience calibrating and using a pH meter, and 3) give them lots and lots of practice completing buffer calculations. The experiment itself is very simple to run and requires minimal setup for the instructor. Once the students prepare their buffer, they then test the buffering capacity of the solution by adding small increments of strong acid or base, measuring the pH after each addition. In the lab report, these measurements are compared to theoretical calculations using the same amount of acid or base. The main point of this lab is to give the students plenty of practice with ICE tables and the Henderson-Hasselbalch equation.

Points to Stress: The production of the buffer solution requires careful measurement of solid salts on the balance.

Stressing proper weighing technique here is important. Also, the proper use of stock chemicals should be reviewed before starting the experiment; otherwise contamination is probably going to occur. (Although, no matter how many times I say not to, someone always puts their excess back into the reagent bottle. Arghhh!)

Remind the students to make sure that after they create their buffer it is at room temperature before starting the test procedure. They should also be reminded to take an initial pH of the

.

buffer before starting each test run. This is important so they can get a baseline acid concentration and can thereby determine how well they made their buffer.

In the background it is mentioned that different pKa values should be used rather than the ones listed in most textbooks. This is because the activity coefficient of the buffer must be taken into account to get an accurate pH. We do not cover activity in our freshman classes so I did not go into detail about activity as part of the lab. But if you cover activity in your course it would be appropriate to stress this point as a practical application of activity coefficients.

Last, but not least, the calculations for the lab report require the student to calculate a theoretical pH for each addition of acid or base to compare to the experimental values determined by the pH meter; a percent error is calculated for each as well. Each acid and base addition normally requires 4–6 mL to change by 0.5 pH units for the full strength buffer, and 2–4 mL to change the diluted buffer. This is a lot of calculations and at least one should be demonstrated in lab so the students get a clear understanding of what is expected in their report. It is also important that the students calculate the initial HA and A- concentrations based on the actual (experimental) pH of the buffer. If they use the pH from their pre-lab, and the buffer pH is “off,” then every experimental pH will have a high percent error. Since the experiment, in part, is supposed to convince the students of the validity of Henderson-Hasselbalch, huge percent errors will be counterproductive.

Potential Problems: We have only run into a couple of problems when running this lab: 1) Trying to get the pH of the buffer to come out correctly, our pH values are always too low. At first we assumed student error, but later found that it was in fact our distilled water. Unfortunately, because of the climate in Florida our distilled water becomes fairly acidic (pH ~5.9) very quickly and thus our buffer pH values are always lower than expected. This is why I suggest that students back calculate their starting concentrations of acid and base from the experimental pH rather than just starting from the concentrations in their pre-lab. If they don’t do this calculation, then they should definitely consider the pH of the water they are using as a potential source of error in their comparisons of theoretical and experimental pH. 2) If the students have calculated their mass values incorrectly in their pre-lab this can cause error throughout the lab and report. I generally have my TAs run through the proper way to complete the calculations for at least one of the buffers. You can always select just one pH for an entire class to work with and go through the calculations for this pH. Rotate through different pH values from class to class to cut down on cheating/copying from class to class.

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Pre-Lab Key

Experiment 16 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) You need to prepare 1.000 L (in a volumetric flask) of 0.50 M phosphate buffer, pH 6.21. Use the Henderson-Hasselbalch equation with a value of 6.64 for pK2 to calculate the quantities of K2HPO4 and KH2PO4 you need to add to the flask.

Record the steps of these calculations and use them as a guide for the calculation you will have to make in the laboratory. (Calculate the intermediate values to at least one more significant figure than required.)

What is the [K2HPO4]/[KH2PO4] ratio you will need?

0.372 What concentrations of [K2HPO4] and [KH2PO4] will you need to make the total concentration 0.50 M? (Units required) [K2HPO4] = 0.135 M [KH2PO4] = 0.365 M How many moles of K2HPO4 and KH2PO4 will you need? (Units required) K2HPO4 = 0.135 mol KH2PO4 = 0.365 mol What mass of K2HPO4 and KH2PO4 will you need? (Units required) (FW of K2HPO4 = 174.2 g/mol. FW of KH2PO4 = 136.1 g/mol.) .

K2HPO4 = 23.6 g KH2PO4 = 49.6 g 2) Buffer capacity refers to the amount of acid or base a buffer can “absorb” without a significant pH change. It is governed by the concentrations of the conjugate acid and base forms of the buffer. A 0.5 M buffer will require five times as much acid or base as a 0.1 M buffer for a given pH change. In this problem you begin with a buffer of known pH and concentration, and calculate the new pH after a particular quantity of acid or base is added. In the laboratory you will carry out some stepwise additions of acid or base and measure the resulting pH values.

Starting with 60 mL of 0.50 M phosphate buffer, pH = 6.83, you add 1.7 mL of 1.00 M HCl. Using the Henderson-Hasselbalch equation with a pK2 for phosphate of 6.64, calculate the following values to complete the ICE table: What is the composition of the buffer to begin with, both in terms of the concentration and the molar quantity of the two major phosphate species?

(Units required) Concentration of HPO42−: 0.304 M Concentration of H2PO4−: 0.196 M Molar quantity of HPO42−: 18.2 mmol Molar quantity of H2PO4−: 11.8 mmol What is the molar quantity of H3O+ added as HCl, and the final molar quantity of HPO42− and H2PO4− at equilibrium?

Molar quantity of H3O+ added: 1.70 mmol Molar quantity of HPO42−: 16.5 mmol Molar quantity of H2PO4−: 13.5 mmol What is the new HPO42−/H2PO4− ratio, and the new pH of the solution? (Note: You can use the molar ratio rather than the concentration ratio because both species are in the same volume.)

HPO42−/H2PO4− ratio: 1.22 log(HPO42−/H2PO4−): 0.0872 pH: 6.73 Now take another 60 mL of the 0.50 M pH 6.83 buffer and add 3.7 mL of 1.00 M NaOH.

.

Using steps similar to those above, calculate the new pH of the solution.

pH = 7.07

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Post-Lab Key

Experiment 16 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

Answer the following question: 1) Human blood contains a buffering system. What are the key molecules in this buffer and why is it necessary?

By far the most important buffer for maintaining acid-base balance in the blood is the carbonic-acid-bicarbonate buffer. The simultaneous equilibrium reactions of interest are: The pH of the buffered solution (i.e., the blood) is dependent only on the ratio of the amount of CO2 present in the blood to the amount of HCO3(bicarbonate ion) present in the blood (at a given temperature, so that pK remains constant). This ratio remains relatively constant because the concentrations of both buffer components (HCO3- and CO2) are very large, compared to the amount of H+ added to the blood during normal activities and moderate exercise. When H+ is added to the blood as a result of metabolic processes, the amount of HCO3- (relative to the amount of CO2) decreases; however, the amount of the change is tiny compared to the amount of HCO3present in the blood. This optimal buffering occurs when the pH is within approximately 1 pH unit from the pK value for the buffering system, i.e., when the pH is between 5.1 and 7.1.

All cells in the body continually exchange chemicals (e.g., nutrients, waste products, and ions) with the external fluid surrounding them. This external .

fluid, in turn, exchanges chemicals with the blood being pumped throughout the body. A dominant mode of exchange between these fluids (cellular fluid, external fluid, and blood) is diffusion through membrane channels, due to a concentration gradient associated with the contents of the fluids. Hence, the chemical composition of the blood (and therefore of the external fluid) is extremely important for the cell. If, for instance, the pH of the blood and external fluid is too low (too many H+ ions), then an excess of H+ ions will enter the cell. Maintaining the proper pH is critical for the chemical reactions that occur in the body. In order to maintain the proper chemical composition inside the cells, the chemical composition of the fluids outside the cells must be kept relatively constant. http://www.chemistry.wustl.edu/~edudev/LabTutorials/Buffer/Buffer.html 2) Using stoichiometry, explain why the capacity of a diluted buffer solution is lower than a more concentrated buffer.

Buffers act by reacting mole to mole (stoichiometrically) with added acid or base. At a diluted concentration the number of moles of buffer is reduced and therefore, the number of acid or base molecules capable of reacting with added acid or base is reduced as well. The buffer molecules are therefore, more quickly used up and the system is not as pH stable as a system with a buffer of higher concentration.

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Experiment 17: Entropy, Free Energy, and Chemical Equilibrium Instructor Notes and Lab Preparation Chemicals and Equipment: 50 mL burets Ring stands Buret clamps 1 L bottle of 0.05 M HCl 1 bottle of calcium hydroxide 1 rack of thermometers 2 dropping bottles of bromothymol blue Hotplates or Bunsen burners 1 pack of 0.1 nm filters (Whatman # 5 or 6, or equivalent) Corks Transfer pipettes Bulbs for transfer pipettes Buchner funnels Filter flasks 1 box of rock salt 1 cooler of ice Ice trays Waste container Setup: Place 1 hotplate on each bench.

The rest of the chemicals are placed on the front counter.

Place a waste container in the hood.

Preparation: To make 250 mL of Bromothymol blue: To a 250 mL volumetric flask, add 0.1 g of bromothymol blue powder in 16 mL of 0.01 M NaOH. Dilute to the mark w/ DI water.

Instruction Notes: The maximum in solubility for Ca(OH)2 is around 4 °C. When we first developed this lab we attempted to get students to find the temperature of maximum solubility, but this really required more data points at low temperature than the students were capable of collecting. Now the lab simply requires them to determine the trend in solubility and calculate the thermodynamic properties related to the solubility. .

Points to Stress: Students will be using the filtration flasks again. Because the determination of concentration is very dependent on good technique in this lab, it is extremely important that they filter their solutions carefully and that they don’t attempt to titrate cloudy solutions as this will throw off the results significantly.

Within the calculations they complete for their lab report, the dilution of the saturated solution before titration should be discussed. Because we are approximate in the volume of water added in the dilution, the students get confused as to how to determine the concentration. It should be stressed that we are titrating to determine the number of moles of hydroxide ion that were in the original 10 mL sample.

Potential Problems: As mentioned above, technique is very important. The most significant problem with this lab has come from students titrating cloudy solutions. Remind them that they only need a small amount of filtrate to titrate so they can filter very slowly and carefully without a significant loss of time.

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Pre-Lab Key

Experiment 17 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) You add an excess of Ca(OH)2 to water maintained at a particular temperature,

stir until the solution is saturated, filter, then determine the [OH−] in the solution by titration with acid. Titration of 10.0 mL of the calcium hydroxide solution to the endpoint requires 4.86 mL of 0.070 M HCl solution.

What is the molar quantity of OH− in the 10.0 mL of solution? 0.340 mmol

What are the concentrations of Ca2+ and OH−? 0.0170M and 0.0340 M, respectively.

What is the solubility of Ca(OH)2 under these conditions? 0.126 g/dL What is the Ksp for Ca(OH)2 under these conditions? 1.97E-05 2) You measure the solubility of a salt at four different temperatures and calculate the following Ksp values: Temperature (°C): Ksp:

ΔGosoln (kJ/mol)

15

31

51

70

2.89×10-2 4.11×10-2 6.08×10-2 8.45×10-2

8.5

8.1

7.5

7.0

ΔGosoln for the solubility process is given by the relationship ΔGosoln = -RTlnKsp where R = 8.314 J/(K*mol) and T is the absolute temperature in degrees K. Calculate ΔGosoln (in kJ/mol) at each of these temperatures. (In table)

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As the temperature increases, how does the solubility of the salt change?

Increases According to Le Châtelier's principle, dissolving this salt should therefore be which type of process? Endothermic In other words, what is the sign of ΔHosoln? Positive 3) It is possible to determine both ΔHosoln and ΔSosoln for the solution process from this data. One or both of two graphical methods can be used: Method 1: From the equation: Method 2: Alternatively, rearranging the o o o ΔG soln = ΔH soln-TΔS soln equation relating Ksp and ΔGosoln: a plot of ΔGosoln versus T will give ΔSosoln as the slope, and ΔHosoln as the intercept. and plotting lnKsp versus will give as the slope, and

as the

intercept.

Using one or both of these graphical methods, determine ΔHosoln (in kJ/mol) and ΔSosoln (in J/(K*mol)) for this salt. (Be sure to pay close attention to your units and signs in both your plots and your calculations.) o H soln = 16 kJ/mol o S soln = 26 J/(K*mol)

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Post-Lab Key

Experiment 17 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Why do you have to filter the Ca(OH)2 solution before titrating it?

As you are making a saturated solution, there will be non-dissolved particles of salt left over. Removal of these particles is essential to get an accurate concentration of the dissolved salt.

2) Why don't you have to accurately record the amount of Ca(OH)2(s) used in the saturated solution?

Since you remove all of the non-dissolved salt and titrate the saturated solution to determine its concentration, accurate measurement of the amount of solid salt placed in the solution is a waste of time.

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Experiment 18: Electrochemistry: The Nernst Equation Instructor Notes and Lab Preparation Chemicals and Equipment: Bottle of 0.1 M CuNO3 solution Bottle of 0.1 M Mg(NO3)2 solution Bottle of 0.1 M Fe(NO3)3 solution Bottle of 0.1 M Pb(NO3)2 solution Bottle of 0.1 M Al(NO3)3 solution Bottle of 0.1 M KNO3 1 bottle of each of the metal strips: copper, lead, aluminum, iron, magnesium 1 container of agar powder (inexpensive type, Fishersci #AC40039-5000 or equivalent) 1 container of cotton 1 sheet of sand paper, cut into about 12 x 15 cm pieces 1 cooler with ice U-tubes Thick-stemmed disposable pipettes Hotplates – 2 per lab bench Multi-meters with alligator clips Setup: Place all chemicals and supplies on the front counter.

Place a metals waste container in the hood.

Place 2 hotplates on each lab bench.

Copyright © Pearson Education, Inc.

Preparation: The salt solutions should be prepared as nitrates in a 0.1 M nitric acid solution.

Instruction Notes: Experiment 19 is all about electrochemistry. Each student builds his or her own unit cell and measures the potential. There are actually only five metals: Fe3+, Cu+, Mg2+, Al3+, Pb2+. You should assign each combination to at least two students (if possible) so that the data is reproduced when pooled together. I normally create a table on the front board so that as the students complete the measurements they can report their findings to the class. They only pool the room temperature data. We have recently dropped the concentration to 0.1 M to save on chemical cost. This appears to work just as well as the 1.0 M solutions. We use small pieces (.3 cm x 1 cm strips) of each metal cut from either thin sheet or ~14 gauge wire. You have to really watch the students so they don’t throw them away and that they are returned to the correct container at the end of the lab.

Points to Stress: The students should have calculated the theoretical potentials for all of the cells before coming into class. The experiment allows them to confirm their calculations and should also help them decide which metal should act as the cathode and which should act as the anode. In other words, if they have prepped their lab properly, the potentials should be positive and similar to the ones they have calculated.

The inclusion of the Nernst equation and temperature dependence of cell potential allows this experiment to bridge the gap between chapter 17 and 18, making another use of the thermodynamic properties enthalpy, entropy, and free energy. The importance and relevance of the relationship of these properties to yet another chemical process should be discussed.

Potential Problems: When setting up the u-tubes, make sure the students soak the cotton in the respective solutions before placing them in the tubes. If the cotton is not soaked through they will get bad readings.

We use multi-meters (set to VDC) for this experiment. Occasionally we have had problems with students not setting the meters to VDC. This error is the first thing to check when students report anomalous readings.

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Pre-Lab Key

Experiment 18 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Using literature values for the half-cell reactions of the following metals,

calculate the total cell potentials for each of the 10 possible cells. Assume complete reduction of the metals. Show all your work please. Copy the resulting potentials into your lab notebook.

Metals: Fe3+, Cu+, Mg2+, Al3+, Pb2+ 3Cu+ + Fe ↔ Fe3+ + 3Cu Ecell = 0.521 V + 0.04 V = 0.561 V 2Fe3+ + 3Mg ↔ 2Fe + 3Mg2+ Ecell = - 0.04 V + 2.37 V = 2.33 V Fe3+ + Al ↔ Fe + Al3+ Ecell = - 0.04 V + 1.66 V = 1.62 V 2Fe3+ + 3Pb ↔ 2Fe + 3Pb2+ Ecell = - 0.04 V + 0.126 V = 0.086 V 2Cu+ + Mg ↔ 2Cu + Mg2+ Ecell = 0.521 V + 2.37 V = 2.89 V 3Cu+ + Al ↔ 3Cu + Al3+ Ecell = 0.521 V + 1.66 V = 2.18 V

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2Cu+ + Pb ↔ 2Cu + Pb2+ Ecell = 0.521 V + 0.126 V = 0.647 V 3Mg + 2Al3+ ↔ 3Mg2+ + 2Al Ecell = -1.66 + 2.37 V = 0.71 V Mg + Pb2+ ↔ Mg2+ + Pb Ecell = -0.126 + 2.37 V = 2.24V 2Al + 3Pb2+ ↔ 2Al3+ + 3Pb Ecell = -0.126 + 1.66 V = 1.53 V 2) The standard potential of a galvanic cell (Eocell) can be determined from thermochemical data such as the free energy of formation (ΔGof) of each of the species in the reaction.

For example, given the reaction: Fe3+(aq) + Al(s)  Al3+(aq) + Fe(s)

Species ΔGof (kJ/mol)

Fe3+(aq) Fe(s) Al3+(aq) Al(s) -10.5

0

-481.2

First use the data to the right to calculate the ΔGorxn. -470.7 kJ/mol Then calculate Eocell from ΔGorxn. 1.626 V 3) Following is the shorthand notation for a galvanic cell: Al(s)|Al3+(1M)||Cu+(1M)|Cu(s) Write the overall balanced equation occurring in this cell. (Indicate the physical states of reactants and products, i.e., Al(s), Cu+(aq), etc.)

3Cu+(aq) + Al(s)  Al3+(aq) + 3Cu(s) Identify the anode and cathode of this cell.

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0

Cu(s) Cathode; Al(s) Anode You make the following voltage measurements at four different temperatures: Temperature (°C):

10

35

55

74

Eocell (V):

2.190

2.176

2.165

2.154

ΔGorxn (kJ/mol) -633.9 -629.9 -626.7 -623.5 Calculate the ΔGorxn at each of these temperatures. (In table) Now graph this data using the relationship ΔG = ΔH - TΔS, and determine ΔHorxn (in kJ/mol) and ΔSorxn (in J/(K*mol)) for this process. (This assumes that ΔHorxn and ΔSorxn do not vary with temperature. Remember to use absolute temperature, and pay attention to units and signs in your plot and your calculations.)

ΔHorxn = -679.8 kJ/mol and ΔSorxn = -162.2 J/(K*mol)

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Post-Lab Key

Experiment 18 Post-Laboratory Question Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) What is a Watt? How is it related to the potential of a battery?

A Watt is a unit of energy defined as a Joule per second. Battery power is defined as the amount of energy that can be delivered by a unit of time. The greater the potential of the battery the more energy that can be delivered in a unit of time, thus the Watt is a unit of energy that can be used to describe the potential (power) of a battery.

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Experiment 19: Analysis of a Hydrate Instructor Notes and Lab Preparation Chemicals and Equipment:          

Barium chloride dihydrate Copper(II) sulfate pentahydrate Magnesium sulfate heptahydrate Bunsen burners Ring stands Ring clamps Clay triangle Crucibles + lids Crucible tongs Analytical balances

Setup: Place hydrates in small plastic bottles and mark as unknowns A – C. Set out crucibles, ring stands, clay triangles, and Bunsen burners for each student.

Instruction Notes: This is a simple dehydration experiment. The hydrates given are sufficiently different in water percentage so the identification should be easy barring experimental error.

Potential Problems: From experience, be warned that the hydrate salts tend to get spilled a lot onto the balances.

Students should be monitored closely to make sure they do not leave a mess.

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Pre-Lab Key

Experiment 19 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

Use the information to answer the questions. Show work, include units, and put your answers in the blanks.

1) Geoffrey weighs an empty beaker and finds it to have a mass of 95.83 g. After putting a spoonful of an unknown hydrate into the beaker, he finds that the mass has increased slightly to 99.87 g. He heats the beaker and its contents twice, and finds that the mass has dropped to 97.22 g. Geoffrey is told by his teacher that the molar mass of the anhydrous salt is 74.10 g. a) What mass of hydrate did Geoffrey start with? _____4.04 g_____________________ b) How much water was driven off from the hydrate during the heating process in units of…

1. grams?

____2.65 g____________________ 2. moles?

___ 0.147 mol_________________ c) How much anhydrous salt remained in the beaker in units of: 1. grams?

____ 1.39 g__________________ 2. moles?

_____0.0188 mol______________

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d)

Write down the mole ratio as decimal numbers:

__0.0188____ moles anhydrous salt : __0.147____ moles water Write down the mole ratio as whole numbers:

__1____ moles anhydrous salt : __~8____ moles water e) What is the formula of the hydrate? (use MN to symbolize the anhydrous salt)

_______MN.8H2O_________________ f) Based on Geoffrey’s data, calculate the percentage of water in the sample of hydrate.

_______65.6%_________________ g) Why must Geoffrey heat the sample twice instead of just once?

To completely rid the sample of water.

2) In an experiment, a student dehydrated a hydrate unknown and found that the mass of hydrate was initially 4.13g. After heating, the mass of the desiccant was 3.52g. (a) Calculate the mass percent of water in the hydrate. 14.7% (b)

The table below lists the possible hydrates that could be your unknown.

Complete the table.

Hydrate

Mass Salt (g) Mass Water (g) Mass Percent (%)

CuSO45H2O

249.68

90.05

36

BaCl22H2O

208.236

36.02

17

MgSO47H2O

246.47

126.07

51

(c)

Compare the mass percent values to determine the identity of your unknown.

Mass percent of the unknown is closest to BaCl22H2O

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Post-Lab Key

Experiment 19 Post-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Give some reasons why this experiment for finding the formula of hydrates might not work?

Underheating or overheating, loss of salt in transfer, etc. will give bad mass information that can skew results.

2) If the Bunsen burner left some black soot on the bottom of your crucible, how would this change your answer?

This answer will depend on the answer to the first, but if the student did not already mention a weighing issue it should be discussed here.

3) What are some unavoidable sources of error (like procedural) to explain why you could not achieve 100% accuracy?

One of the most unavoidable sources of error that is not accounted for in this procedure is the humidity in the atmosphere. If the procedure is completed in the Southeast where humidity is high, the results will differ from those completed in the Midwest where humidity is much lower since the hydrates can reabsorb water from the atmosphere as they cool.

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Experiment 2: Atomic Spectra Instructor Notes and Lab Preparation Chemicals and Equipment:

6 M HCl 1 bottle of mossy zinc 1 striker 2 Petri dishes w/opening labeled as unknown 1 & unknown 2 2 coffee cans w/opening for flames Spectroscopes Helium lamp labeled as helium Sodium flame (NaCl) labeled as unknown 1 Lithium (LiCl) flame labeled as unknown 2 Hydrogen lamp labeled as unknown 3 Krypton lamp labeled as unknown 4 Mercury lamp labeled as unknown 5 2 Bunsen burners with vacuum tubing Setup:

On each of the two front lab benches, set up one flame unknown. Place the unknown salt in the Petri dish. Cover and place the Bunsen burner over the opening in the Petri dish. Then cover with the coffee can. When ready to start experiment, place a few pieces of zinc and a few mL of HCl in the Petri dish, cover, and light the Bunsen burner. Place one lamp on each of the remaining four lab benches.

Place the zinc, HCl, striker, and spectroscopes on the front counter.

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Example Coffee can flame set-up and lamp set-up Spectroscopes are inexpensive (< $10) and are available from both Fisher Scientific and Sargent-Welch as well as many other sources. Lamps and bulbs can also be bought from Fischer Scientific and bulbs are available for many other elements as well.

Instruction Notes:

We generally set up five unknown light sources for this lab. Every other semester, I change a few of the sources to keep the experiment fresh. As mentioned above, there are a variety of light bulbs available for the lamps and most of the alkali metals and alkaline earth metals can be vaporized using mossy zinc and HCl to produce colored flames.

Having the students calculate their own wavelengths in the pre-lab and then determine the identity of those elements in the lab really exemplifies the use of spectrometry for identification of materials both near and far.

Points to Stress:

We ask the students to work in pairs to assure that the measurements taken are reproducible. This introduction to the necessity of reproducibility in science should be pointed out. Another important point is the introduction of the need to calibrate instrumentation. To this point, we have assumed that the instruments and equipment being used are in good condition and operating correctly. This experiment is the first time the fallacy of that assumption is explained and .

corrected. The creation of a calibration graph is an important tool that the students will need in future experiments and should not be overlooked as a teaching point in an effort to complete the rest of the lab.

Potential Problems:

The majority of the problems associated with this lab come from inappropriate use of the spectrometers. Make sure the students are aligning the slit in the spectrometer, not the grating, with the light source. Also, too much ambient light will make the lines of the source difficult to see. Keep the light in the lab dim if at all possible.

Unfortunately, you cannot turn off all the lights because then the calibration lines in the spectrometer are invisible as well. Goggles make the reading of the spectrometers more difficult. I generally allow students to work around the lamp sources without the goggles, but because of the acid, have them use the goggles around the flames.

Pre-Lab Key

Experiment 2 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Some of the elements that may be used in this experiment are listed below. For each element write the complete electron configuration. a. Helium 1s2 b. Hydrogen 1s1 c. Mercury 1s 2 2s 2 2p 6 3s 2 3p 6 3d 1 0 4s 2 4p 6 4d 1 0 4f 1 4 5s 2 5p 6 5d 1 0 6s 2 d. Krypton 1s 2 2s 2 2p 6 3s 2 3p 6 3d 1 0 4s 2 4p 6

.

e. Sodium 1s22s22p6 3s1 f. Lithium 1s22s1

2) What makes the wavelength of the N=4 to N=2 transition for hydrogen different from an N=4 to N=2 transition for mercury? How is this difference related to the spectra we observe? Explain.

The spacing between energy levels is unique in each atom based on the number of protons in the nucleus and the shielding that each electron feels. Thus the distances between the shells vary and therefore the wavelength of light emitted or absorbed varies based on this difference in energy and distance.

3) You have just carried out the calibration of your spectroscope's diffraction grating and you get the values shown in the table for the visible emission lines of helium. Plot these values and produce a calibration graph that can be used to correct the values in the table below: Color

Spectrometer Reading

Violet

5.3

Blue

5.7

Green

6.1

Yellow

6.8

Orange

7.7

Red

8.3

Spectrometer reading

Corrected Wavelength (nm)

Color of the Frequency Emitted Light (hz)

6.5 5.4 7.9 4.6 8.5

4) The following atomic line spectra display six of the more prominent lines observed in the visible region for the indicated element. The number by each line represents the relative intensity. A photograph of the actual spectrum of the element may show many more lines, or sometimes fewer lines, depending on the overall intensity of the spectrum.

If two lines are within 1 nm of each other, they are shown as a single line. You can find a complete listing of the lines for each element at the NIST (National Institute of Standards and Technology) Atomic Spectra Database.

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Some spectral lines of Cadmium:

Identify the line in the spectrum of Cd with each of the following characteristics: a.

Light with a photon energy of 2.953 × 10-19 J/photon. F (672.18 nm)

b. Light with a photon energy of 185.8 kJ/mol. E (643.85 nm) c.

Light with a frequency of 5.574 × 1014 s-1. C (537.84 nm)

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Post-Lab Key

Experiment 2 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1. What is the difference between an emission spectrum and an absorption spectrum?

An emission spectrum is produced when an atom is excited by an input of energy, which promotes electrons into higher energy shells. As the atom relaxes back into its normal configuration, energy is released as photons at energies (wavelengths) distinct to the atom’s electronic configuration. An absorption spectrum is produced when light passes through a medium that absorbs photons (radiation) of certain wavelengths (energies). Not all wavelengths of the light are absorbed. The spectrum that is produced by those absorbed is called the absorption spectrum.

2. In this experiment you compared your findings to known line spectra for the unknowns.

The lines given were only the few “strongest” lines in the visible spectrum. Explain what is meant by the “strongest” lines and why some lines are stronger than others.

Within the electronic configurations of each atom there are many paths that will allow the atom to “relax” when returning from an excited to relaxed or normal state. Some of those paths are more common than others resulting in a higher number of emissions of that wavelength. The larger the number of emissions of a particular wavelength the stronger the line will appear in its spectrum.

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Experiment 20: Qualitative Analysis: Cations, Anions, and Complex Ions Instructor Notes and Lab Preparation Chemicals and equipment:                   

Well plates Wax pencils Bunsen burners Hoses Wire pieces 2 flame tests (in the hood) Metals waste jar Kimwipes Multiple sets of the reagents: 6 M NaOH 6 M NH3 6 M HCL 6 M nitric acid Concentrated sulfuric acid 0.1 M nitrates: Na, Mg, Ni, Cr, Zn, Ag, Pb 0.1 M NaNO3, NaCI, NaCL, NaSO4 0.5 M NaI (acidified) 0.5 M FeSO4*7H2O Unknown containing 2 metals

Setup:    

Set up 2 flame tests in the hood Place waste jar in the hood Place 1 set of reagents on each lab bench Place all other supplies and chemicals on front lab bench

Preparation: Prepare 3 or 4 unknowns, each containing 2 different metals. Label as unknown 1–7. Place each chemical solution in a dropping bottle. Note: If students will be sharing chemicals, we have found it helpful to mark the paper covering the counter with the names of each chemical so that students return the bottles back to where they find them.

Set up a hot water bath in the hood. .

Set out small weighing boats.

Instruction Notes: This is a shortened, one day, all-inclusive qualscheme lab. We have cut way down on both waste and chemical costs by microscaling the tests used to confirm the unknown content.

Points to Stress: Good note taking and careful observations are very important in this lab. Make sure that students know to not only write about the color of the precipitates but also their texture, etc.

Make sure that the students understand that having a plan of attack when starting on the unknown is essential to complete the lab in a timely fashion, and to be confident about their results. This is supposed to be a qualscheme activity so they should proceed through a standard set of reactions to separate the various ions, and to avoid false positives.

Potential Problems: We have white well plates which can sometimes make seeing the detail of white precipitates difficult. We tell students that they can also run the tests in a test tube if they have trouble making good observations. Students also often forget about the anions in the mixture and, since they should be finding both, must be reminded of this aspect of the lab.

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Pre-Lab Key

Experiment 20 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) The qualitative analysis scheme depends on the relative solubilities of the salts of the cations with various anions. The aqueous solubility of a salt can be estimated from Ksp, the solubility product constant, describing the equilibrium reaction between the solid salt and the ions in solution: MxAy(s)  xMy+(aq) + yAx-(aq),

Ksp = [My+]x[Ax-]y

(MxAy does not appear in the expression because the pure solid is considered to be in its standard state with an activity of 1.)

The actual solubility may differ slightly from the calculated value, though, if the undissociated salt itself is slightly soluble, and if there are intermediate ion-pair structures formed in solution. Nevertheless, the calculated value is still very useful in comparing solubility relationships of various ions, and predicting whether or not a precipitate will form when two ionic solutions are mixed.

For the salt BaSO4, Ksp = 1.1×10-10.

What is [Ba2+] for a saturated solution of BaSO4? 1.0e-5 M What is the calculated solubility for BaSO4? 2.4e-4 g/dL Addition of a soluble sulfate, such as Na2SO4, will suppress the solubility of the Ba2+ ion because of the common ion effect.

What will [Ba2+] be in the presence of 0.15 M Na2SO4? 4.4e-10M

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2) For the salt SrF2, Ksp = 4.3×10-9.

What is [Sr2+] for a saturated solution of SrF2? 1.0e-3M What is the calculated solubility for SrF2? 1.3e-2 g/dL What will [Sr2+] be in the presence of 0.10 M NaF? 4.3e-7M 3) To estimate whether a salt will precipitate under a particular set of conditions, a reaction quotient, Qsp, can be calculated from given concentrations of the cation and anion from the equation: Qsp = [My+]x[Ax-]y If Qsp is less than or equal to Ksp, no precipitate will form, but if Qsp is greater than Ksp, a precipitate will form and the ion concentrations will decrease until they reach values that satisfy the Ksp expression.

In the case of metal hydroxides, the anion is the hydroxide ion, and its concentration depends on the pH of the solution. Whether a precipitate occurs will therefore depend both on the Ksp of the hydroxide and on the pH. For the reaction: Mg(OH)2  Mg2+ + 2OH−, Ksp = 5.6×10-12.

You prepare a solution of 0.20 M Mg2+ and adjust the pH to 6.0.

What is the value of the reaction quotient, Qsp, under these conditions? 2.0e-17 Will a precipitate form? No What will be the equilibrium concentration of the Mg2+ ion? 2.0e-1M

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4) Insoluble salts can sometimes be brought into solution by the formation of a soluble complex ion. For example AgCl(s) can be dissolved in ammonia solutions by the formation of the soluble Ag(NH3)2+ complex ion. The ammonia competes with the Cl− ion for reaction with Ag+ as follows: AgCl(s)  Ag+(aq) + Cl−(aq) Ag+(aq) + 2NH3(aq)  Ag(NH3)2+(aq)

Sum:

AgCl(s) + 2NH3(aq)  Ag(NH3)2+(aq) + Cl−(aq)

Assuming all the Cl− comes from the dissociation of AgCl, then [Ag(NH3)2+] = [Cl−], and or What is the maximum [Ag(NH3)2+] that can be obtained by dissolving AgCl in 1.5M ammonia? 8.0e-2M 5) For another silver salt, AgBrO3, Ksp = 5.4×10-5.

What is Knet for the reaction AgBrO3 + 2NH3  Ag(NH3)2+ + BrO3−? 8.6e2 What is the maximum [Ag(NH3)2+] that can be obtained by dissolving AgBrO3 in 4.1 M ammonia? 1.4e2M

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Post-Lab Key

Experiment 20 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Which cation forms a 2-ligand complex with NH3?

Silver 2) Which anion does not form a precipitate with lead?

Nitrate 3) Write the balanced equation for the dissociation of the compound in question 2 in water.

Pb(NO3)2(aq) Pb2+(aq) + 2NO3-(aq)

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Experiment 21: Isomerism in Coordination Chemistry Preparation of a Coordination Compound: Tetraamminecopper(II) Sulfate Monohydrate

Instructor Notes and Lab Preparation Chemicals and Equipment: 2 50 mL beakers 1 400 mL beaker 1 glass stir rod 1 25 mL filter funnel 1 125 mL side-arm flask for filtration 1 ring stand 1 Pasteur pipette and bulb 1 watch glass 1 vacuum trap 5 mL of 15 M ammonia solution 5 mL of deionized water ~0.400 g copper sulfate pentahydrate 15 mL of 95% ethanol 10 mL of 1:1 (by volume) ethanol/ammonia solution Ice

Setup: Place copper sulfate pentahydrate on the front counter along with weigh boats.

Place ammonia and ethanol solutions in 1 L containers in the hood.

Instruction Notes: Tetraamminecopper(II) sulfate is readily formed in high yield provided that the “mother liquor” is not too dilute with respect to the reactants. Students should be reminded not to use more solvent than described in the procedures. Although the drying time can be reduced with heating, the use of a drying oven is discouraged because of the volatility of ammonia.

Points to Stress: The procedure given is able to determine the stoichiometry of the compounds within 1–5% without rigorous analytical technique. Students should be reminded however that the more meticulous their technique, the better both their yield and their data will be.

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Potential Problems: The only issues occur if students add too much liquid during the synthesis of the crystals.

Students should be warned to be cautious not to overly “wet” the salt.

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Pre-Lab Key

Experiment 21 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

Use the information to answer the questions. Show work, include units, and put your answers in the blanks.

1) Write a balanced chemical equation for the reaction of CuSO4 · 5H2O with ammonia to form [Cu(NH3)4] SO4 · H2O.

CuSO45H2O + 4NH3  [Cu(NH3)4] SO4 · H2O 2) How many moles of ammonia are present in 5 mL of a 15 M NH3 solution?

0.075 mol 3) How many moles of Cu2+ are present in 1.5 grams of CuSO4 · 5H2O?

0.006 mol 4) Calculate the molar mass of [Cu(NH3)4] SO4 · H2O.

MW = 245.75 g/mol

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Post-Lab Key

Experiment 21 Post-Laboratory Question Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Based on the color of the filtrate, is it possible to obtain a 100% yield on this reaction?

Explain why/why not?

This answer will depend on the student’s results but in general they should respond that a 100% yield is not possible. Reasons may include inadequate mixing, purification, etc.

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Experiment 22: Light Emitting Diode Instructor Notes and Lab Preparation Chemicals and Equipment: Ruthenium(III) chloride trihydrate Sodium tetrafluoroborate Gallium indium eutectic Hypophosphorous acid, 50 wt % in water 2,2’-Dipyridyl Sodium hydroxide Acetonitrile Tin oxide coated glass. This comes with the ICE Nanocrystalline Solar Cell Kit; the kit has ten pieces and this experiment only requires one Voltmeter (with leads) 6-volt power supply, such as a lantern battery Cotton swab

Preparation: “RuCl3” from commercial RuCl3·3H2O Place the commercial RuCl3·3H2O (~3 g) in a mortar and pestle. Place this in an oven (120 to 130 °C) for 3 hours, allow it to cool, and grind it. Place the ground material back in the oven overnight prior to the laboratory experiment. There is most likely some water remaining, as anhydrous RuCl3 is insoluble. The product should be dark brown when dry.

NaH2PO2 preparation from H3PO2 Hypophosphorous acid is monoprotic (two of the hydrogens are attached directly to phosphorus).

It does not fume, but with a pKa of 1.1, it is stronger than phosphoric acid. Use caution when handling, and keep it away from oxidizers (do not store it in a cabinet where the usual mineral acids are kept!). Measure 10 mL of 50 % H3PO2 into a beaker with a stirbar, add 6 mL of water, and begin stirring. Slowly add pellets of NaOH until the pH is about 6–8 (wait until the previous pellet dissolves before adding the next). This will require about 3.9 g of NaOH. The final solution is about 6 M in NaH2PO2. If the solution becomes too basic, additional drops of H3PO2 can be added to bring the pH in the correct range.

Synthesis Dissolve 0.083 g (0.40 mmol) “RuCl3” in 8 mL of water in a 30-mL beaker equipped with a Teflon-coated stirbar. Add 0.188 g (1.204 mmol) 2,2’-dipyridyl and, using a syringe or Mohr pipette, 0.44 mL of NaH2PO2. The liquid will be deep greyish-green. Mark the beaker to indicate the level of liquid, place a watch glass on top of the beaker, and heat so it boils for 30 min. Add

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water as necessary to bring the level to the mark. The liquid will now be deep orange. Allow the solution to cool to room temperature. Dissolve 0.333 g of NaBF4 in 1.5 mL of water and add the resulting solution to the orange solution. Solid should form. Heat the mixture until the solid dissolves, adding water as needed. Cool in ice; crystals should form. Collect these by filtration and wash with ethanol. Draw air through the orange solid to dry.

Instruction Notes: This experiment is fairly sophisticated and uses chemicals more suited to an organic laboratory.

However, if Parts I–III are completed for the student, the production of the light emitting diode is fast and simple enough to complete in the general chemistry lab.

Points to Stress: The relevance of this experiment to what is currently being researched is important. Thin organic materials such as the diode created here have only been produced in a meaningful way within the last 10 years or so and the production of better “diodes” is an area of very active research. The ability to make a light from a thin, easily produced chemical compound has use for flat-screen displays from cars to watches, to safety striping on streets. Understanding the chemical physics behind the functionality of these compounds is therefore important and relevant to any chemist.

Potential Problems: None noted to date.

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Pre-Lab KEY

Experiment 22 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

Use the information to answer the questions. Show work, include units, and put your answers in the blanks.

1) Calculate the theoretical yield of [Ru(bpy)3](BF4)2 (MM = 743.29 g) recovered. Use the RuCl3 as the limiting reactant.

2) Identify the central atom, ligand(s), and coordination number in the coordination compound [Ru(bpy)3](BF4) 2. Draw a sketch of the complex ion.

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3) Would the photoluminescence in this reaction be classified as fluorescence or phosphorescence? Why?

It is phosphorescence because it took several seconds for the light to appear whereas if it were fluorescence it would happen almost instantly.

4) When the negative and positive leads were switched, no light is emitted. Why is this and what does this prove?

This shows that the circuit only allows current to flow one way, and proves that the device is a diode.

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Post-Lab Key

Experiment 22 Post-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Calculate the percent yield of [Ru(bpy)3](BF4)2 (MM = 743.29 g) recovered. Use the RuCl3 as the limiting reactant.

The answer to this question will depend on the student’s data and should use the theoretical yield calculated in their pre-lab problem 1.

2) Does the [Ru(bpy)3]2+ complex exhibit a small or large crystal-field splitting energy? Explain. (Hint: Think about the color of the pure compound.)

It has a large crystal- field splitting energy. The color of a compound is the wavelength of light that an object reflects. Since it is orange, the compound reflects a relatively longer wavelength of light. The longer the wavelength reflected, the shorter the wavelength absorbed, or the larger the crystal-field splitting energy.

Copyright © 2014 Pearson Eduaction, Inc.

Copyright © 2014 Pearson Eduaction, Inc.

Experiment 23: Amino Acid Analysis by Ninhydrin Instructor Notes and Lab Preparation Chemicals and Equipment:        

Pencils Rulers TLC Plates (1/student). Avicel microcrystalline cellulose 20x20cm 250microns. Cat# 05011 (Analtech uniplate, analtech.com) TLC Reservoirs: placed in hood.

Hot Plates/Heat Gun/oven (set to 110oC) Glass capillaries (Kimax-51 EF3773B, size 0.8-1.1x100mm) Bunson burners and strikers Waste container for used solvent (maybe recycled)



3 Different Amino Acid Standards (Glutamic Acid, Lysine, & Serine) o To obtain 12 mg/mL stock solutions, weigh out 1.2 g of each amino acid and dissolve it in 100 mL of ddH2O. Each of the three standards goes in its own well-labeled reagent bottle with the concentration and name of amino acid. Note: For Glutamic acid (which has a solubility of 0.864g/100ml (8mg/ml)), make an 8mg/ml solution.



5 Different Unknowns (5 combinations with each containing two of the standard amino acids and supplied in a reagent bottle labeled as “Unknown A”, “Unknown B”, etc.). o Production of these “unknowns” is nearly identical to the procedure for the standard solutions given previously. The only difference is that in these solutions there will be 1.2 g of EACH amino acid that will be in the unknowns.  Unknown A = 0.8 g of glutamic acid dissolved in 100-mL ddH2O.  Unknown B = 1.2 g of lysine dissolved in 100-mL ddH2O.  Unknown C = 1.2 g of serine dissolved in 100-mL ddH2O.



Chromatography Solvent (1-Propanol:di water 1:1 volume ratio) o Prepare enough for 200 mls/student (~4L/class), place it in 1 L bottles, and keep in the hood.

(Can be recycled.)



0.2% Ninhydrin Solution (w/v): 0.2 g in 100 mls ethanol. The ninhydrin is sprayed onto the TLC plate using a Spra-tool spray gun. Pour the ninhydrin solution into the the polyjar and attach to the spraygun. Place in hood along with cardboard box spray shield and tray.

Setup:  

Standard solutions, unknown solutions, pencils, rulers, TLC plates, and heat gun should all be placed on the front counter.

The TLC reservoirs should be placed in the hood, along with a waste container for used solvent so that it may be recycled.

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

4 functional hot plates should be placed at each bench.

Instruction Notes: Each student will have one of three unknowns to identify: (1) Unknown A = Glutamic Acid; (2) Unknown B = Lysine; and (3) Unknown C = Serine. The two essential values which must be calculated and discussed are the unknown’s molecular weights and isoelectric point. The molecular weight can be deduced from the equivalence volume obtained, and the weight of unknown titrated, while the isoelectric point can be determined from the student’s generated curve, which must be turned in showing the pertinent data as outlined in the graphs on the following pages.

Unknown A (Glutamic Acid)

Unknown C (Serine) Unknown B (Lysine)

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Pre-Lab Key

Experiment 23 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) How would you expect the amino acids glycine, alanine, and phenylalanine to travel relative to each other on the silica plates using a mixture of chloroform/methanol/ammonia? Explain your answer.

As glycine, alanine, and phenylalanine all possess side chains that are essentially nonpolar (glycine is an exception in some texts), these amino acids will be separated only by their molecular weight as the mobile phase in this experiment is extremely polar.

Essentially, the amino acids will “spot” from greatest molecular weight to smallest (the heaviest amino acid will be towards the bottom of the plate, while the lightest will be towards the top). Thus, there will be three spots moving up the silica plate, the first spot will be phenylalanine, the second alanine, and the third glycine.

2) You spot an unknown sample on a TLC plate and elute it with chloroform.

After visualizing the plate, you detect only one spot with an Rf of 0.97.

Does this indicate that the unknown material is a pure compound? What could be done to verify the purity?

Yes and no! The unknown material could be a pure compound if it truly “spotted” immediately before the solvent front. However, the unknown material may also not be a pure compound on the basis that components making up the unknown were nearly identical to those used in the solvent, and were unable to efficiently separate. To truly validate the purity of the compound a different solvent could be utilized to run the separation again, or a sensitive analytical technique such as GC-MS could be utilized.

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3) You spot a sample of an unknown pure compound on a TLC plate and elute it with hexane. It’s Rf is 0.25. Would you expect its Rf to increase, decrease, or remain the same if you elute with acetone? Why? Hexane is non-polar whereas acetone is polar and since the unknown had a small Rf, its affinity for the non-polar solvent was relatively low and did not want to move up the plate with the solvent. On the other hand, if acetone was utilized as the solvent of choice, the Rf of the unknown would most likely increase as its affinity for the solvent would now be greater than what it was for the non-polar hexane.

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Post-Lab Key

Experiment 23 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

Answer the following questions: 1) Why is it important that the spots you made on the TLC plate not be more than 1 mm in diameter? What might have happened if the spots had been larger?

Large spots cause bleedover between lanes.

2) We saw in the mechanism for the ninhydrin reaction that the molecule reacts with amino acids via their amine group. What other molecules might give a “false positive” reaction with ninhydrin?

Cleaning products like bleach or ammonia can react with ninhydrin. Any other molecule that contains an amine group is possible.

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Experiment 3: Halogen Reactions Instructor Notes and Lab Preparation Chemicals and Equipment: Solid Salts: KI, KBr, KCl, NaI, NaBr, NaCl 0.1 M solutions: KI, KBr, KCl, NaI, NaBr, NaCl Saturated I2 1:1 Clorox/H2O Any brand of household bleach should be ok. Ultra products are 6.25% NaOCl. Many of the procedures I use specify freshly opened bleach, but I think a new jug should last for a one-week lab cycle. Scented products are not good idea.

1.0 M HCl Hexane—place in hood Re-pipette dispenser with aqueous Br2 Waste container labeled “Halogens Waste” Crayons or colored pencils (otional) Setup: 

Br2, Hexane, and the Halogens waste container need to be placed in the hood.



The remaining chemicals can be placed on the front lab bench in an organized fashion.

Prepation: Prepare 0.1 M solutions of each of the reagents.

Saturated I2- add solid iodine crystals and dilute with DI water, excess crystals should remain after shaking.

Br2- Prepare in the hood, add 2.5 mL of concentrated bromine to a 1 L volumetric flask and dilute to the mark with DI water. Transfer solution to a re-pipette container. Adjust the setting on the re-pipette container to dispense 1 mL. Instruction Notes: The Halogens experiment is a personal favorite of mine because it provides one of the few opportunities in a first semester lab to actually teach chemistry. So much of what we teach in the

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first semesters of chemistry are physical rather than chemical based, that it is a nice break to actually discuss the chemical basis for the viability of a reaction. I have always worked with both sodium and potassium salts to help the students understand that based on periodic trends groups have similar physical and chemical properties. If cost is an issue a single salt type could be used.

There are a multitude of teaching points to this lab: 1) Electronegativity should be discussed as a foundation for polarity.

2) Polar and non-polar solvents should be discussed with respect to the solubility of polar and ionic solutes; “like dissolves like.”

3) The density of the two solvents should be discussed in conjunction with the discussion about why the two solvents form layers.

4) Electron-transfer reactions can be reviewed as these reactions are also oxidationreduction processes.

5) Quantitative experimentation versus qualitative experimentation can be discussed.

6) The distinction between observations and conclusions should be mentioned.

The experiment runs as follows: Each student should be assigned a group of salts and salt solutions to test. The first test is simple solubility. A small amount of salt is placed in a test tube containing water or hexane. This test determines the solubility of the halide salts.

The second part of the experiment is the most important. Each aqueous halogen is placed in a test tube and then a layer of hexane added. The halogens then migrate into the hexane layer. It is very important that the students get a good determination of the color of each of the halogens in the hexane layer as this will be how they determine if a reaction does or does not occur in the third part of the experiment.

The colors for each halogen are: chlorine (yellow), bromine (orange), and iodine (pink). I started providing crayons for the students to use a few years ago when I realized how very bad they were at describing colors accurately. I have them draw the test tube and color in what they observe. It is especially helpful during the last part of the experiment.

In the third part of the experiment, each of the halogen solutions are combined with each of the halide solutions. Here is what should occur: Cl2 + 2Br-  2Cl- + Br2

Since chlorine is more electronegative than bromine, a reaction will occur and the color observed in the hexane layer will be orange.

Cl2 + 2I-  2Cl- + I2

Since chlorine is more electronegative than iodine, a reaction will occur and the color observed in the hexane layer will be pink.

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Br2 + Cl-  No Rxn

Since bromine is less electronegative than chlorine, a reaction does not occur and the color observed in the hexane layer will be orange.

Br2 + 2I-  2Br- + I2

Since bromine is more electronegative than iodine, a reaction will occur and the color observed in the hexane layer will be pink.

I2 + Cl-  No Rxn

Since iodine is less electronegative than chlorine, a reaction does not occur and the color observed in the hexane layer will be pink.

I2 + Br-  No Rxn

Since iodine is less electronegative than bromine, a reaction will not occur and the color observed in the hexane layer will be pink.

Potential Problems: Two problems that occur somewhat frequently when running this experiment are the faint coloration of the chlorine solution and improperly mixing the solutions in the test tubes. The chlorine solution is made by using hydrochloric acid in reaction with bleach. If too much bleach is used the yellow color of the chlorine is very faint. Two ways to make the yellow color easier to see are to reduce the amount of bleach used and to use 0.5 mL of hexane instead of 1 mL. The smaller hexane layer will concentrate the color. The second problem will require that you show the students the proper way to mix solutions in a test tube and make sure they mix the solution for more than a few seconds. The iodine solution will become a vibrant hot pink if mixed thoroughly, but can form a brown layer at the interface of the water and hexane layers if not mixed properly.

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Pre-Lab Key

Experiment 3 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) When you go into your laboratory to carry out the halogen displacement reactions, you already know from lecture that the order of reactivity, and electronegativity, is Cl2>Br2>I2, so you are able to predict which results to expect from the various combinations of halogen and halide salt.

Let us suppose, though, that you are a student with a color perception disorder, so that you see the color of the halogens differently from other students. And suppose the halogen samples aren't identified for you, but are given as unknowns, labeled X2, Y2, and Z2 for the halogens, and X−,Y−, and Z− for the corresponding halides. You are to identify the relative reactivity of the unknowns, and hence their identity, by mixing different combinations of halogen and halide, and identifying the halogen product from the color of the hexane layer after extraction.

As a start, you first add the free halogens to your test tube, extract into hexane, and get the following results: Z2 hexane layer is red; X2 hexane layer is yellow; Y2 hexane layer is blue.

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Next, you mix the following combinations of reagents and record the resulting color of the hexane in your test tubes, giving the results shown in the table below.

Using the results of the tests provided, write the balanced reactions for each of the six possible combinations of X2, Y2, Z2, X-, Y-, and Z-.

X2 + 2Y-  2X- + Y2 X2 + 2Z-  NR Y2 + 2X-  NR Y2 + 2Z-  NR Z2 + 2X-  2Z- + X2 Z2 + 2Y-  2Z- + Y2

Reagent Combination

Resulting Color

Y2 + Z−

red

X2 + Y−

blue

Z2 + X−

yellow

2) Complete and balance the following reactions as a "molecular" equation. If no reaction occurs, write NR in place of giving products.

2NaI + F2  2 Na+ + 2F- + I2 (note: Na+ is spectator and may not be in rxn) Complete and balance the following reaction as a "net ionic" equation. If no reaction occurs, write NR in place of giving products.

NaBr + At2  NR

3) For each of the following solvents, determine if they are polar or non-polar.

Explain your conclusions.

a.

b.

Non-Polar, due to symmetry

Polar, electronegative oxygen and lack of symmetry

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Post-Lab Key

Experiment 3 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) This experiment uses a qualitative colorimetric process. What is meant by this statement?

Qualitative means that the quality of the observations are used to determine the outcome of an experiment rather than a measured (numerical) or quantitative value.

2) All of the halogens with the exception of fluorine make strong acids. Using electronegativity, explain why hydrofluoric acid is a weak acid.

Fluorine is the most electronegative atom in the periodic table and as such creates a relatively strong bond to the positive hydrogen atom. This lowers the dissociation constant of HF making it a weak acid. Another consideration is HF’s ability to hydrogen bond helps to stabilize the molecule, making it less likely to dissociate. Finally, F- is a very unstable conjugate base for HF because of its small ionic size. This instability also makes the dissociation of HF less energetically favorable and thus HF is a weak acid.

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Experiment 4 Paper Chromatography: An M&M’s® True Colors Instructor Notes and Lab Preparation Chemicals and Equipment: 1 bag of M&Ms 1 box of toothpicks 1 box of food coloring 1 2 L bottle of chromatography solution 1 waste jar 25 1000 mL beakers (***in the lab already***) 1 box of pencils 1 box of saran wrap 25 rulers 1 bag of rubber bands 2 boxes small filter paper Setup: Place a beaker of M&Ms on the front counter.

Chromatography solution should be kept in the hood to reduce fumes.

Preperation: Chromatography solution: 3:2 solution of ethanol and water Instruction Notes: This is a simple chromatography experiment and is particularly nice because there is no need to create “unknowns” for the procedure. The other nice feature is that you can run the experiment on a time rather than distance basis, so you can allow the students to leave the lab and return in time to remove the chromatogram from the solution. Since most real research is done by a process of setup and then allowed to run for extended periods of time and then removed later, this process gives the students a feel for how many actual experiments would be completed in a research lab. I often think students feel all experimental procedures are completed within three hours; the fact that you might have an experiment run for days or weeks is completely foreign to them. I use this as a teaching point for this lab.

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Potential Problems: The only problem that occurs occasionally is a student over saturating a spot with color or chocolate. To properly spot the chromatogram with sufficient shell color they need to spot, and then allow the spot to completely dry, and then spot again over the same point. The students often are impatient to complete the setup and spot again too soon. This causes the spot to spread and become too large. This produces weaker bands and also tends to have lane bleed over.

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Pre-Lab Key

Experiment 4 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Below is a paper chromatogram comparing four standard dyes (in lanes 1–4, Red, Green, Blue, Yellow, respectively) with six dye mixtures (lanes 5–10). The position of the solvent front is noted by the wavy blue line. (A number of factors can affect the mobility of the solvent, including non-uniformity in the paper and temperature variations within the chamber while the chromatogram is being run. Ideally the front should be horizontal, but this problem illustrates a case where it is not.)

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The relative mobility of substances in thin layer and paper chromatography is recorded as an Rf value. The Rf value is simply the ratio of the distance the substance has moved from the origin (measured to the front of the spot), divided by the distance the solvent front has moved. It is important to record the solvent front position before the chromatogram dries completely. When the solvent front is irregular, as shown in the example above, the solvent front distance should be measured above the spot. Identical substances should have the same Rf value even if they do not line up exactly on a horizontal line.

Measure the Rf values of the the four standard dyes in lanes 1–4 to 2 significant figures.

Rf of the red spot in lane 1: ___0.45__

Rf of the blue spot in lane 3: ___0.55__

Rf of the green spot in lane 2: _0.79__

Rf of the yellow spot in lane 4: __0.20__

Each of the dye mixtures in lanes 5 through 9 contains only one of the four standard dyes.

Identify the standard dye in the mixture by comparing its Rf .

Which of the four standard dyes is contained in the mixture in lane 7? ___Blue______ .

Which of the four standard dyes is contained in the mixture in lane 10? ___Red___ Which of the four standard dyes is contained in the mixture in lane 9? __Yellow___ Which of the four standard dyes is contained in the mixture in lane 5? __Green___ 2) For each of the dyes shown in the background, answer the following questions: a) What is the molecular weight? (Unit=g/mol) b) How many double bonds are present? (Include the S=O double bonds.) c) What is the total number of charges on Blue No. 1? (Do not count the Na+ counterions.) d) How many sulfonate (-SO3−) groups are there? e) Some dyes are classified as "azo" dyes because they have the azo group (-N=N-) as part of the structure. Indicate which of the dyes are azo dyes.

Yellow # 5 Blue #1 Red #40 Red #3

a) 534.36 g/mol; b) 14; c) 3; d) 2; e) Yes a) 792.85 g/mol; b) 22; c) 4; d) 3; e) No a) 496.42 g/mol; b) 13; c) 2; d) 2; e) Yes a) 534.36 g/mol; b) 14; c) 3; d) 2; e) No

3) Using the information from 2) above and your knowledge of polarity, make a prediction of the rankings of the food dyes from largest to smallest Rf values and explain your reasoning. Copy your predictions into your lab notebook for later use.

This ranking will vary based on the student’s reasoning of how much the structure, charges, and mass play into the movement of the dyes. I generally give full credit for any well-reasoned answer.

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Post-Lab Key

Experiment 4 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Using the internet or a chemistry textbook, find at least three other types of chromatography techniques and list them.

Ion-exchange Chromatography, HPLC, Reversed Phase Chromatography, Gas Chromatography, etc.

2) Explain the role of the solvent in this experiment. Why was a 3:2 ethanol water mixture used?

The solvent in chromatography support the molecules to be separated up the chromatogram. The solvent is designed to attract the dyes but not bind tightly to them so that they can “fall out” of the solvent when their mass/structure becomes too much for their attraction to the solvent to overcome.

3) Now that you have completed the experiment, compare your actual results to the predictions you made regarding the Rf values of the food dyes. Explain any discrepancies.

This should be a well thought out explanation and comparison to their predictions from the pre-lab. Each dye should be considered and supporting information, including Rf values, used to support their arguments.

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Experiment 5: Molecular Models Instructor Notes and Lab Preparation Chemicals and Equipment:

 12–14 protractors  6 different molecular models Instruction Notes: This lab was developed after many years of trying to get students to build models for themselves and meeting with heavy resistance. In order to make the lab more investigative, I started building the models myself and asked the students to identify them instead. I use the large scale models normally used for demonstrations in lecture. They are glued together to make them more durable for use from class to class. You can also use styrofoam balls and dowels purchased from the local craft shop if you don’t want to purchase the demonstration model kits. We spray painted the models ourselves to expand on the normal number of central atoms available. I don’t recommend using the smaller model kits since they reduce the number of students that can work on a model at one time. I set out six models at a time and switch out the models from section to section.

Potential Problems: When using the protractors it is important to center them over the central atom. If the protractor is set to the outside of the central atom the angles are not accurate.

The students should also be sure to determine the number of valence electrons based on the atoms in the structure first. This allows for better identification of the ions.

Below is a list of the molecules we build and the color code we use for each element: UNKNOWN A–BeF2 UNKNOWN B–CO2 UNKNOWN C–BH3

UNKNOWN H–SO42UNKNOWN I–NH4+ UNKNOWN J–PF5

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UNKNOWN D–SO3 UNKNOWN E–NO3UNKNOWN F–CO32UNKNOWN G–H3O+

UNKNOWN K–I3UNKNOWN L- XeO2F2 UNKNOWN M- XeOF4 UNKNOWN N- ICl4-

Bonds: Double – red Single – yellow Elements: O – red N – blue C – black S – yellow P – brown Xe – silver Cl – green I – purple F – orange Be – gold B – light blue H– white Lone-pair electrons – clear sticks (Important for ionic structures)

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Pre-Lab Key

Experiment 5 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Make six or more copies of the worksheet provided at the end of the procedure section and bring them with you to lab.

This question is here so that you can give it “teeth” by removing points if they do not remember to bring the forms with them.

2) Sulfur is a Period 3 element and as such can create structures that exceed the octet rule.

Because of this ability, there are 2 resonance structures for the sulfuric acid molecule. a) Draw the two possible structures; b) Calculate the formal charges; and c) Determine the best structure and explain your choice.

O H

O

S

O

H

O

b) H (FC) = 0; S (FC) = +2; O (FC) = 2 at -1, 2 at 0

b) H (FC) = 0; S (FC) = 0; O (FC) = all at 0

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c) The second structure is best as it minimizes the formal charges on the structure.

3) For each of the molecules listed below, complete the following: a. b. c. d. e. f. g.

Calculate the number of valence electrons Draw a good Lewis structure and any resonance structures Calculate the formal charges for each atom Calculate the oxidation numbers for each atom Determine the electron cloud and molecular geometries for the central atom(s) List the bond angles Determine if the structure is polar or non-polar

Molecules: Dichromate ion, Iron II Hydroxide, Disulfur Decafluoride Dichromate Ion a. Valence electrons = 54 b.

c. Formal charges = -1 on each oxygen with 3 lone pairs; 0 on the center oxygen; +2 on each Chromium d. Oxidation Numbers: Cr = +6 ; O = -2 e. Each Cr is tetrahedral f. All 109 degrees for the O to Cr bonds; 180 degrees from Cr-Cr g. Polar Iron II Hydroxide a. Valence Electrons = 14 b.

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c. Formal charges = -1 on each oxygen; 0 on each hydrogen; +2 on iron d. Oxidation Numbers: Fe = +2; O = -2; H = +1 e. N/A ionic bonding f. N/A ionic bonding g. Ionic (extreme polarity) Disulfur Decafluoride a. Valence Electrons = 82 b.

c. Formal charges = 0 on each fluorine; 0 on each sulfur d. Oxidation Numbers: S = +5 ; F = -1 e. Each S is trigonal bipyramidal f. Equatorial F 120 degrees apart; S-F equatorial bonds 90 degrees; SF axial bonds 180 degrees g. Non-polar

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Post-Lab Key

Experiment 5 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) What physical forces create the three-dimensional structure of a molecule?

The electron repulsion is the predominant force that creates a molecule’s structure as the electrons find the lowest energy state that puts their negative charges as distant from each other as possible.

2) Now that you know more about molecular structure, discuss how the structure of a molecule influences its polarity.

Polarity is a combination of the pull of an electronegative atom on the electrons in a structure towards itself with or against the pull of other atoms in the molecule. If a structure contains two or more equally electronegative atoms arranged in a symmetrical manner they will essentially cancel each other out yielding a non-polar molecule.

3) Oxidation Numbers and Formal Charges are often confused by first-year chemistry students. If you were teaching the course, explain how you would differentiate the two types of values so that the concepts would be clear to your students.

Oxidation numbers are based on the electronegativity of an atom and are therefore assigned in a hierarchical manner with more electronegative atoms being assigned the bulk of the electron density or charge. In formal charges, the charges are assigned as if each atom shares the electron density equally with no preference being given to the more electronegative atom. Thus, formal charges are a property of the molecule the atom finds

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itself in and oxidation numbers are a physical property of the atom depending only on its electronegativity ranking.

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Experiment 6: Limiting and Excess Reagents Instructor Notes and Lab Preparation: Chemicals and Equipment: Centrifuges Disposable pipet (and bulbs if glass pipet are used) Grease pencils or labeling tape 1 pair of heat protective gloves Dessicators Dessicant Large filter paper Aqueous Metals Waste jar (in hood) Dessicant waste jar this can be reused indefinitely) 0.5 M Cu(NO3)2 0.5 M KI Methanol

Setup: 

Place one bottle of each reagent on each lab bench.



Each lab bench should have a working centrifuge.



Ovens should be set to 70 degrees.



Place waste jar in hood labeled as “Aqueous Metals Waste.”

Preparation: Dessicant: Mix 2/3 white dessicant and 1/3 blue dessicant 0.5 M Cu(NO3)2: For 1 L, slowly add 93.78g of Cu(NO3)2 and dissolve with DI water. (Supply for ~6 sections of 24 students)

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0.5 M KI: For 1 L, slowly add 83.01g of KI and dissolve with DI water Methanol is stored in the flammable cabinet.

Instruction Notes: There are several nice points to this experiment that can be pointed out to the students as you go along. The first is that a solution does not have to be colorless to be considered “clear.” In all the experiments I performed as a first year-student, we never came across a supernatant that wasn’t colorless. The supernatant in this experiment is a clear green/yellow. You should remind your students that “clear” in this case means without cloudiness, not without color. Another common misconception among young chemists is that equal amounts of reagents yields the maximum amount of product. This probably arises because most reactions used in first-year labs are 1:1 reactions. We developed this lab after several failed attempts to find other combinations that were not 1:1. We tried PbI, which makes a great yellow precipitate, but who really wants that much lead all over the place? We also tried a hydroxide precipitate but it cooked down to oxide in the oven, complicating the issue. The copper iodide reaction is relatively non-toxic and not 1:1, and the precipitate is stable so it is a good reaction for this experiment.

The graph we ask the students to draw appears as a lop-sided pyramid. This also reinforces the notion that the maximum output of product does not mean identical or even amounts of reactants.

Potential Problems: If your lab uses a variety of test tubes (pyrex, non-pyrex, lipped, and not) have your students check to make sure that the test tubes they will be centrifuging will fit in the centrifuge before starting the experiment. They should also make sure the general appearance is the same, i.e., both lipped or not, as this will affect the weight of the tubes and if too disparate, unbalance the centrifuges.

Occasionally, a student will have an allergy to iodine. Most of these allergies are consumption based, but we issue gloves and have the student work at the hood instead of a bench just to be safe.

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Pre-Lab Key

Experiment 6 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

Your laboratory experiment on "Limiting and Excess Reagents" involves the reaction between aqueous solutions of Cu(NO3)2 and KI to form CuI and I2. The CuI precipitates from the reaction as a solid, and the other products (KNO3 and I2) remain in aqueous solution. In the lab you will vary the relative amounts of each reactant, determining the amount of solid product in each case. In this pre-lab you will explore some features of this reaction.

1) This reaction is an oxidation-reduction (ReDox) reaction in which one reactant loses electrons (is oxidized) and another reactant gains electrons (is reduced). To recognize a ReDox reaction you need to check the oxidation states of all the elements in both the reactants and products. If any change their oxidation state, the reaction is a ReDox. An increase in oxidation state is an oxidation; a decrease in oxidation state is a reduction.

(Note: If an element is a reactant or product, converted to or formed from a compound, the reaction must be a ReDox reaction.)

Indicate the oxidation state of the elements in each of the following compounds: Reactants

Products

Cu(NO3)2

KI

CuI

KNO3

I2

Cu +2

K +1

Cu +1

K +1

I 0

N +5

I

-1

N +5

O

-1

I

-2

O -2

Which reactant is oxidized, and which is reduced?

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Iodine is oxidized, copper is reduced.

2) Calculating Molarity: What is the formula weight of CuI? 190.45 g/mol What quantity of CuI is present in 50.0 mL of a 0.50 M solution? 4.76 g What is the formula weight of Cu(NO3)2? 187.57 g/mol What mass of Cu(NO3)2 is present in 225 mL of a 0.30 M solution? 12.66 g What is the formula weight of KI? (Enter unit as g/mol) 166.00 g/mol What mass of KI is needed to add to a 1000 mL volumetric flask to produce a solution that is 0.75 M? 124.5 g What is the formula weight of I2? 253.8 g/mol What is the molarity of a solution produced by dissolving 53 g of I2 in a 250 mL volumetric flask? 0.84 M 3) Limiting Reagent Calculations: You mix 3.5 mL of 0.50 M Cu(NO3)2 with 2.0 mL of 0.50 M KI, and collect and dry the CuI precipitate. Calculate the following: Moles of Cu(NO3)2 used: 1.75 mmol Moles of KI used: 1.00 mmol Moles of CuI precipitate expected from Cu(NO3)2: 1.75 mmol Moles of CuI precipitate expected from KI: 0.50 mmol What mass of CuI do you expect to obtain from this reaction? 0.50 mmol If the actual mass of the precipitate you recover is 0.133 g, what is the percent recovery of the precipitate? 26.6 %

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Post-Lab Key

Experiment 6 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) How would a decrease in the concentration of the reactants affect the moles of reactant?

If the concentration is lowered the number of moles is reduced.

2) Explain the role of limiting and excess reagents in an experiment.

When designing an experiment the scientist always attempts to minimize waste. To do this he/she will calculate the masses or volumes of reactants in such a manner as to produce the amount of product needed without significant amounts of either reactant left over. Knowledge of how reactants will be in excess or limiting, i.e,. the stoichiometry of the reaction, is therefore crucial to a scientist’s ability to conserve chemicals and produce products efficiently.

3) Copper compounds are fairly common. Using the internet or any other reliable source, look up at least five other copper compounds. List the name, chemical formula, and uses of each compound. Be sure to cite your source in acceptable format.

There are a multitude of copper compounds that can be listed. Just make sure the student notes where he/she took the information.

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Experiment 7: Redox Reactions in Voltaic Cells: Construction of a Potential Series Instructor Notes and Lab Preparation Chemicals and Equipment: Foam cups Pieces of white card stock 1 box of large (15 cm circles) coarse filter paper Sheets of sand paper Multi- or volt-meters & alligator clips 1.0 M NH4NO3 1.0 M Cu(NO3)2 1.0 M Mg(NO3)2 1.0 M Pb(NO3)2 1.0 M Fe(NO3)3 1.0 M AL(NO3)3 0.01 M HNO3 Pieces of each metal: Cu; Mg; Fe; Pb; Al Aqueous Metals Waste jar (in hood)

Preparation: Prepare solutions in 0.01 M HNO3. Due to the large quantity of substance required, prepare in advance as it will take several hours to completely dissolve. Heat may be applied to speed up the process. Also, add small quantities at a time for easier dissolving.

1.0 M NH4NO3—80.04 g per 1 liter 1.0 M Cu(NO3)2—187.56 g per 1 liter 1.0 M Mg(NO3)2 —148.31 g per 1 liter 1.0 M Pb(NO3)2—331.00 g per 1 liter

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1.0 M Fe(NO3)3 -241.86 g per 1 liter 1.0 M AL(NO3)3 –213.00 g per 1 liter In class instructions: Do NOT discard the foam cups or metals. Metals MUST be cleaned, dried ,and returned to the correct container. ***Mg and Fe are often confused***

Instruction Notes: This is a very simple lab for the students to perform and only takes about an hour to complete.

We use foam cups to minimize the volume of solution we need to use. 100 mL beakers can be used instead but will require almost 100 mL of solution in each to enable the filter papers to make good contact with the solution. The foam cups only require about 20–30 mL cutting the chemical cost way down. We ask the students to clean and dry the foam cups at the end of the lab so that we can re-use them from class to class to minimize the waste.

Dependent on the number of multi- or voltmeters you have available, you can make this an individual, pairs, or group lab experiment. You should make the assignments of the 10 metal combinations so that all ten are completed and reported to the class as a whole. If you have sufficient numbers of students/groups you can also introduce more statistics into the experiment based on more than one experimental potential value being determined for each of the metal combinations.

Potential Problems: The only problem we have ever experienced with this lab is experimental potentials not matching up well to those predicted by calculation. This is generally caused by a poor connection to the voltmeter or poor contact via one or more of the salt bridges. Make sure that both the metals and alligator clips are clean and making good contact with each other and the solution. Also make sure that the salt bridge is soaked through with solution. If the filter paper is wound too tight, it often appears wet but may be dry on the inside.

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Pre-Lab Key

Experiment 7 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

Following is a diagram of a voltaic cell.

1) Identify the parts of the cell shown above by selecting the appropriate letter from the figure.

__E__ __E__ __E__ __B__ __A__

This is the positive electrode.

Reduction occurs at this electrode.

This electrode is the cathode Anions depart from the salt bridge here.

Oxidation occurs at this electrode.

2) Write the shorthand notation for the following voltaic cell reaction: Fe(s) + Ni2+(aq)  Fe2+(aq) + Ni(s)

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Fe(s)|Fe2+(1M)||Ni2+(1M)|Ni(s) Write the reaction represented by the following voltaic cell notation: Mn(s)|Mn2+(1M)||Pb2+(1M)|Pb(s)

Mn(s) + Pb2+(aq) ->Mn2+(aq) + Pb(s) 3) The list below shows the reduction potentials for ten different metal ions.

Calculate ΔE for the ReDox reaction: Ni(s) + Mn2+(aq)  Ni2+(aq) + Mn(s)

ΔE = -0.93 Will this reaction occur spontaneously? No Calculate ΔE for the ReDox reaction: Co(s) + Sn2+(aq)  Co2+(aq) + Sn(s)

ΔE = +0.14 Will this reaction occur spontaneously? Yes

Half-Reaction

Eo(V)

Cu2+(aq) + 2e−  Cu(s) +0.34 Pb2+(aq) + 2e−  Pb(s)

-0.13

2+

−

-0.14

2+

−

-0.25

2+

−

Sn (aq) + 2e  Sn(s) Ni (aq) + 2e  Ni(s)

Co (aq) + 2e  Co(s) -0.28 Cd2+(aq) + 2e−  Cd(s) -0.40 Fe2+(aq) + 2e−  Fe(s)

-0.44

Zn2+(aq) + 2e−  Zn(s) -0.76 Mn2+(aq) + 2e−  Mn(s) -1.18 Mg2+(aq) + 2e−  Mg(s) -2.37

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Post-Lab Key Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) What is a reactivity series and how is it related to this experiment?

A reactivity series of metals is a list of metals ranked in order of decreasing reactivity to displace hydrogen gas from water and acid solutions. It can also be used to predict which metals will displace other metals in aqueous solutions. For this experiment it helps to determine which metal will accept electrons and which will donate electrons by determining the cathode and anode for each cell.

2) Create a potential series ranking all of the metals tested by the class arranged from strongest oxidizing agent to weakest oxidizing agent.

This should be a list reverse order of the reduction potentials shown above for the metals used in your lab.

3) Explain your reasoning behind the rankings you list in problem 2.

The identification of each metal as either the cathode or anode in each cell created develops the direction of flow for the electrons. The metal that is assigned as the cathode in all the combinations will be the strongest oxidizing agent, the metal that is assigned as the cathode in all but one combination is the second strongest and so forth.

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Experiment 8: Reactions in Aqueous Solutions: Strong Acids and Bases Instructor Notes and Lab Preparation Chemicals and Equipment:           

0.5 M HCl 1 liter, dyed red 0.5 M NaOH 1 liter, dyed blue 1 mL graduated pipette (plastic) 3-way rubber bulb 25 mL volumetric flasks 50 mL beaker as needed hotplates pH meter power supply calibration instructions buffer solutions

Setup: The 0.5 M solutions are purchased in standardized form and a few drops of food color are added to make it easier for the students to differentiate the solutions. A hotplate and a pH meter should be placed out for every two students.

Instruction Notes: Points to Stress This is a relatively straight-forward lab that allows students to observe the reaction of a strong acid and base, and how the concentrations of each relate to their pH. We also use the lab to point out the predictability of the production of salt from the acid/base reaction. This reinforces the limiting and excess reagents lab just completed. The lab also introduces the use of pipettes, serial dilution, pH meters, and calibration.

Potential Problems The only problem we have had with the experiment is the pH of the DI water not being 7 or even very close to 7. This complicates the calculation of pH and makes for much greater errors between the student’s calculated and experimental values.

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Pre-Lab Key

Experiment 8 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Constant boiling HCl has a concentration of 11.6 M. Your laboratory assistant is going to dilute this acid for you to create a stock solution of HCl for you to use in this experiment.

What volume of the 11.6 M HCl must the assistant add to a one-liter volumetric flask so that the concentration of the stock solution will be 0.62 M after dilution to the mark with water?

53.4 mL What is the pH of this stock solution? (Remember the significant figure rules for logarithms!)

pH = 0.21 You will now use this 0.62 M HCl solution to carry out a series of dilutions, measuring the pH of each diluted sample. Calculate the [Cl-] and [H+] concentrations, and the pH of each sample. (Remember the significant figure rules for logarithms!)

Sample A will be produced by pipetting 1.00 mL of the stock solution into a 25 mL volumetric flask, and diluting to volume with distilled water.

[Cl-]= 0.025M [H+]=0.025M pH= 1.61 Sample B will be produced by pipetting 1.00 mL of sample A into a 25 mL volumetric flask, and diluting to volume with distilled water.

[Cl-]= 9.9e-4M [H+]= 9.9e-4M pH= 3.00

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Sample C will be produced by pipetting 1.00 mL of sample B into a 25 mL volumetric, and diluting to volume with distilled water.

[Cl-]= 4.0e-5M [H+]= 4.0e-5M pH= 4.40 Sample D will be produced by pipetting 1.00 mL of sample C into a 25 mL volumetric, and diluting to volume with distilled water.

[Cl-]= 1.6e-6M [H+]= 1.6e-6M pH= 5.80 Sample E will be produced by pipetting 1.00 mL of sample D into a 25 mL volumetric, and diluting to volume with distilled water.

[Cl-]= 6.3e-8M [H+]= 1.4e-7M pH= 6.86 Sample F will be produced by pipetting 1.00 mL of sample E into a 25 mL volumetric, and diluting to volume with distilled water.

[Cl-]= 2.5e-9M [H+]= 1.0e-7M pH= 6.99 2) Your laboratory assistant is going to prepare a stock solution of sodium hydroxide (NaOH) for you. What mass of NaOH must the assistant add to a one-liter volumetric flask so that the stock solution has a concentration of 0.35 M after dilution to the mark with water?

14.0 g What is the pH of this stock solution? (Remember the significant figure rules for logarithms!)

pH= 13.54 You will now use this 0.35 M NaOH solution to carry out a series of dilutions in the same manner as in problem 1, measuring the pH of each diluted sample. Calculate the [Na+] and [OH-] concentrations and the pH of each sample A–F. Sample A will be produced by pipetting 1.00 mL of the stock solution into a 25 mL volumetric, and diluting to volume with distilled water.

[Na+]= 1.4e-2M [OH-]= 1.4e-2M

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pH= 12.15 Sample B will be produced by pipetting 1.00 mL of sample A into a 25 mL volumetric, and diluting to volume with distilled water.

[Na+]= 5.6e-4M [OH-]= 5.6e-4M pH= 10.75 Sample C will be produced by pipetting 1.00 mL of sample B into a 25 mL volumetric, and diluting to volume with distilled water.

[Na+]= 2.2e-5M [OH-]= 2.2e-5M pH= 9.35 Sample D will be produced by pipetting 1.00 mL of sample C into a 25 mL volumetric, and diluting to volume with distilled water.

[Na+]= 9.0e-7M [OH-]= 9.1e-7M pH= 7.96 Sample E will be produced by pipetting 1.00 mL of sample D into a 25 mL volumetric, and diluting to volume with distilled water.

[Na+]= 3.6e-8M [OH-]= 1.2e-7M pH= 7.08 Sample F will be produced by pipetting 1.00 mL of sample E into a 25 mL volumetric, and diluting to volume with distilled water.

[Na+]= 1.4e-9 [OH-]= 1.0e-7 pH= 7.00

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Post-Lab Key

Experiment 8 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Explain how this experiment can be related to limiting and excess reagent problems.

The determination of the salt yield is based on the amount of reagents that are added.

2) The acid and base used in this experiment are monoprotic. How would the experimental results vary if a diprotic acid had been used instead?

If the acid were diprotic then it would take twice as much base to neutralize the acid since the molar ratio between H:OH would be 2:1. If the same volumes were used then half as much product would be produced since the base would run out.

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Experiment 9: Calorimetry & Hess’s Law Instructor Notes and Lab Preparation Chemicals and Equipment: Styrofoam cups Thermometers Cardboard lids with hole in center Bottle of 1.0 M HCL Container of Mg metal 2.5 cm strips cut from Mg ribbon Bottle of MgO Bunsen burners Vacuum tubing Pieces of wire gauze Small weighing boats Aqueous metals waste container (in hood) **adjust pH before disposing**

Setup: Place 1 styrofoam cup and 1 piece of cardboard per pair of students Place the aqueous waste container in the hood The rest of the supplies should be set up on the front counter

Preparation: 1.0 M HCL—83.3mL of concentrated HCL per 2L of 1.0 M HCL solution prepared **** Prepare a large quantity****

Instruction Notes: After years of running the typical coffee cup calorimetry, solving for the heat capacity of a metal, I decided that the experiment didn’t cover enough of the material discussed in the thermodynamic chapter in the textbook. The process of solving for heats of reaction and use of Hess’s Law to determine an otherwise undeterminable heat of reaction provides a more complete coverage of the thermodynamics discussed in the book and also gives the students a realistic use for Hess’s Law.

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Points to Stress: The first part of the experiment asks the students to determine the calorimeter constant. We currently use a single 12 ounce styrofoam coffee cup for the calorimetry runs. (We tested the experiment with both single and doubled cups with no appreciable difference in results.) The first part of the experiment is repeated because the students often have difficulty with the technique. This leads to differences in the results of the two experimental runs. It is important to discuss with the students whether they should use the results of both experimental runs or to favor one run over the other and how to justify this decision in their conclusion.

Potential Problems: The only difficulty with the experiment occurs when the students do not stir continuously when adding the MgO powder. If you use stir bars, this problem can be avoided. When you find that you do not have enough stir plates to use this fix, we stress the need for one student to “swirl” the cup continuously while adding the powder.

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Pre-Lab Key

Experiment 9 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Use the graph of Temperature (oC) versus Time (s) provided to answer the following questions: o

What is the initial temperature? 23.9 C o

What is the final temperature? 39.2 C The Cp for the calorimeter (including the HCl solution and the Mg) was calculated to be 0.226 kJ/°C. For the temperature curve shown above, 0.175 g Mg was reacted with HCl. What is the ΔH of the reaction? 480 kJ/mol 2) To calibrate your calorimeter cup, you first put 53 mL of cold water in the cup, and measure its temperature to be 21.1 °C. You then pour 41 mL of hot water, temperature = 53.7 °C, into the cup and measure the temperature every thirty seconds over a 10 minute period. You extrapolate this “cooling curve” back to the time of addition and find that the “final temperature” after mixing is 34.4 °C. What is the heat change of the hot water, qHW? -3310.8 J (Assume the density of the water is 1.00 g/mL, and remember that the specific heat of water is 4.184 J/g-K or J/g- °C.)

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What is the heat change of the cold water, qCW? 2949.3 J What is the heat change of the calorimeter cup, qcup? 362 J What is the heat capacity of the calorimeter cup, Ccup? 27 J/K 3) You would like to measure the heat of neutralization of an acid with a base. You mix 225 mL each of 0.77 M HCl and 0.77 M NaOH, both at a temperature of 18.7 °C in a calorimeter cup equilibrated to that same temperature. After following the temperature change for 10 minutes and extrapolating it back to the time of addition, you find that the “final temperature” after mixing was 23.7 °C. Previously you measured the heat capacity of the calorimeter cup (Ccup) to be 39 J/ °C. Assuming that the density of the two solutions is 1.00 g/mL, and that their specific heats are the same as water, 4.184 J/g-°C, what is the heat absorbed by the two solutions, qsoln?

-9414 J What is the heat absorbed by the calorimeter cup, qcup? 195 J What is the heat change for the neutralization reaction, qrxn?

-9609 J What is the ΔH for the neutralization? -55.5 kJ/mol Write the net ionic equation for this neutralization reaction, including the states of matter (i.e., H+(aq)) for the reactants and products.

H+(aq) + OH-(aq)  H2O(l) 4) Given the following data: SO2(g) → S(s) + O2(g)

∆H = +296.8 kJ

CO2(g) → C(s) + O2(g)

∆H = +393.5 kJ

C(s) + 2S(s) → CS2(l)

∆H = +89.4 kJ

Find the ΔH of the following reaction: CO2(g) + 2SO2(g) → CS2(l) + 3O2(g)

ΔH = 1076.5 kJ .

Post-Lab Key

Experiment 9 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) What is the difference between the heat (q) of a reaction and the reaction enthalpy (HRXN) ?

The heat of a reaction is simply the energy created by a temperature difference. The enthalpy is the sum of all internal energy differences for a reaction.

2) Why is Tfinal taken from extrapolating back to the beginning of the experimental run rather than from the last temperature reading taken?

Tfinal is the theoretical temperature that would be reached in a reaction were all substances to reach thermal equilibrium instantaneously. The extrapolation based on the temperature decay line allows for the estimation of what this temperature would be. The last temperature reading taken is at the end of the temperature decay and does not directly indicate the temperature of instantaneous equilibrium.

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Experiment 1: Conservation of Matter Instructor Notes and Lab Preparation Chemicals and Equipment: 3.0 M HCl in 1-Liter bottles with pour spout Bottle of anhydrous copper (II) sulfate Weighing boats Bottle of zinc metal (mossy) 50 mL volumetric flasks with stopper 125 mL flasks with side arm #4 black rubber stoppers *** DO NOT USE CORK *** 1 bag of 7-inch balloons 1 box of calipers Long tweezers Disposable pipets Parafilm 1 box of baking soda for acid spills Setup:    

Make sure there are Kimwipes near the balance(s).

Set out equipment and chemicals.

Place 1 bottle of HCl on each lab bench.

Place 50 mL volumetric flasks on each lab bench.

Preparation: 3.0 M HCl – Use 247.93 mL of concentrated HCl per 1 L of solution.

Prepare a large carboy: To prepare 12 L of 3.0 M HCl: Fill a carboy with approximately 5 L of water. Add 2.98 L of concentrated HCl (in the hood). Then dilute to 12L with D.I. water. (Supply for 20 classes sections @ 24 students) Mix well.

To prepare 6 L of 3.0 M HCl: Fill a carboy with approximately 3 L of water. Add 1.49 L of concentrated HCl (in the hood). Then dilute to 6 L with D.I. water. (Supply for 10 classes sections @ 24 students) Mix well.

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Instruction Notes: The Conservation of Matter experiment was originally developed for our non-majors course. For that course we prepare the copper (II) sulfate solution for the students. When we adapted it for the majors we decided to have them prepare their own solution and introduce the use of volumetric glassware. If you feel your students are not ready to mix their own solution, this is an easy change that both shortens and simplifies the experiment. The experiment is very visual as both the color of the solution changes dramatically in a short period of time and the balloon is inflated to a large size. This leaves no doubt that a reaction has taken place.

Points to stress: Students are working with 3.0 M HCl. This is strong acid and it should be stressed repeatedly that all safety precautions must be taken to avoid burns. All spills of any size should be cleaned up by lab personnel, not the students. I provide my teaching assistants with baking soda and gloves to handle the clean-up.

When the students are preparing the solution, they should weigh the CuSO4 in a weigh boat and then use a small amount of HCl to assist in the transfer of the salt to the volumetric flask. CuSO4 is a “sticky, wet” salt and thus clumps, making dry transfer difficult (the neck of the 50 mL volumetric flask is small). You can also use paper funnels but the “stickiness” of the salt may cause a loss of mass there as well. We have found that using the HCl to aid in the transfer reduces the loss of mass best and also keeps the lab benches much cleaner.

Potential Problems: As mentioned above, the necks of the volumetric flasks are fairly small so students should be shown the proper way to transfer the salt so that they don’t make a huge mess on the bench tops and, more importantly, don’t lose mass of salt.

The only other problem that sometimes occurs is a hole in a balloon or a balloon that is not securely fastened to the side-arm of the flask. Have your students blow the balloon up once before attaching it to the side-arm. This will allow them to determine if the balloon is air tight and it also stretches the balloon so that the resistance is lowered when the reaction starts to fill it with hydrogen gas.

You might note tweezers in the list of items needed for this lab. Because our lab balances are only capable of weighing up to ~100g of material, at the end of the lab we have to remove the pellets from the solution and weigh them separately using a tared weigh boat. If you have the same difficulty, tweezers work well as long as the students are gentle. The removal and separate measurement also reinforces the notion that the total mass of the products is what is important; not that they all remain in the same state or container.

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Pre-Lab Key

Experiment 1 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Based on your reading of the laboratory Introduction and Background, prepare a hypothesis regarding what results you expect from your experiment. Be detailed.

This answer will vary based on the student’s interpretation of the reason for the experiment but should contain a reference to both the concepts of conservation of matter and also the techniques that will be employed to explore the concepts.

2) Calculate the number of moles or grams in each of the following quantities of salt.

5.233 g of Cu(NO3)2 · 6H2O =

________0.01770 mol____

7.77 g of zinc nitrate trihydrate =

________0.0319 mol_____

6.215×10-2 mol of Na2SO4 · 10H2O =

________20.02 g________

2.0×10-2 mol of zinc nitrate hexahydrate =

________6.0 g_________

3) You would like to make up 100 mL of a 0.23 M solution of Na2SO4 · 10H2O. How many moles of the salt do you need? How many grams of the salt do you need?

0.023 mol .

You have a 2.5 M solution of Cu(NO3)2 · 6H2O. What volume of the solution (in mL) must you measure in order to have 0.50 moles of the salt?

200 mL You have a 0.75 M solution of ZnSO4 · 7H2O. What volume of the solution (in mL) must you measure in order to have 32.5 grams of the salt?

Mol amount = 0.11 mol; Vol = 151 mL 4) In your laboratory you will be required to make a solution of 0.25 M CuSO4 · 5H2O in a 25.00 mL volumetric flask. This exercise will walk you through the steps involved. After completion, write down the values for your answers in your notebook and take them to the laboratory to use in preparing your solution. a. What is the formula weight of CuSO4 · 5H2O? 249.684 g/mol b. How many moles of CuSO4 · 5H2O are required to make 25.00 mL of a 0.25 M solution?

0.010 mol c. What mass of CuSO4 · 5H2O will be needed to make this solution?

1.56 g d. You will add the salt to a weighing boat, which you determine has a mass of 1.9785 g. What will be the reading on the balance when you have put sufficient CuSO4·5H2O in the boat?

3.54 g

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Post-Lab Key

Experiment 1 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Based on your experimental results, does the Law of Conservation of Mass apply to this experiment? Explain.

The answer here will depend on the student’s results. If they have a high value of recovery they should indicate that the Conservation of Mass does apply. If they get a lower value of recovery it will most likely lead them to assume that the conservation of mass is not upheld.

2) If the balloon had expanded more how would the mass of H2 and the mass percent recovered have changed?

Both the mass obtained and the percent recovery would increase.

3) How would increasing the molar concentration of CuSO4 have affected this reaction?

Assuming that there was an excess of zinc, the increase in copper sulfate concentration would lead to a higher product amount but the conservation of mass should stay relatively constant.

4) Zinc metal reacts with acid solution to produce hydrogen gas and zinc ion as follows: Zn(s) + 2H+(aq)  Zn2+(aq) + H2(g) You add 6.825 grams of Zn metal to 29.0 mL of 1.70 M HCl and set up the flask to

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catch the hydrogen in a balloon as in the following picture: How many moles of Zn have you added to the flask?

0.1044 mol How many moles of H+ are in the solution?

0.0493 mol How many moles of Zn will have reacted with the H+ when the reaction is complete?

0.02465 mol How many moles of Zn will remain unreacted?

0.07975 mol What mass of Zn will remain unreacted?

5.214 g How many moles of H2 gas will be produced?

0.02465 mol What mass of H2 gas will be produced?

0.04969 g Assuming that one mole of H2 gas would occupy a volume of 22 L at the temperature and pressure in the balloon, what would be the volume of the balloon?

0.5423 L

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Experiment 2: Atomic Spectra Instructor Notes and Lab Preparation Chemicals and Equipment:

6 M HCl 1 bottle of mossy zinc 1 striker 2 Petri dishes w/opening labeled as unknown 1 & unknown 2 2 coffee cans w/opening for flames Spectroscopes Helium lamp labeled as helium Sodium flame (NaCl) labeled as unknown 1 Lithium (LiCl) flame labeled as unknown 2 Hydrogen lamp labeled as unknown 3 Krypton lamp labeled as unknown 4 Mercury lamp labeled as unknown 5 2 Bunsen burners with vacuum tubing Setup:

On each of the two front lab benches, set up one flame unknown. Place the unknown salt in the Petri dish. Cover and place the Bunsen burner over the opening in the Petri dish. Then cover with the coffee can. When ready to start experiment, place a few pieces of zinc and a few mL of HCl in the Petri dish, cover, and light the Bunsen burner. Place one lamp on each of the remaining four lab benches.

Place the zinc, HCl, striker, and spectroscopes on the front counter.

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Example Coffee can flame set-up and lamp set-up Spectroscopes are inexpensive (< $10) and are available from both Fisher Scientific and Sargent-Welch as well as many other sources. Lamps and bulbs can also be bought from Fischer Scientific and bulbs are available for many other elements as well.

Instruction Notes:

We generally set up five unknown light sources for this lab. Every other semester, I change a few of the sources to keep the experiment fresh. As mentioned above, there are a variety of light bulbs available for the lamps and most of the alkali metals and alkaline earth metals can be vaporized using mossy zinc and HCl to produce colored flames.

Having the students calculate their own wavelengths in the pre-lab and then determine the identity of those elements in the lab really exemplifies the use of spectrometry for identification of materials both near and far.

Points to Stress:

We ask the students to work in pairs to assure that the measurements taken are reproducible. This introduction to the necessity of reproducibility in science should be pointed out. Another important point is the introduction of the need to calibrate instrumentation. To this point, we have assumed that the instruments and equipment being used are in good condition and operating correctly. This experiment is the first time the fallacy of that assumption is explained and .

corrected. The creation of a calibration graph is an important tool that the students will need in future experiments and should not be overlooked as a teaching point in an effort to complete the rest of the lab.

Potential Problems:

The majority of the problems associated with this lab come from inappropriate use of the spectrometers. Make sure the students are aligning the slit in the spectrometer, not the grating, with the light source. Also, too much ambient light will make the lines of the source difficult to see. Keep the light in the lab dim if at all possible.

Unfortunately, you cannot turn off all the lights because then the calibration lines in the spectrometer are invisible as well. Goggles make the reading of the spectrometers more difficult. I generally allow students to work around the lamp sources without the goggles, but because of the acid, have them use the goggles around the flames.

Pre-Lab Key

Experiment 2 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Some of the elements that may be used in this experiment are listed below. For each element write the complete electron configuration. a. Helium 1s2 b. Hydrogen 1s1 c. Mercury 1s 2 2s 2 2p 6 3s 2 3p 6 3d 1 0 4s 2 4p 6 4d 1 0 4f 1 4 5s 2 5p 6 5d 1 0 6s 2 d. Krypton 1s 2 2s 2 2p 6 3s 2 3p 6 3d 1 0 4s 2 4p 6

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e. Sodium 1s22s22p6 3s1 f. Lithium 1s22s1

2) What makes the wavelength of the N=4 to N=2 transition for hydrogen different from an N=4 to N=2 transition for mercury? How is this difference related to the spectra we observe? Explain.

The spacing between energy levels is unique in each atom based on the number of protons in the nucleus and the shielding that each electron feels. Thus the distances between the shells vary and therefore the wavelength of light emitted or absorbed varies based on this difference in energy and distance.

3) You have just carried out the calibration of your spectroscope's diffraction grating and you get the values shown in the table for the visible emission lines of helium. Plot these values and produce a calibration graph that can be used to correct the values in the table below: Color

Spectrometer Reading

Violet

5.3

Blue

5.7

Green

6.1

Yellow

6.8

Orange

7.7

Red

8.3

Spectrometer reading

Corrected Wavelength (nm)

Color of the Frequency Emitted Light (hz)

6.5 5.4 7.9 4.6 8.5

4) The following atomic line spectra display six of the more prominent lines observed in the visible region for the indicated element. The number by each line represents the relative intensity. A photograph of the actual spectrum of the element may show many more lines, or sometimes fewer lines, depending on the overall intensity of the spectrum.

If two lines are within 1 nm of each other, they are shown as a single line. You can find a complete listing of the lines for each element at the NIST (National Institute of Standards and Technology) Atomic Spectra Database.

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Some spectral lines of Cadmium:

Identify the line in the spectrum of Cd with each of the following characteristics: a.

Light with a photon energy of 2.953 × 10-19 J/photon. F (672.18 nm)

b. Light with a photon energy of 185.8 kJ/mol. E (643.85 nm) c.

Light with a frequency of 5.574 × 1014 s-1. C (537.84 nm)

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Post-Lab Key

Experiment 2 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1. What is the difference between an emission spectrum and an absorption spectrum?

An emission spectrum is produced when an atom is excited by an input of energy, which promotes electrons into higher energy shells. As the atom relaxes back into its normal configuration, energy is released as photons at energies (wavelengths) distinct to the atom’s electronic configuration. An absorption spectrum is produced when light passes through a medium that absorbs photons (radiation) of certain wavelengths (energies). Not all wavelengths of the light are absorbed. The spectrum that is produced by those absorbed is called the absorption spectrum.

2. In this experiment you compared your findings to known line spectra for the unknowns.

The lines given were only the few “strongest” lines in the visible spectrum. Explain what is meant by the “strongest” lines and why some lines are stronger than others.

Within the electronic configurations of each atom there are many paths that will allow the atom to “relax” when returning from an excited to relaxed or normal state. Some of those paths are more common than others resulting in a higher number of emissions of that wavelength. The larger the number of emissions of a particular wavelength the stronger the line will appear in its spectrum.

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Experiment 3: Halogen Reactions Instructor Notes and Lab Preparation Chemicals and Equipment: Solid Salts: KI, KBr, KCl, NaI, NaBr, NaCl 0.1 M solutions: KI, KBr, KCl, NaI, NaBr, NaCl Saturated I2 1:1 Clorox/H2O Any brand of household bleach should be ok. Ultra products are 6.25% NaOCl. Many of the procedures I use specify freshly opened bleach, but I think a new jug should last for a one-week lab cycle. Scented products are not good idea.

1.0 M HCl Hexane—place in hood Re-pipette dispenser with aqueous Br2 Waste container labeled “Halogens Waste” Crayons or colored pencils (otional) Setup: 

Br2, Hexane, and the Halogens waste container need to be placed in the hood.



The remaining chemicals can be placed on the front lab bench in an organized fashion.

Prepation: Prepare 0.1 M solutions of each of the reagents.

Saturated I2- add solid iodine crystals and dilute with DI water, excess crystals should remain after shaking.

Br2- Prepare in the hood, add 2.5 mL of concentrated bromine to a 1 L volumetric flask and dilute to the mark with DI water. Transfer solution to a re-pipette container. Adjust the setting on the re-pipette container to dispense 1 mL. Instruction Notes: The Halogens experiment is a personal favorite of mine because it provides one of the few opportunities in a first semester lab to actually teach chemistry. So much of what we teach in the

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first semesters of chemistry are physical rather than chemical based, that it is a nice break to actually discuss the chemical basis for the viability of a reaction. I have always worked with both sodium and potassium salts to help the students understand that based on periodic trends groups have similar physical and chemical properties. If cost is an issue a single salt type could be used.

There are a multitude of teaching points to this lab: 1) Electronegativity should be discussed as a foundation for polarity.

2) Polar and non-polar solvents should be discussed with respect to the solubility of polar and ionic solutes; “like dissolves like.”

3) The density of the two solvents should be discussed in conjunction with the discussion about why the two solvents form layers.

4) Electron-transfer reactions can be reviewed as these reactions are also oxidationreduction processes.

5) Quantitative experimentation versus qualitative experimentation can be discussed.

6) The distinction between observations and conclusions should be mentioned.

The experiment runs as follows: Each student should be assigned a group of salts and salt solutions to test. The first test is simple solubility. A small amount of salt is placed in a test tube containing water or hexane. This test determines the solubility of the halide salts.

The second part of the experiment is the most important. Each aqueous halogen is placed in a test tube and then a layer of hexane added. The halogens then migrate into the hexane layer. It is very important that the students get a good determination of the color of each of the halogens in the hexane layer as this will be how they determine if a reaction does or does not occur in the third part of the experiment.

The colors for each halogen are: chlorine (yellow), bromine (orange), and iodine (pink). I started providing crayons for the students to use a few years ago when I realized how very bad they were at describing colors accurately. I have them draw the test tube and color in what they observe. It is especially helpful during the last part of the experiment.

In the third part of the experiment, each of the halogen solutions are combined with each of the halide solutions. Here is what should occur: Cl2 + 2Br-  2Cl- + Br2

Since chlorine is more electronegative than bromine, a reaction will occur and the color observed in the hexane layer will be orange.

Cl2 + 2I-  2Cl- + I2

Since chlorine is more electronegative than iodine, a reaction will occur and the color observed in the hexane layer will be pink.

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Br2 + Cl-  No Rxn

Since bromine is less electronegative than chlorine, a reaction does not occur and the color observed in the hexane layer will be orange.

Br2 + 2I-  2Br- + I2

Since bromine is more electronegative than iodine, a reaction will occur and the color observed in the hexane layer will be pink.

I2 + Cl-  No Rxn

Since iodine is less electronegative than chlorine, a reaction does not occur and the color observed in the hexane layer will be pink.

I2 + Br-  No Rxn

Since iodine is less electronegative than bromine, a reaction will not occur and the color observed in the hexane layer will be pink.

Potential Problems: Two problems that occur somewhat frequently when running this experiment are the faint coloration of the chlorine solution and improperly mixing the solutions in the test tubes. The chlorine solution is made by using hydrochloric acid in reaction with bleach. If too much bleach is used the yellow color of the chlorine is very faint. Two ways to make the yellow color easier to see are to reduce the amount of bleach used and to use 0.5 mL of hexane instead of 1 mL. The smaller hexane layer will concentrate the color. The second problem will require that you show the students the proper way to mix solutions in a test tube and make sure they mix the solution for more than a few seconds. The iodine solution will become a vibrant hot pink if mixed thoroughly, but can form a brown layer at the interface of the water and hexane layers if not mixed properly.

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Pre-Lab Key

Experiment 3 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) When you go into your laboratory to carry out the halogen displacement reactions, you already know from lecture that the order of reactivity, and electronegativity, is Cl2>Br2>I2, so you are able to predict which results to expect from the various combinations of halogen and halide salt.

Let us suppose, though, that you are a student with a color perception disorder, so that you see the color of the halogens differently from other students. And suppose the halogen samples aren't identified for you, but are given as unknowns, labeled X2, Y2, and Z2 for the halogens, and X−,Y−, and Z− for the corresponding halides. You are to identify the relative reactivity of the unknowns, and hence their identity, by mixing different combinations of halogen and halide, and identifying the halogen product from the color of the hexane layer after extraction.

As a start, you first add the free halogens to your test tube, extract into hexane, and get the following results: Z2 hexane layer is red; X2 hexane layer is yellow; Y2 hexane layer is blue.

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Next, you mix the following combinations of reagents and record the resulting color of the hexane in your test tubes, giving the results shown in the table below.

Using the results of the tests provided, write the balanced reactions for each of the six possible combinations of X2, Y2, Z2, X-, Y-, and Z-.

X2 + 2Y-  2X- + Y2 X2 + 2Z-  NR Y2 + 2X-  NR Y2 + 2Z-  NR Z2 + 2X-  2Z- + X2 Z2 + 2Y-  2Z- + Y2

Reagent Combination

Resulting Color

Y2 + Z−

red

X2 + Y−

blue

Z2 + X−

yellow

2) Complete and balance the following reactions as a "molecular" equation. If no reaction occurs, write NR in place of giving products.

2NaI + F2  2 Na+ + 2F- + I2 (note: Na+ is spectator and may not be in rxn) Complete and balance the following reaction as a "net ionic" equation. If no reaction occurs, write NR in place of giving products.

NaBr + At2  NR

3) For each of the following solvents, determine if they are polar or non-polar.

Explain your conclusions.

a.

b.

Non-Polar, due to symmetry

Polar, electronegative oxygen and lack of symmetry

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Post-Lab Key

Experiment 3 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) This experiment uses a qualitative colorimetric process. What is meant by this statement?

Qualitative means that the quality of the observations are used to determine the outcome of an experiment rather than a measured (numerical) or quantitative value.

2) All of the halogens with the exception of fluorine make strong acids. Using electronegativity, explain why hydrofluoric acid is a weak acid.

Fluorine is the most electronegative atom in the periodic table and as such creates a relatively strong bond to the positive hydrogen atom. This lowers the dissociation constant of HF making it a weak acid. Another consideration is HF’s ability to hydrogen bond helps to stabilize the molecule, making it less likely to dissociate. Finally, F- is a very unstable conjugate base for HF because of its small ionic size. This instability also makes the dissociation of HF less energetically favorable and thus HF is a weak acid.

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Experiment 4 Paper Chromatography: An M&M’s® True Colors Instructor Notes and Lab Preparation Chemicals and Equipment: 1 bag of M&Ms 1 box of toothpicks 1 box of food coloring 1 2 L bottle of chromatography solution 1 waste jar 25 1000 mL beakers (***in the lab already***) 1 box of pencils 1 box of saran wrap 25 rulers 1 bag of rubber bands 2 boxes small filter paper Setup: Place a beaker of M&Ms on the front counter.

Chromatography solution should be kept in the hood to reduce fumes.

Preperation: Chromatography solution: 3:2 solution of ethanol and water Instruction Notes: This is a simple chromatography experiment and is particularly nice because there is no need to create “unknowns” for the procedure. The other nice feature is that you can run the experiment on a time rather than distance basis, so you can allow the students to leave the lab and return in time to remove the chromatogram from the solution. Since most real research is done by a process of setup and then allowed to run for extended periods of time and then removed later, this process gives the students a feel for how many actual experiments would be completed in a research lab. I often think students feel all experimental procedures are completed within three hours; the fact that you might have an experiment run for days or weeks is completely foreign to them. I use this as a teaching point for this lab.

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Potential Problems: The only problem that occurs occasionally is a student over saturating a spot with color or chocolate. To properly spot the chromatogram with sufficient shell color they need to spot, and then allow the spot to completely dry, and then spot again over the same point. The students often are impatient to complete the setup and spot again too soon. This causes the spot to spread and become too large. This produces weaker bands and also tends to have lane bleed over.

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Pre-Lab Key

Experiment 4 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Below is a paper chromatogram comparing four standard dyes (in lanes 1–4, Red, Green, Blue, Yellow, respectively) with six dye mixtures (lanes 5–10). The position of the solvent front is noted by the wavy blue line. (A number of factors can affect the mobility of the solvent, including non-uniformity in the paper and temperature variations within the chamber while the chromatogram is being run. Ideally the front should be horizontal, but this problem illustrates a case where it is not.)

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The relative mobility of substances in thin layer and paper chromatography is recorded as an Rf value. The Rf value is simply the ratio of the distance the substance has moved from the origin (measured to the front of the spot), divided by the distance the solvent front has moved. It is important to record the solvent front position before the chromatogram dries completely. When the solvent front is irregular, as shown in the example above, the solvent front distance should be measured above the spot. Identical substances should have the same Rf value even if they do not line up exactly on a horizontal line.

Measure the Rf values of the the four standard dyes in lanes 1–4 to 2 significant figures.

Rf of the red spot in lane 1: ___0.45__

Rf of the blue spot in lane 3: ___0.55__

Rf of the green spot in lane 2: _0.79__

Rf of the yellow spot in lane 4: __0.20__

Each of the dye mixtures in lanes 5 through 9 contains only one of the four standard dyes.

Identify the standard dye in the mixture by comparing its Rf .

Which of the four standard dyes is contained in the mixture in lane 7? ___Blue______ .

Which of the four standard dyes is contained in the mixture in lane 10? ___Red___ Which of the four standard dyes is contained in the mixture in lane 9? __Yellow___ Which of the four standard dyes is contained in the mixture in lane 5? __Green___ 2) For each of the dyes shown in the background, answer the following questions: a) What is the molecular weight? (Unit=g/mol) b) How many double bonds are present? (Include the S=O double bonds.) c) What is the total number of charges on Blue No. 1? (Do not count the Na+ counterions.) d) How many sulfonate (-SO3−) groups are there? e) Some dyes are classified as "azo" dyes because they have the azo group (-N=N-) as part of the structure. Indicate which of the dyes are azo dyes.

Yellow # 5 Blue #1 Red #40 Red #3

a) 534.36 g/mol; b) 14; c) 3; d) 2; e) Yes a) 792.85 g/mol; b) 22; c) 4; d) 3; e) No a) 496.42 g/mol; b) 13; c) 2; d) 2; e) Yes a) 534.36 g/mol; b) 14; c) 3; d) 2; e) No

3) Using the information from 2) above and your knowledge of polarity, make a prediction of the rankings of the food dyes from largest to smallest Rf values and explain your reasoning. Copy your predictions into your lab notebook for later use.

This ranking will vary based on the student’s reasoning of how much the structure, charges, and mass play into the movement of the dyes. I generally give full credit for any well-reasoned answer.

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Post-Lab Key

Experiment 4 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Using the internet or a chemistry textbook, find at least three other types of chromatography techniques and list them.

Ion-exchange Chromatography, HPLC, Reversed Phase Chromatography, Gas Chromatography, etc.

2) Explain the role of the solvent in this experiment. Why was a 3:2 ethanol water mixture used?

The solvent in chromatography support the molecules to be separated up the chromatogram. The solvent is designed to attract the dyes but not bind tightly to them so that they can “fall out” of the solvent when their mass/structure becomes too much for their attraction to the solvent to overcome.

3) Now that you have completed the experiment, compare your actual results to the predictions you made regarding the Rf values of the food dyes. Explain any discrepancies.

This should be a well thought out explanation and comparison to their predictions from the pre-lab. Each dye should be considered and supporting information, including Rf values, used to support their arguments.

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Experiment 5: Molecular Models Instructor Notes and Lab Preparation Chemicals and Equipment:

 12–14 protractors  6 different molecular models Instruction Notes: This lab was developed after many years of trying to get students to build models for themselves and meeting with heavy resistance. In order to make the lab more investigative, I started building the models myself and asked the students to identify them instead. I use the large scale models normally used for demonstrations in lecture. They are glued together to make them more durable for use from class to class. You can also use styrofoam balls and dowels purchased from the local craft shop if you don’t want to purchase the demonstration model kits. We spray painted the models ourselves to expand on the normal number of central atoms available. I don’t recommend using the smaller model kits since they reduce the number of students that can work on a model at one time. I set out six models at a time and switch out the models from section to section.

Potential Problems: When using the protractors it is important to center them over the central atom. If the protractor is set to the outside of the central atom the angles are not accurate.

The students should also be sure to determine the number of valence electrons based on the atoms in the structure first. This allows for better identification of the ions.

Below is a list of the molecules we build and the color code we use for each element: UNKNOWN A–BeF2 UNKNOWN B–CO2 UNKNOWN C–BH3

UNKNOWN H–SO42UNKNOWN I–NH4+ UNKNOWN J–PF5

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UNKNOWN D–SO3 UNKNOWN E–NO3UNKNOWN F–CO32UNKNOWN G–H3O+

UNKNOWN K–I3UNKNOWN L- XeO2F2 UNKNOWN M- XeOF4 UNKNOWN N- ICl4-

Bonds: Double – red Single – yellow Elements: O – red N – blue C – black S – yellow P – brown Xe – silver Cl – green I – purple F – orange Be – gold B – light blue H– white Lone-pair electrons – clear sticks (Important for ionic structures)

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Pre-Lab Key

Experiment 5 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Make six or more copies of the worksheet provided at the end of the procedure section and bring them with you to lab.

This question is here so that you can give it “teeth” by removing points if they do not remember to bring the forms with them.

2) Sulfur is a Period 3 element and as such can create structures that exceed the octet rule.

Because of this ability, there are 2 resonance structures for the sulfuric acid molecule. a) Draw the two possible structures; b) Calculate the formal charges; and c) Determine the best structure and explain your choice.

O H

O

S

O

H

O

b) H (FC) = 0; S (FC) = +2; O (FC) = 2 at -1, 2 at 0

b) H (FC) = 0; S (FC) = 0; O (FC) = all at 0

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c) The second structure is best as it minimizes the formal charges on the structure.

3) For each of the molecules listed below, complete the following: a. b. c. d. e. f. g.

Calculate the number of valence electrons Draw a good Lewis structure and any resonance structures Calculate the formal charges for each atom Calculate the oxidation numbers for each atom Determine the electron cloud and molecular geometries for the central atom(s) List the bond angles Determine if the structure is polar or non-polar

Molecules: Dichromate ion, Iron II Hydroxide, Disulfur Decafluoride Dichromate Ion a. Valence electrons = 54 b.

c. Formal charges = -1 on each oxygen with 3 lone pairs; 0 on the center oxygen; +2 on each Chromium d. Oxidation Numbers: Cr = +6 ; O = -2 e. Each Cr is tetrahedral f. All 109 degrees for the O to Cr bonds; 180 degrees from Cr-Cr g. Polar Iron II Hydroxide a. Valence Electrons = 14 b.

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c. Formal charges = -1 on each oxygen; 0 on each hydrogen; +2 on iron d. Oxidation Numbers: Fe = +2; O = -2; H = +1 e. N/A ionic bonding f. N/A ionic bonding g. Ionic (extreme polarity) Disulfur Decafluoride a. Valence Electrons = 82 b.

c. Formal charges = 0 on each fluorine; 0 on each sulfur d. Oxidation Numbers: S = +5 ; F = -1 e. Each S is trigonal bipyramidal f. Equatorial F 120 degrees apart; S-F equatorial bonds 90 degrees; SF axial bonds 180 degrees g. Non-polar

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Post-Lab Key

Experiment 5 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) What physical forces create the three-dimensional structure of a molecule?

The electron repulsion is the predominant force that creates a molecule’s structure as the electrons find the lowest energy state that puts their negative charges as distant from each other as possible.

2) Now that you know more about molecular structure, discuss how the structure of a molecule influences its polarity.

Polarity is a combination of the pull of an electronegative atom on the electrons in a structure towards itself with or against the pull of other atoms in the molecule. If a structure contains two or more equally electronegative atoms arranged in a symmetrical manner they will essentially cancel each other out yielding a non-polar molecule.

3) Oxidation Numbers and Formal Charges are often confused by first-year chemistry students. If you were teaching the course, explain how you would differentiate the two types of values so that the concepts would be clear to your students.

Oxidation numbers are based on the electronegativity of an atom and are therefore assigned in a hierarchical manner with more electronegative atoms being assigned the bulk of the electron density or charge. In formal charges, the charges are assigned as if each atom shares the electron density equally with no preference being given to the more electronegative atom. Thus, formal charges are a property of the molecule the atom finds

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itself in and oxidation numbers are a physical property of the atom depending only on its electronegativity ranking.

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Experiment 6: Limiting and Excess Reagents Instructor Notes and Lab Preparation: Chemicals and Equipment: Centrifuges Disposable pipet (and bulbs if glass pipet are used) Grease pencils or labeling tape 1 pair of heat protective gloves Dessicators Dessicant Large filter paper Aqueous Metals Waste jar (in hood) Dessicant waste jar this can be reused indefinitely) 0.5 M Cu(NO3)2 0.5 M KI Methanol

Setup: 

Place one bottle of each reagent on each lab bench.



Each lab bench should have a working centrifuge.



Ovens should be set to 70 degrees.



Place waste jar in hood labeled as “Aqueous Metals Waste.”

Preparation: Dessicant: Mix 2/3 white dessicant and 1/3 blue dessicant 0.5 M Cu(NO3)2: For 1 L, slowly add 93.78g of Cu(NO3)2 and dissolve with DI water. (Supply for ~6 sections of 24 students)

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0.5 M KI: For 1 L, slowly add 83.01g of KI and dissolve with DI water Methanol is stored in the flammable cabinet.

Instruction Notes: There are several nice points to this experiment that can be pointed out to the students as you go along. The first is that a solution does not have to be colorless to be considered “clear.” In all the experiments I performed as a first year-student, we never came across a supernatant that wasn’t colorless. The supernatant in this experiment is a clear green/yellow. You should remind your students that “clear” in this case means without cloudiness, not without color. Another common misconception among young chemists is that equal amounts of reagents yields the maximum amount of product. This probably arises because most reactions used in first-year labs are 1:1 reactions. We developed this lab after several failed attempts to find other combinations that were not 1:1. We tried PbI, which makes a great yellow precipitate, but who really wants that much lead all over the place? We also tried a hydroxide precipitate but it cooked down to oxide in the oven, complicating the issue. The copper iodide reaction is relatively non-toxic and not 1:1, and the precipitate is stable so it is a good reaction for this experiment.

The graph we ask the students to draw appears as a lop-sided pyramid. This also reinforces the notion that the maximum output of product does not mean identical or even amounts of reactants.

Potential Problems: If your lab uses a variety of test tubes (pyrex, non-pyrex, lipped, and not) have your students check to make sure that the test tubes they will be centrifuging will fit in the centrifuge before starting the experiment. They should also make sure the general appearance is the same, i.e., both lipped or not, as this will affect the weight of the tubes and if too disparate, unbalance the centrifuges.

Occasionally, a student will have an allergy to iodine. Most of these allergies are consumption based, but we issue gloves and have the student work at the hood instead of a bench just to be safe.

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Pre-Lab Key

Experiment 6 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

Your laboratory experiment on "Limiting and Excess Reagents" involves the reaction between aqueous solutions of Cu(NO3)2 and KI to form CuI and I2. The CuI precipitates from the reaction as a solid, and the other products (KNO3 and I2) remain in aqueous solution. In the lab you will vary the relative amounts of each reactant, determining the amount of solid product in each case. In this pre-lab you will explore some features of this reaction.

1) This reaction is an oxidation-reduction (ReDox) reaction in which one reactant loses electrons (is oxidized) and another reactant gains electrons (is reduced). To recognize a ReDox reaction you need to check the oxidation states of all the elements in both the reactants and products. If any change their oxidation state, the reaction is a ReDox. An increase in oxidation state is an oxidation; a decrease in oxidation state is a reduction.

(Note: If an element is a reactant or product, converted to or formed from a compound, the reaction must be a ReDox reaction.)

Indicate the oxidation state of the elements in each of the following compounds: Reactants

Products

Cu(NO3)2

KI

CuI

KNO3

I2

Cu +2

K +1

Cu +1

K +1

I 0

N +5

I

-1

N +5

O

-1

I

-2

O -2

Which reactant is oxidized, and which is reduced?

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Iodine is oxidized, copper is reduced.

2) Calculating Molarity: What is the formula weight of CuI? 190.45 g/mol What quantity of CuI is present in 50.0 mL of a 0.50 M solution? 4.76 g What is the formula weight of Cu(NO3)2? 187.57 g/mol What mass of Cu(NO3)2 is present in 225 mL of a 0.30 M solution? 12.66 g What is the formula weight of KI? (Enter unit as g/mol) 166.00 g/mol What mass of KI is needed to add to a 1000 mL volumetric flask to produce a solution that is 0.75 M? 124.5 g What is the formula weight of I2? 253.8 g/mol What is the molarity of a solution produced by dissolving 53 g of I2 in a 250 mL volumetric flask? 0.84 M 3) Limiting Reagent Calculations: You mix 3.5 mL of 0.50 M Cu(NO3)2 with 2.0 mL of 0.50 M KI, and collect and dry the CuI precipitate. Calculate the following: Moles of Cu(NO3)2 used: 1.75 mmol Moles of KI used: 1.00 mmol Moles of CuI precipitate expected from Cu(NO3)2: 1.75 mmol Moles of CuI precipitate expected from KI: 0.50 mmol What mass of CuI do you expect to obtain from this reaction? 0.50 mmol If the actual mass of the precipitate you recover is 0.133 g, what is the percent recovery of the precipitate? 26.6 %

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Post-Lab Key

Experiment 6 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) How would a decrease in the concentration of the reactants affect the moles of reactant?

If the concentration is lowered the number of moles is reduced.

2) Explain the role of limiting and excess reagents in an experiment.

When designing an experiment the scientist always attempts to minimize waste. To do this he/she will calculate the masses or volumes of reactants in such a manner as to produce the amount of product needed without significant amounts of either reactant left over. Knowledge of how reactants will be in excess or limiting, i.e,. the stoichiometry of the reaction, is therefore crucial to a scientist’s ability to conserve chemicals and produce products efficiently.

3) Copper compounds are fairly common. Using the internet or any other reliable source, look up at least five other copper compounds. List the name, chemical formula, and uses of each compound. Be sure to cite your source in acceptable format.

There are a multitude of copper compounds that can be listed. Just make sure the student notes where he/she took the information.

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Experiment 7: Redox Reactions in Voltaic Cells: Construction of a Potential Series Instructor Notes and Lab Preparation Chemicals and Equipment: Foam cups Pieces of white card stock 1 box of large (15 cm circles) coarse filter paper Sheets of sand paper Multi- or volt-meters & alligator clips 1.0 M NH4NO3 1.0 M Cu(NO3)2 1.0 M Mg(NO3)2 1.0 M Pb(NO3)2 1.0 M Fe(NO3)3 1.0 M AL(NO3)3 0.01 M HNO3 Pieces of each metal: Cu; Mg; Fe; Pb; Al Aqueous Metals Waste jar (in hood)

Preparation: Prepare solutions in 0.01 M HNO3. Due to the large quantity of substance required, prepare in advance as it will take several hours to completely dissolve. Heat may be applied to speed up the process. Also, add small quantities at a time for easier dissolving.

1.0 M NH4NO3—80.04 g per 1 liter 1.0 M Cu(NO3)2—187.56 g per 1 liter 1.0 M Mg(NO3)2 —148.31 g per 1 liter 1.0 M Pb(NO3)2—331.00 g per 1 liter

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1.0 M Fe(NO3)3 -241.86 g per 1 liter 1.0 M AL(NO3)3 –213.00 g per 1 liter In class instructions: Do NOT discard the foam cups or metals. Metals MUST be cleaned, dried ,and returned to the correct container. ***Mg and Fe are often confused***

Instruction Notes: This is a very simple lab for the students to perform and only takes about an hour to complete.

We use foam cups to minimize the volume of solution we need to use. 100 mL beakers can be used instead but will require almost 100 mL of solution in each to enable the filter papers to make good contact with the solution. The foam cups only require about 20–30 mL cutting the chemical cost way down. We ask the students to clean and dry the foam cups at the end of the lab so that we can re-use them from class to class to minimize the waste.

Dependent on the number of multi- or voltmeters you have available, you can make this an individual, pairs, or group lab experiment. You should make the assignments of the 10 metal combinations so that all ten are completed and reported to the class as a whole. If you have sufficient numbers of students/groups you can also introduce more statistics into the experiment based on more than one experimental potential value being determined for each of the metal combinations.

Potential Problems: The only problem we have ever experienced with this lab is experimental potentials not matching up well to those predicted by calculation. This is generally caused by a poor connection to the voltmeter or poor contact via one or more of the salt bridges. Make sure that both the metals and alligator clips are clean and making good contact with each other and the solution. Also make sure that the salt bridge is soaked through with solution. If the filter paper is wound too tight, it often appears wet but may be dry on the inside.

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Pre-Lab Key

Experiment 7 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

Following is a diagram of a voltaic cell.

1) Identify the parts of the cell shown above by selecting the appropriate letter from the figure.

__E__ __E__ __E__ __B__ __A__

This is the positive electrode.

Reduction occurs at this electrode.

This electrode is the cathode Anions depart from the salt bridge here.

Oxidation occurs at this electrode.

2) Write the shorthand notation for the following voltaic cell reaction: Fe(s) + Ni2+(aq)  Fe2+(aq) + Ni(s)

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Fe(s)|Fe2+(1M)||Ni2+(1M)|Ni(s) Write the reaction represented by the following voltaic cell notation: Mn(s)|Mn2+(1M)||Pb2+(1M)|Pb(s)

Mn(s) + Pb2+(aq) ->Mn2+(aq) + Pb(s) 3) The list below shows the reduction potentials for ten different metal ions.

Calculate ΔE for the ReDox reaction: Ni(s) + Mn2+(aq)  Ni2+(aq) + Mn(s)

ΔE = -0.93 Will this reaction occur spontaneously? No Calculate ΔE for the ReDox reaction: Co(s) + Sn2+(aq)  Co2+(aq) + Sn(s)

ΔE = +0.14 Will this reaction occur spontaneously? Yes

Half-Reaction

Eo(V)

Cu2+(aq) + 2e−  Cu(s) +0.34 Pb2+(aq) + 2e−  Pb(s)

-0.13

2+

−

-0.14

2+

−

-0.25

2+

−

Sn (aq) + 2e  Sn(s) Ni (aq) + 2e  Ni(s)

Co (aq) + 2e  Co(s) -0.28 Cd2+(aq) + 2e−  Cd(s) -0.40 Fe2+(aq) + 2e−  Fe(s)

-0.44

Zn2+(aq) + 2e−  Zn(s) -0.76 Mn2+(aq) + 2e−  Mn(s) -1.18 Mg2+(aq) + 2e−  Mg(s) -2.37

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Post-Lab Key Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) What is a reactivity series and how is it related to this experiment?

A reactivity series of metals is a list of metals ranked in order of decreasing reactivity to displace hydrogen gas from water and acid solutions. It can also be used to predict which metals will displace other metals in aqueous solutions. For this experiment it helps to determine which metal will accept electrons and which will donate electrons by determining the cathode and anode for each cell.

2) Create a potential series ranking all of the metals tested by the class arranged from strongest oxidizing agent to weakest oxidizing agent.

This should be a list reverse order of the reduction potentials shown above for the metals used in your lab.

3) Explain your reasoning behind the rankings you list in problem 2.

The identification of each metal as either the cathode or anode in each cell created develops the direction of flow for the electrons. The metal that is assigned as the cathode in all the combinations will be the strongest oxidizing agent, the metal that is assigned as the cathode in all but one combination is the second strongest and so forth.

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Experiment 8: Reactions in Aqueous Solutions: Strong Acids and Bases Instructor Notes and Lab Preparation Chemicals and Equipment:           

0.5 M HCl 1 liter, dyed red 0.5 M NaOH 1 liter, dyed blue 1 mL graduated pipette (plastic) 3-way rubber bulb 25 mL volumetric flasks 50 mL beaker as needed hotplates pH meter power supply calibration instructions buffer solutions

Setup: The 0.5 M solutions are purchased in standardized form and a few drops of food color are added to make it easier for the students to differentiate the solutions. A hotplate and a pH meter should be placed out for every two students.

Instruction Notes: Points to Stress This is a relatively straight-forward lab that allows students to observe the reaction of a strong acid and base, and how the concentrations of each relate to their pH. We also use the lab to point out the predictability of the production of salt from the acid/base reaction. This reinforces the limiting and excess reagents lab just completed. The lab also introduces the use of pipettes, serial dilution, pH meters, and calibration.

Potential Problems The only problem we have had with the experiment is the pH of the DI water not being 7 or even very close to 7. This complicates the calculation of pH and makes for much greater errors between the student’s calculated and experimental values.

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Pre-Lab Key

Experiment 8 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Constant boiling HCl has a concentration of 11.6 M. Your laboratory assistant is going to dilute this acid for you to create a stock solution of HCl for you to use in this experiment.

What volume of the 11.6 M HCl must the assistant add to a one-liter volumetric flask so that the concentration of the stock solution will be 0.62 M after dilution to the mark with water?

53.4 mL What is the pH of this stock solution? (Remember the significant figure rules for logarithms!)

pH = 0.21 You will now use this 0.62 M HCl solution to carry out a series of dilutions, measuring the pH of each diluted sample. Calculate the [Cl-] and [H+] concentrations, and the pH of each sample. (Remember the significant figure rules for logarithms!)

Sample A will be produced by pipetting 1.00 mL of the stock solution into a 25 mL volumetric flask, and diluting to volume with distilled water.

[Cl-]= 0.025M [H+]=0.025M pH= 1.61 Sample B will be produced by pipetting 1.00 mL of sample A into a 25 mL volumetric flask, and diluting to volume with distilled water.

[Cl-]= 9.9e-4M [H+]= 9.9e-4M pH= 3.00

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Sample C will be produced by pipetting 1.00 mL of sample B into a 25 mL volumetric, and diluting to volume with distilled water.

[Cl-]= 4.0e-5M [H+]= 4.0e-5M pH= 4.40 Sample D will be produced by pipetting 1.00 mL of sample C into a 25 mL volumetric, and diluting to volume with distilled water.

[Cl-]= 1.6e-6M [H+]= 1.6e-6M pH= 5.80 Sample E will be produced by pipetting 1.00 mL of sample D into a 25 mL volumetric, and diluting to volume with distilled water.

[Cl-]= 6.3e-8M [H+]= 1.4e-7M pH= 6.86 Sample F will be produced by pipetting 1.00 mL of sample E into a 25 mL volumetric, and diluting to volume with distilled water.

[Cl-]= 2.5e-9M [H+]= 1.0e-7M pH= 6.99 2) Your laboratory assistant is going to prepare a stock solution of sodium hydroxide (NaOH) for you. What mass of NaOH must the assistant add to a one-liter volumetric flask so that the stock solution has a concentration of 0.35 M after dilution to the mark with water?

14.0 g What is the pH of this stock solution? (Remember the significant figure rules for logarithms!)

pH= 13.54 You will now use this 0.35 M NaOH solution to carry out a series of dilutions in the same manner as in problem 1, measuring the pH of each diluted sample. Calculate the [Na+] and [OH-] concentrations and the pH of each sample A–F. Sample A will be produced by pipetting 1.00 mL of the stock solution into a 25 mL volumetric, and diluting to volume with distilled water.

[Na+]= 1.4e-2M [OH-]= 1.4e-2M

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pH= 12.15 Sample B will be produced by pipetting 1.00 mL of sample A into a 25 mL volumetric, and diluting to volume with distilled water.

[Na+]= 5.6e-4M [OH-]= 5.6e-4M pH= 10.75 Sample C will be produced by pipetting 1.00 mL of sample B into a 25 mL volumetric, and diluting to volume with distilled water.

[Na+]= 2.2e-5M [OH-]= 2.2e-5M pH= 9.35 Sample D will be produced by pipetting 1.00 mL of sample C into a 25 mL volumetric, and diluting to volume with distilled water.

[Na+]= 9.0e-7M [OH-]= 9.1e-7M pH= 7.96 Sample E will be produced by pipetting 1.00 mL of sample D into a 25 mL volumetric, and diluting to volume with distilled water.

[Na+]= 3.6e-8M [OH-]= 1.2e-7M pH= 7.08 Sample F will be produced by pipetting 1.00 mL of sample E into a 25 mL volumetric, and diluting to volume with distilled water.

[Na+]= 1.4e-9 [OH-]= 1.0e-7 pH= 7.00

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Post-Lab Key

Experiment 8 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Explain how this experiment can be related to limiting and excess reagent problems.

The determination of the salt yield is based on the amount of reagents that are added.

2) The acid and base used in this experiment are monoprotic. How would the experimental results vary if a diprotic acid had been used instead?

If the acid were diprotic then it would take twice as much base to neutralize the acid since the molar ratio between H:OH would be 2:1. If the same volumes were used then half as much product would be produced since the base would run out.

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Experiment 9: Calorimetry & Hess’s Law Instructor Notes and Lab Preparation Chemicals and Equipment: Styrofoam cups Thermometers Cardboard lids with hole in center Bottle of 1.0 M HCL Container of Mg metal 2.5 cm strips cut from Mg ribbon Bottle of MgO Bunsen burners Vacuum tubing Pieces of wire gauze Small weighing boats Aqueous metals waste container (in hood) **adjust pH before disposing**

Setup: Place 1 styrofoam cup and 1 piece of cardboard per pair of students Place the aqueous waste container in the hood The rest of the supplies should be set up on the front counter

Preparation: 1.0 M HCL—83.3mL of concentrated HCL per 2L of 1.0 M HCL solution prepared **** Prepare a large quantity****

Instruction Notes: After years of running the typical coffee cup calorimetry, solving for the heat capacity of a metal, I decided that the experiment didn’t cover enough of the material discussed in the thermodynamic chapter in the textbook. The process of solving for heats of reaction and use of Hess’s Law to determine an otherwise undeterminable heat of reaction provides a more complete coverage of the thermodynamics discussed in the book and also gives the students a realistic use for Hess’s Law.

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Points to Stress: The first part of the experiment asks the students to determine the calorimeter constant. We currently use a single 12 ounce styrofoam coffee cup for the calorimetry runs. (We tested the experiment with both single and doubled cups with no appreciable difference in results.) The first part of the experiment is repeated because the students often have difficulty with the technique. This leads to differences in the results of the two experimental runs. It is important to discuss with the students whether they should use the results of both experimental runs or to favor one run over the other and how to justify this decision in their conclusion.

Potential Problems: The only difficulty with the experiment occurs when the students do not stir continuously when adding the MgO powder. If you use stir bars, this problem can be avoided. When you find that you do not have enough stir plates to use this fix, we stress the need for one student to “swirl” the cup continuously while adding the powder.

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Pre-Lab Key

Experiment 9 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Use the graph of Temperature (oC) versus Time (s) provided to answer the following questions: o

What is the initial temperature? 23.9 C o

What is the final temperature? 39.2 C The Cp for the calorimeter (including the HCl solution and the Mg) was calculated to be 0.226 kJ/°C. For the temperature curve shown above, 0.175 g Mg was reacted with HCl. What is the ΔH of the reaction? 480 kJ/mol 2) To calibrate your calorimeter cup, you first put 53 mL of cold water in the cup, and measure its temperature to be 21.1 °C. You then pour 41 mL of hot water, temperature = 53.7 °C, into the cup and measure the temperature every thirty seconds over a 10 minute period. You extrapolate this “cooling curve” back to the time of addition and find that the “final temperature” after mixing is 34.4 °C. What is the heat change of the hot water, qHW? -3310.8 J (Assume the density of the water is 1.00 g/mL, and remember that the specific heat of water is 4.184 J/g-K or J/g- °C.)

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What is the heat change of the cold water, qCW? 2949.3 J What is the heat change of the calorimeter cup, qcup? 362 J What is the heat capacity of the calorimeter cup, Ccup? 27 J/K 3) You would like to measure the heat of neutralization of an acid with a base. You mix 225 mL each of 0.77 M HCl and 0.77 M NaOH, both at a temperature of 18.7 °C in a calorimeter cup equilibrated to that same temperature. After following the temperature change for 10 minutes and extrapolating it back to the time of addition, you find that the “final temperature” after mixing was 23.7 °C. Previously you measured the heat capacity of the calorimeter cup (Ccup) to be 39 J/ °C. Assuming that the density of the two solutions is 1.00 g/mL, and that their specific heats are the same as water, 4.184 J/g-°C, what is the heat absorbed by the two solutions, qsoln?

-9414 J What is the heat absorbed by the calorimeter cup, qcup? 195 J What is the heat change for the neutralization reaction, qrxn?

-9609 J What is the ΔH for the neutralization? -55.5 kJ/mol Write the net ionic equation for this neutralization reaction, including the states of matter (i.e., H+(aq)) for the reactants and products.

H+(aq) + OH-(aq)  H2O(l) 4) Given the following data: SO2(g) → S(s) + O2(g)

∆H = +296.8 kJ

CO2(g) → C(s) + O2(g)

∆H = +393.5 kJ

C(s) + 2S(s) → CS2(l)

∆H = +89.4 kJ

Find the ΔH of the following reaction: CO2(g) + 2SO2(g) → CS2(l) + 3O2(g)

ΔH = 1076.5 kJ .

Post-Lab Key

Experiment 9 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) What is the difference between the heat (q) of a reaction and the reaction enthalpy (HRXN) ?

The heat of a reaction is simply the energy created by a temperature difference. The enthalpy is the sum of all internal energy differences for a reaction.

2) Why is Tfinal taken from extrapolating back to the beginning of the experimental run rather than from the last temperature reading taken?

Tfinal is the theoretical temperature that would be reached in a reaction were all substances to reach thermal equilibrium instantaneously. The extrapolation based on the temperature decay line allows for the estimation of what this temperature would be. The last temperature reading taken is at the end of the temperature decay and does not directly indicate the temperature of instantaneous equilibrium.

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Experiment 10: Gas Laws and Air Bags Instructor Notes and Lab Preparation Chemicals and Equipment: 1 bottle of 6 M acetic acid 1 bottle of solid sodium bicarbonate 1 box each of 3 different types of zipper-type plastic bags: sandwich, freezer, storage Small weighing boats

Preparation: 6 M acetic acid—Fill the large carboy labeled 6 M acetic acid partially with DI water. In the hood, add 6.90 L of concentrated acetic acid. Dilute to the 20 L mark with DI water and mix thoroughly.

Instruction Notes: The concept of this lab is very simple and yet it combines gas law, stoichiometry, and limiting and excess reagents. The lab is short, normally taking no more than 1 hour to complete. A successful airbag should have a full pillow-like appearance and there should be no particulate residue from the bicarbonate left unreacted. Although we require that the bags be tested in the hood, the reaction is not really fast enough to “explode.”

The heavier zipper-style bags with a double zipper work better than the lighter single zipper bags. They are slightly more expensive but are far less likely to leak. To save on expense, the bags can be used more than once as long as they are well rinsed and dried in between classes.

Points to Stress: The comparison of the “airbag” the students are building in the laboratory to the ones found in cars is a very valid one. The requirement for an exact calculation of the materials needed to completely fill but not overfill the bag is an important point to make.

Potential Problems: The only problem in the experiment occurs when the bag leaks. The reason we started determining the volume of the bags by filling them with water instead of simply converting the measurements given on the box is to test the bag for leaks, i.e., if the water leaks so will the gas.

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Pre-Lab Key

Experiment 10 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Airbags in cars are devices that inflate very rapidly (in about 30 milliseconds) when there is a collision, creating a cushion that deflates slowly and cushions the impact of the collision on the driver or the passenger. The bag is inflated by a gas which is generated by the rapid heating of sodium azide (NaN3) to about 300 °C, causing it to decompose into sodium metal and nitrogen gas.

Complete and balance the equation for this reaction, indicating the states of the reactant (NaN3(s)) and products (Na(s) and N2(g)).

2NaN3(s) 2Na(s) + 3N2(g) Enough nitrogen must be generated in the bag to create a total pressure of 3.50 atm, which then drops as the bag slowly deflates. Assuming the volume of the bag is 65.0 L and the temperature in the car is 28 °C, calculate the molar quantity of N2 that must be generated.

9.21 mol What molar quantity of sodium azide is required to generate this much nitrogen?

6.14 mol What mass of sodium azide is required?

400 g

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The sodium metal produced in the reaction is highly reactive and potentially explosive, so it is removed in a reaction with excess KNO3(s) that produces K2O(s) and Na2O(s), and additional nitrogen. Write a balanced equation for this process, showing the states of the reactants and products.

10 Na(s) + 2 KNO3(s)  K2O(s) + 5 Na2O(s) + N2(g) 2) In this laboratory experiment you are to generate sufficient carbon dioxide to inflate a plastic ziploc bag. You first must carry out some calculations to determine the quantity of reagents needed.

Carbon dioxide will be generated by reacting solid sodium bicarbonate (NaHCO3) with a solution of acetic acid (CH3COOH) as follows: NaHCO3(s) + CH3COOH(aq)  H2CO3(aq) + CH3COO−(aq) + Na+(aq) H2CO3(aq)  H2O(l) + CO2(g) You measure the plastic ziploc bag dimensions to be 6.90 inches by 7.15 inches.

Assuming the average thickness of the bag when inflated will be 1.45 inches, what do you calculate the volume of the bag to be?

71.5 in3 = 1172 mL Of course, this volume is only approximate as the width of the inflated bag is not uniform. To get a better value for the volume, you fill the bag with water, then measure the volume of the water to be 1174.5 mL. Now you measure room temperature to be 24 °C and the barometer on the wall gives a reading of 753 torr for the atmospheric pressure. Using the ideal gas law, calculate the number of moles of CO2 that will be required to fill the bag at this temperature and pressure.

0.0477 mol Now inspect the reactions above generating CO2 and calculate the molar quantity of NaHCO3 and CH3COOH needed.

Quantity of NaHCO3: 0.0477 mol Quantity of CH3COOH: 0.0477 mol What mass of NaHCO3 will be needed? 0.401 g You have a solution of 5.00 M acetic acid. What volume of this solution will be needed?

9.54 mL

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Post-Lab Key

Experiment 10 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Is the gas produced in this experiment really an “ideal” gas? Explain.

An ideal gas is one that completely obeys the ideal gas law. Most gases deviate in some fashion from the law in its absolute form due to intermolecular forces, and thus the gas we are working with is most likely not an “ideal” gas.

2) The gas production in this experiment is fairly slow, but the production in a car airbag is quite rapid. What is the reason for the difference in the two reactions?

The sodium azide in an airbag is ignited and thus a combustion reaction takes place. The input of energy creates a much faster reaction than the simple mixing that takes place during our experiment.

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Experiment 11: Intermolecular Forces and the Triple Point of CO2 Instructor Notes and Lab Preparation Chemicals and Equipment: Liquids: 1 liter bottles of each, labeled only as unknowns A Methanol B Acetone C Water D 2-propanol E Hexane (1) box of small (4.25”) 1 m filter paper (1) bag of small rubber bands Buret or ring stands and clamps Thermometers glass, preferably non-Hg (not digital temperature probes) Clear plastic cups Thick-stemmed pipettes Thin-stemmed pipettes Pliers Bottle of blue food coloring Metric rulers Hot glue gun with glue sticks Bucket of dry ice ~ 1–1.5 lbs for a class of 28 students String

Preparation: Place equipment and chemicals on front lab bench.

Place 5 waste containers in the hood, labeled A, B, C, D, and E Waste in jars C and D can be disposed of in the sink drain with running water. Waste from jars A, B, and E goes into an organic waste jar.

Instruction Notes: This is currently the first wet lab (emphasis on wet) of our second semester of general chemistry at Florida State University. It is a good lab to start with because it requires a minimal amount of .

lab skills (no new techniques) and the “exploding” cups of water refresh the students’ memories regarding the need for goggles in the lab.

The first part of the experiment is about as exciting as watching paint dry, but does reinforce the concepts of intermolecular forces and their effect on physical properties quite effectively. The second part of the experiment dealing with the triple point of CO2 is a student favorite (mostly because of the exploding cups of water).

The setup for this experiment is easy since there are no solutions to make and the equipment is pretty standard. I have found that the buret stands and clamps are best for supporting the thermometers. It also allows simultaneous determination of the temperature change for two solvents at one time. Students can work as groups or benches to determine the T values for the solvent unknowns. This saves on chemical cost. It is advisable to have the two parts of the experiment running at opposite ends of the bench or table due to the possible splash of water.

Also, make sure the students keep their lab notebooks and other personal items clear of the “impact” area. I generally have students work in teams of four with two students working on the first part of the lab while the other two complete part two. At some point, the pairs switch to complete all of the experiment. This also cuts the need for thermometers and pliers in half.

The dry ice needed for the experiment is not as great as it might seem. One or two ~2 cm pellets are sufficient to fill the pipette bulb 1/3 full. We use about 30 pounds during a week for 20 sections of ~28 students. We collect a large amount (~10 lbs at a time) and store it all day in a sealed cooler for easy access.

I suggest purchasing glue guns and pliers from a store such as Walmart or K-mart. The mini glue guns and glue sticks are generally less than $5 and you only need 1 or 2 for a class of 28 students. We actually purchased the pliers at a local flea market for $1 per pair.

Points to stress: The correlation between the strength of the intermolecular forces and the magnitude of the observed change in temperature is a difficult concept for the students to grasp. I have done the best I can to explain it in the background, but a discussion in lab is probably wise.

Students often rush through the observations of the CO2 phase change, but it should be mentioned that very few people get the opportunity to observe liquid CO2.

Potential Problems: There are no known difficulties with either part of this experiment.

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Pre-Lab Key

Experiment 11 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

Copy the information gathered in part 1 and 2 of this pre-lab into your lab notebook for later use.

1) For the following liquids: methanol, isopropanol, n-hexane, acetone and water, (a) write the molecular formula; (b) draw a Lewis structure for each compound; and (c) calculate the molecular weight of each compound.

methanol : CH3OH

MW = 32 g/mol H H

C

O

H

H

2-propanol : (CH3)2CHOH

MW = 60 g/mol H H

C

H

H

C

O

H

C

H

H

.

H

n-hexane : C6H14 MW = 86 g/mol H

H

H

H

H

H

H

C

C

C

C

C

C

H

H

H

H

H

H

H

acetone : C3H4O MW = 58 g/mol O H

H

C C

C

H

water : H2O

H

H

H

MW = 18 g/mol O H

H

2) The relative strength of the intermolecular forces depends on the size (i.e., MW), and structure of the molecules, and include in rough order of strength: ion-ion, ion-dipole, hydrogen bonding, dipole-dipole (polar), dipole-induced dipole, and induced dipole-induced dipole (dispersion) forces. The latter increases with molecular weight and can sometimes become the most predominant force. There are no ions in these liquids, and all the molecules are alike, leaving primarily hydrogen bonding, dipole-dipole, and disperson forces to consider.

What are the strongest intermolecular forces in acetone? Dipole-Dipole What are the strongest intermolecular forces in methanol? H-Bonding What are the strongest intermolecular forces in 2-propanol? H-Bonding What are the strongest intermolecular forces in water? H-Bonding What are the strongest intermolecular forces in hexane? Dispersion 3) The figure at the right shows your experimental setup for estimating the triple point of CO2. Carefully mark off your capillary tube in ten 1-cm segments, and note the position of the drop of dye solution when the capillary is inserted in the open plastic tube. When the top of the plastic tube is sealed, the pressure inside will increase as the CO2 sublimes, and the drop of dye will be forced up the capillary, compressing the air inside it. Note the position of the drop of dye when liquid CO2 starts to form. The pressure inside the capillary should then correspond to the pressure at the triple point of CO2.

You can calculate the pressure from a modification of Boyle's law, P1V1=P2V2.

The volume of the capillary V = hA where h is the height and A is the cross.

sectional area, which is constant. Boyle's law therefore reduces to P1h1=P2h2 where P1 is the pressure in the open tube, h1 is the distance of the drop from the top of the capillary in the open tube, h2 is the distance of the drop from the top of the capillary in the closed tube, and P2 is the pressure in the closed tube—the pressure at the triple point.

What are the heights of the air columns in the two capillaries? (Measure from the top of the dye droplet.)

Height of column in the open tube (h1): 10 cm Height of column in the closed tube (h2): 2 cm Calculate the pressure in the closed tube, assuming the pressure in the open tube is 1.0 atm.

Pressure in the closed tube: 4.5 atm

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Post-Lab Key

Experiment 11 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) How is the change in temperature we observe as the various solvents evaporate related to their Hvap? Explain this relationship.

The Hvap is a physical quantity that measures the amount of energy required to vaporize the liquid. Evaporation is an endothermic process that results in a drop in temperature of the surroundings. The compound requiring the least amount of energy (smallest Hvap) and thus the weakest intermolecular forces will exhibit the largest temperature drop. This occurs because as the molecules evaporate they carry heat away from the thermometer, which registers each loss as a drop in temperature. Molecules with stronger intermolecular forces require a larger amount of energy to accumulate before they can evaporate so their removal of heat from the thermometer is a much slower process. The length of time between evaporation events allows the thermometer to recover heat from the surrounding air and thus we do not observe as great a drop in temperature for these molecules.

2) In our use of the microgauge to determine the triple point pressure of CO2, we only measure the height change (distance the dye moves) in the gauge, not the actual volume. Why does this approximation not ruin our results?

From the rules of multiplication and division, we know that any equivalent value found on two sides of an equality can be ignored in the calculation as it cancels out. Since the diameter of the gauge is a constant, it is an equivalent value that appears on both sides of the P1V1 = P2V2 equation and thus it cancels out and can be left out of the overall calculation. .

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Experiment 12: The Purification of Water Instructor Notes and Lab Preparation Chemicals and Equipment: pH probes and meters (students work individually, but share a pH meter) pH 4 buffer pH 7 buffer dropping bottles of 3 M HCl 1 L bottles of 10 ppm phosphate stock solution dropping bottles of ammonium molybdate reagent dropping bottle of stannous chloride dropping bottles of sodium hypochlorite reagent 1 jar of lime powder (CaO) 1 jar of aluminum sulfate powder (Al2(SO4)3) 1 jar of sand 1 jar of activated charcoal (fish tank rocks) 1 box of 0.1 m filter paper Buchner funnels 250 mL filtering flasks vacuum tubing ring stands clamps spectrometers cuvettes 10 mL glass pipettes pipette pumps or bulbs 25 mL volumetric flasks

.

Preparation: Ammonium molybdate reagent: Prep for 1 L: In a 1 L volumetric flask, dilute 280 mL of concentrated sulfuric acid by adding it to 400 mL of DI. water. Meanwhile, dissolve 25 g of ammonium molybdate (NH4)6Mo7O24.4H2O in about 150mL of D.I. water. When the acid mixture has cooled down, combine it with the molybdate and dilute the entire solution to 1 L with DI water.

For 100 mL: 2.5 g of ammonium molybdate and 28 mL of concentrated sulfuric acid; prepared as stated.

0.1 M stannous chloride: In a 1 L beaker, dissolve 12.0 g of tin(II) chloride dihydrate in 500 mL of glycerol.

Warm the solution on a hotplate to speed up the process.

10 ppm stock phosphate solution: Dissolve 0.07 g of K2HPO4 in 7 L of DI water.

Sodium hypochlorite reagent: Dissolve 250 mL of non-scented commercial grade (5.25%) bleach in a 1 L volumetric flask and dilute with DI to the mark.

Phenolphthalein indicator: Dissolve 0.5 g of phenolphthalein in 50 mL ethanol and 50 mL of DI water.

Setup: On every bench place: A Spectronic 20® (or similar spectrometer), a 25 mL volumetric flask, (2) pH 4 buffers, (2) pH 7 buffers, 1 box of Kimwipes®.

Place an aqueous metals waste container in the hood. *** Paper towels and Kimwipes need to be replenished often*** ***pH buffers need to be replaced after every two classes***

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Instruction Notes: This is a lab that was created to emulate a typical water quality laboratory and purification plant, as described in Sec 14.13, 4th edition of General Chemistry by McMurry and Fay. It also serves to introduce the spectrophotometer for a purpose, rather than just as an exercise. This is a fulllength lab intended to run for approximately 2 ½ to 3 hours. Students should work in groups for the creation of the standard solutions and when using the spectrophotometer. This will reduce the amount of standard solution and the number of spectrophotometers needed for a class. They may also share a pH meter since its use is limited. pH paper can be used instead if you prefer.

Because this lab has several parts and is longer than other labs, I generally set time goals for completion of each section to minimize student “dawdling.” If they can complete the serial dilutions and initial absorbance measurements by the end of one hour, they will be able to complete the purification process and re-test the phosphate concentration within the 3 hour limit.

Points to stress: The first part of this lab requires the students to go into the field and collect their own water samples. Even though it is stated in the manual, you should re-emphasize not using detergents to clean the sample bottle. We ask that students bring the collection bottles they will be using into the lab a week early in order to use a mild acid wash to clean the bottle(s). This will help them remember the need for sample collection the week before the lab is run. This need for advance preparation which is so typical of actual research is an important point to stress.

There are many different techniques used in this lab that should be carefully presented to the students: 1) the calibration and use of a pH meter; 2) the proper use of a spectrophotometer; 3) the setup and use of a filter funnel system.

True serial dilution is a process that can be invaluable in an actual research lab due to the everrising cost of chemicals, but was very seldom taught in laboratories I had seen previously. As a public university, we are also trying to lower the cost per student for our labs. When we changed the process from a simple creation of solutions to a serial dilution process, we reduced the amount of stock solution needed by 65%. This point should me made to the students as well.

The steps of the purification process are based on the typical process used by water treatment plants.

There are three principal stages in water purification: 1. Primary treatment–Collecting and screening including pumping from rivers and initial storage.

2. Secondary treatment–Removal of fine solids and the majority of contaminants using filters, coagulation, flocculation, and membranes.

3. Tertiary treatment–pH adjustment, carbon treatment to remove taste and smells, disinfection, and temporary storage to allow the disinfecting agent to work.

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The EPA has a nice website: http://www.epa.gov/safewater/kids/watertreatmentplant/ that shows the typical process used to treat waste water, and will demonstrate the similarities between the real process and the procedure the students are completing in the lab.

Potential Problems: We have had some problems with our distilled water blank in combination with the coloring agent, probably due to some contamination in the water.

Students often bring in water from “duck ponds.” Unfortunately, some municipalities treat these duck ponds with antibacterial agents and other chemicals to reduce odor and clear the water.

These chemicals can produce unknown reactions when processed in the lab. (We had one water sample that foamed like “elephant toothpaste” for 20 minutes when we added the CaO and Al powder). Just try to be aware of any water sources in your area that may be “quality” controlled and warn your students away from them.

The only other consistent source of error for the students is in the filtration process. They tend to get in a hurry and swamp the filter paper. It is important to show the students the proper way to use the filtration system.

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Pre-Lab Key

Experiment 12 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Calculate the [PO43-] concentrations corresponding to the following ppm values: Concentration = 0.41 ppm; Concentration = 275 ppm; Concentration = 21 ppm;

[PO43-] = 4.3e-6 M [PO43-] = 2.9e-3 M [PO43-] = 2.2e-4 M

Calculate the ppm values corresponding to the following [PO43-] concentrations: [PO43-] = 4.56 × 10-6 M; Concentration = 0.433 ppm [PO43-] = 3.88 × 10-3 M; Concentration = 369 ppm [PO43-] = 7.66 × 10-5 M; Concentration = 7.28 ppm 2) Use the following table of standard concentrations and absorbances for the phosphate ion to create a calibration graph. What are the slope and intercept of the calibration curve?

Phosphate Analysis What is the concentration of phosphate in a solution that has an absorbance of 0.708 AU? 6.54 ppm One of your samples shows a reading of 1.207 AU, which is beyond the range of the calibration curve. The lab is over and it is too late to redo the calibration over a broader range, or to dilute the sample and determine it a second time. Assume that Beer's Law still holds up to the value of the sample, and calculate the concentration using the equation for the calibration line. 12.0 ppm (Path length = 1.0 cm)

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Concentration (ppm)

Absorbance (AU)

0

0.109

2

0.288

4

0.476

6

0.651

8 10

0.847 0.997

3) Phosphate standards used to generate the calibration curve are usually made by first preparing a stock solution of known concentration, then diluting it to the required concentration.

First prepare a stock solution that is 1.00 x 10-2 M in phosphate ion. How many g of Na3PO4 would you add to a 250 mL volumetric flask to to prepare this solution?

Mass Na3PO4: 0.410 g What is this concentration in ppm of phosphate? (That is as PO43−, not Na3PO4).

PO43− concentration: 950. ppm What volume of this ppm solution would you dilute with water into a 1.000 L volumetric flask to make a solution of 10.0 ppm? 10.5 mL This much of the work will be done for you by the laboratory assistant. You will be given a stock solution of 10.0 ppm phosphate to use in calibrating your phosphate assay. To preserve the ammonium molybdate reagent, you will create the calibrating standards by doing a serial dilution, using an aliquot of each diluted solution to create the next dilution.

After adding 1 mL of ammonium molybdate reagent and 2 drops of stannous chloride to 25 mL of the 10.0 ppm solution to develop the color, use a 20 mL aliquot to dilute with 5 mL of water to make 25 mL of a solution of 8.0 ppm.

Using the dilution equation: V1 × Conc1 = V2 × Conc2 and rearranging: 20mL  25mL

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8 ppm 10 ppm

To create 25 mL of a 6.0 ppm solution, what volume of the 8.0 ppm solution and what volume of water should be mixed? 18.75 mL solution and 6.25 mL water To create 25 mL of a 4.0 ppm solution, what volume of the 6.0 ppm solution and what volume of water should be mixed? 16.67 mL solution and 8.33 mL water To create 25 mL of a 2.0 ppm solution, what volume of the 4.0 ppm solution and what volume of water should be mixed? 12.50 mL solution and 12.50 mL water

4) What is the molar concentration of each of your calibration solutions?

Concentration = 10 ppm; [PO43-] = 1.05E-04 M Concentration = 8 ppm; [PO43-] = 8.42E-05 M Concentration = 6 ppm; [PO43-] = 6.32E-05 M Concentration = 4 ppm; [PO43-] = 4.21E-05 M Concentration = 2 ppm; [PO43-] = 2.11E-05 M

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Post-Lab Key

Experiment 12 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Why should we care about the amount of PO43- found in a natural water source?

High concentrations of PO43- are indicative of human effects on an ecosystem (i.e., pollution). Generally, phosphorous is a limiting reagent in an ecosystem.

In this way, it regulates the amount of growth various species of plants and animals can achieve. Addition of excess phosphorous into the ecosystem can lead to explosive growth (especially in plants), and this overgrowth can endanger the balance of the ecosystem.

1) If you really wanted to be sure your water was drinkable, what further purification step(s) would you take?

Distillation predominantly, but allow for other examples.

2) Blue light has a wavelength between 450–495 nm. Why don't we measure the absorbance of a blue solution within this wavelength range?

The solution appears blue because it transmits light in the blue wavelength region, but it absorbs light at other wavelengths. Therefore very little absorption would be seen within the region of blue light.

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Experiment 13: Kinetics: The Iodine Clock Instructor Notes and Lab Preparation Chemicals and Equipment: 1 bottle of 0.2 M KI 1 bottle of 0.05 M (NH4)2S2O8 1 bottle of 0.00175 M Na2S2O3 Dropping bottles of 3% starch solution Hotplates Cooler of ice 1 box of salt Waste container in hood

Preparation: For an average section of up to 28 students, prepare: 0.2 M KI – 66.40 g of KI in a 2 L volumetric flask, dilute to the mark with DI water 0.05M (NH4)2S2O8 – add 22.82 g in a 2 L volumetric flask, dilute to the mark with DI water 0.00175M Na2S2O3 – add 0.553 g of anhydrous Na2S2O3 to a 2 L volumetric flask, dilute to the mark with DI water 3% Starch – in a 2 L beaker bring 1000 mL of DI water to a boil; weigh 30 g of starch in a large weighing boat. Dissolve in the weighing boat with DI water. Once the water is boiling, SLOWLY add the liquid starch, while stirring. Place a large watch glass on top of beaker.

Continue heating and stirring to dissolve. Be careful not to over boil. If not prepared fresh, keep refrigerated. Check during the week simply by combining a few drops of each reagent and observing the dark color.

Setup: Place a hotplate on each bench.

Place the chemicals on the front counter.

Place the waste container in the hood. ***Chemicals will need to be replenished after every lab*** ***Waste will need to be disposed of after every lab***

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Instruction Notes: The iodine clock experiment has been used off and on at Florida State University for the past 20 years. We go away from it every semester or so to try something new, but we always seem to return to this experiment because of its simplicity and its coverage of both simple kinetics and the Arrhenius Equation.

The experiment itself can be completed in about 1–1 ½ hours. There are no special techniques needed to complete the experiment so instructional time at the front end of lab is minimal. The time that remains available at the end of the lab should be used to help students understand how to process their data. I have always advised my teaching assistants to spend the first few minutes of lab on technique and wait until the end of lab to deal with concepts. Students generally find it easier to understand how to analyze their data once they have collected it.

Points to Stress: The final run of the experiment changes the concentration of the limiting reagent, S2O32-.

Students often assume that this will change the rate of the reaction. It should be pointed out that for this experiment, only variations in temperature change the rate.

Potential Problems: The reaction procedure is very simple and thus we haven’t seen very many problems occur.

Difficulties can arise, however, if the reagents are not fresh enough. Do not make substantial amounts of reagent in advance of the class. Small amounts as described above should be produced instead.

Results can be damaged by contamination of the reagents as well. Make sure students do not pour excess reagent back into the lab supply. I say this repeatedly but I am always catching students who accidentally take too much reagent trying to return it to the reagent bottle. On occasion we have had an entire class’s results become “skewed” and deduced that this type of contamination was the culprit.

Finally, the students need to be careful to use different graduated cylinders to measure each reagent. If they cross-contaminate their cylinders, the reaction may start early and completely destroy the timing of the reaction.

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Pre-Lab Key

Experiment 13 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) For KI, Na2S2O3, and (NH4)2S2O8, calculate the initial concentration (M) based on the table of solution concentrations and volumes given in the procedure section. Copy the completed table into your lab notebook for later use. The shaded columns will be completed in lab.

Run Number

[KI] [Na2S2O3] [(NH4)2S2O8] Time Rate of Temp.

K (M) (M) (M) (sec) Reaction (K) (determine (M/sec) units)

1 2 3 4 5 6 7 8

0.075 0.050 0.025 0.075 0.075 0.075 0.075 0.075

0.00044 0.00044 0.00044 0.00044 0.00044 0.00044 0.00044 0.00022

0.019 0.019 0.019 0.012 0.0062 0.019 0.019 0.019

2) A first order chemical reaction has a rate constant of 9.1 × 101 s-1 at 12 °C and a rate constant of 4.1 × 102 s-1 at 36 °C. Calculate the activation energy, Ea, for the reaction.

4.6E-01 kJ/mol

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3) In the iodine clock reaction, you are measuring the rate of the reaction of iodide ion (I−) with peroxydisulfate ion (S2O82−): S2O82− + 2I−  2SO42− + I2

(Reaction 1)

The rate of the reaction is measured by measuring the amount of I2 formed in a particular time, or Δ[I2]/Δt. The I2 is measured by “titrating” it with a fixed amount of thiosulfate ion (S2O32−) which rapidly converts the I2 formed in Reaction 1 back to iodide ion: I2 + 2S2O32−  2I− + S4O62−

(Reaction 2)

(While Reaction 2 is said to be "fast" with respect to Reaction 1, it actually can proceed no faster than the rate at which I2 is produced, so its actual rate is controlled by the rate of Reaction 1.)

When the S2O32− is completely used up, then I2 begins to accumulate. It combines with starch to form a dark blue complex, so the “end-point” of the “titration” occurs when the solution turns blue. The time it takes for the blue color to appear is called the period of the clock, or Δt, the time it takes for Reaction 2 to go to completion. The rate of S2O32− disappearance is therefore given by its initial concentration divided by the clock period, or [S2O32−]/Δt. Because the rate of disappearance of I2 in Reaction 2 is controlled by its rate of appearance in Reaction 1, the rate of both reactions can be determined from the value of ½ × [S2O32−]/Δt. (Two S2O32− ions are consumed for every I2 consumed.)

(Actually, the I2 in solution combines with iodide ion to form triiodide ion, I3−, which is the species that forms the blue color with starch. That step is not shown to keep the equations simple, but adding it would not change the conclusions.)

The following graphs illustrate what is happening to the concentration of each of the species involved in these two reactions.

Show you understand the reaction by identifying the species associated with each of the lines in the graphs.

Initial conditions: [S2O82−] = [I−] = 0.055 M; [S2O32−] = 0.0015 M

[SO42−] =

____F_______

[I−] =

____E_______

[I2] =

____D_______

[S2O32−] = ____B_______ [S4O62−] = ____A_______ [S2O82−] = ____C_______

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Post-Lab Key

Experiment 13 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Which of the eight values of k should be the same (assuming no experimental error)? 1-5 and 8 2) Explain what factors affect the rate law constant.

Temperature, concentrations of reactants, catalysts, surface area of a solid reactant, pressure of gaseous reactants, or products all effect the rate of a reaction. The rate law constant is only constant at a particular set of conditions.

3) Which reaction proceeds faster, one with a large Ea or a small Ea? Explain why.

A small Ea means less energy is required for the reaction to start so the process can move faster.

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Experiment 14: Photometric Determination of an Equilibrium Constant Instructor Notes and Lab Preparation Chemicals and Equipment: 0.200 M Fe(NO3)3 acidified (0.5 M nitric acid) 0.002 M Fe(NO3)3 acidified (0.5 M nitric acid) 0.002 M KSCN 0.5 M nitric acid 10 mL volumetric pipettes Bulbs for transfer pipettes Cuvettes Spectrophotometers Kimwipes Setup:   

Place one Spectrophotometer on each lab bench Place a waste container in the hood Place remaining chemicals and supplies on the front counter

*** Make sure to replenish Kimwipes often*** Preparation: 0.200 M Fe(NO3)3 – To prepare 7 liters in amber carboy: Dissolve 565.60 g of Fe(NO3)3.H2O in 0.5 M HNO3 and dilute to the mark with 0.5 M nitric acid. Fill 1 liter amber bottles with the reagent.

0.002 M Fe(NO3)3 – To prepare 7 liters in amber carboy: Place 70 mL of 0.200 M Fe(NO3)3 and dilute to the mark with 0.5 M nitric acid. Fill 1 liter amber bottles with the reagent. *** LIGHT SENSITIVE*** 0.002 M KSCN – To prepare 7 liters, dissolve 1.36 g of KSCN in DI water and dilute to the mark.

Instruction Notes: This experiment is extremely simple to execute and takes approximately 1 ½ hours to complete.

The time left at the end of the lab can be put to good use going over the calculations required to complete the lab report. The experiment again uses a spectrophotometer and requires the production of a graph to determine the extinction coefficient. This process has been done several

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times in previous experiments, so students should require minimal instruction to complete this portion of the report. The calculation of the equilibrium constant, however, is best completed using a spreadsheet program such as Excel. The calculations are not difficult but their repetitive nature means they can be tedious.

Points to stress: Students will be using Mohr pipettes to make accurate concentrations of solution. This is the only technique being introduced in the experiment. Whether using bulbs or pumps, the proper technique should be demonstrated. The correct determination of the equilibrium constant depends on the accuracy of the solution concentrations.

Potential Problems: For obvious reasons, contamination of the solutions is one of the most serious sources of error.

Luckily, the solution changes color if contamination occurs and can therefore be corrected.

Another problem we have had is students not rinsing the cuvette completely between measurements. This can lead to skewed results.

The solutions themselves are light/heat sensitive. If the iron(III) nitrate solution loses its color, the results are no longer reliable. The solution should be kept in an amber bottle and made fresh as often as possible or practical.

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Pre-Lab Key

Experiment 14 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) You would like to determine the equilibrium constant for the binding of a ligand to a metal ion: M3+ + L− <=> ML2+ where K = [([ML2+])/([M3+][L-])]. The ML2+ complex has a unique absorbance in the visible region, so its concentration at equilibrium can be determined by measuring the absorbance of the solution at an appropriate wavelength. The concentrations of M3+ and L− can be determined from their initial concentrations and the concentration change in going to equilibrium using an ICE table. The following experiment is carried out first in order to determine the extinction coefficient of the complex at the wavelength chosen.

Calculate the complex concentration in each solution, and then its extinction coefficient.

0.200 M Solution Metal Ion (mL)

0.1 M HNO3 (mL)

0.0011 M Ligand Ion (mL)

Solution Concentration (M)

Absorbance

Extinction Coefficient (M-1cm-1)

A

5.0

4.0

1.0

0.00011

0.113

1027

B

5.0

3.0

2.0

0.00022

0.227

1032

C

5.0

2.0

3.0

0.00033

0.337

1021

D

5.0

1.0

4.0

0.00044

0.431

980

E

5.0

0.0

5.0

0.00055

0.562

1022

Plot the absorbance versus the solution concentration and determine the slope and intercept of the resulting line.

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2) Now you will determine the equilibrium constant, K, at several starting concentrations of M3+ and L−. The following table summarizes your measurements.

0.0023 M Solution Metal Ion (mL)

[M3+]init

0.0011 M Ligand Ion (mL)

[L-]init Absorbance

F

3.0

0.00069

7.0

0.00077

0.286

G

4.0

0.00092

6.0

0.00066

0.312

H

5.0

0.00115

5.0

0.00055

0.308

I

6.0

0.00138

4.0

0.00044

0.276

J

7.0

0.00161

3.0

0.00033

0.221

Using an ICE table, calculate the equilibrium concentrations of ML2+, M3+ , L− and K for each of the five solutions.

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Solution

[ML2+]=x

[M3+]equil

[L-]equil

F

0.000285481

0.00040

0.000485

1456.556

G

0.000311434 0.0006086

0.000349

1468.157

H

0.000307441 0.0008426

0.000243

1504.334

I

0.000275499 0.0011045

0.000165

1516.302

J

0.000220599 0.0013894

0.000109

1451.29

What value of K do you get for Solution H? 1504.334 What average value of K do you get? 1479.328

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K

Post-Lab Key

Experiment 14 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) How constant were your Kc values at room temperature? Explain any variation.

The answer to this question will depend on the student’s data. Their explanation should include references to any errors committed in measurements or calculation.

2) See if you can find literature and/or Internet references for the equilibrium constant for this equilibrium. Please cite the reference(s). How does your result compare with the one(s) you cited? Perhaps you will even find a similar experiment, with sample results, posted out on the Web.

The literature value most often found is around 138.

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Experiment 15: Titration of 7-Up© Instructor Notes and Lab Preparation Chemicals and Equipment: pH probes and meters Bottles of pH 4 buffer Bottles of pH 7 buffer Ring stands Clamps 50 mL burets 1 L bottle of 0.01 M NaOH 1 L bottle of 0.01 M HCL 2 L bottle of 7-Up 1 L bottle of 0.01 M phosphoric acid Dropping bottles of phenolphthalein Preparation: 0.05 g of phenolphthalein in 50 mL of ethanol and 50 mL of DI water.

0.01 M HCL and 0.01 M ***NaOH*** are purchased pre-made. The solutions are stored in sealed plastic carboys.

7-Up – pour into a 2 L beaker and boil to remove carbonation. Allow to cool and return to container.

0.01 M ***Phosphoric Acid*** – to make 7 liters in the large carboy: fill carboy with some DI water, add 4.73 mL of concentrated phosphoric acid, dilute to the mark. ***Do NOT label with Molarity***

Instruction Notes: This is a simple titration experiment which gives students experience titrating weak and strong, mono- and poly-protic acids, and analyzing their plots. We used to spend weeks drying and titrating KHP, but the students seemed to lose focus regarding why we were completing the experiment around week two. By combining the standardization of NaOH along with the titration of different acids as unknowns, there is an obvious and practical reason for completing each and every part of the experiment. The experiment requires approximately 2 ½ hours to

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complete all four titrations. If students add NaOH in no more than 1 mL increments at a time, the graphs are very nice. Each titration will require no more than ~ 45 mL to complete, so if the students refill the buret before beginning each titration, they will avoid the need to refill in the middle of a titration.

Points to stress: During the first titration, the students are asked to move through the process fairly quickly.

Before starting the second titration, they should develop a strategy as to where they will “slow down” to produce the most accurate plot of the equivalence point.

Once the concentration of the NaOH is determined during the first two titrations, the students should be able to determine the maximum pH that can be reached (about 12). This will prevent students from futilely continuing a titration trying to reach a higher pH. Potential Errors: We have not had any significant problems with this experiment.

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Pre-Lab Key

Experiment 15 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1)

To standardize your NaOH solution, you titrate 25.0 mL of 0.0822 M HCl solution, measuring the pH after each small addition of NaOH. Following is the graph you obtain.

The equivalence point occurs where there is a steep rise in pH.

What molar quantity of HCl is being titrated?

0.002 mol What molar quantity of NaOH has been added at the equivalence point? 0.002

mol What volume of NaOH solution has been added at the equivalence point? 23.9 mL Expanded view of the equivalence point: What is the concentration of the NaOH solution?

0.0084M

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2) A triprotic acid requires three moles of base to neutralize it, and the protons are removed one at a time as follows: H3A + OH− H2A− + H2O H2A− + OH−  HA2− + H2O HA2− + OH−  A3− + H2O A titration curve, which plots the pH of the acid solution versus quantity of base that is added, can reveal several things about the nature of the acid. First of all, some or all of the equivalence points can be observed in the graph as places where there is a maximum in the slope of the curve.

As an example, you have a 40.0 mL solution of a triprotic acid, H3A, with a concentration of 0.0588 M. You titrate it with a 0.350 M solution of NaOH. How many millimoles of H3A are you titrating? (Give units.) 2.35 mmol How many millimoles of NaOH are required to complete the titration? (Give units.)

7.056 mmol What volume of NaOH will be needed to completely titrate the acid? 20.16 mL What volume of NaOH will be needed to reach the first equivalence point? 6.72 mL What volume of NaOH will be needed to reach the second equivalence point? 13.44 mL 3) Citric acid is a weak organic acid found in citrus fruits and some soft drinks. It is an important intermediate in the citric acid cycle and therefore occurs in the metabolism of almost all living things. It is a triprotic acid, requiring three moles of OH− to remove all the acidic protons.

Titration of citric acid, as with all triprotic acids, occurs in a stepwise fashion in which one proton is removed at a time. The stepwise dissociation can therefore be represented by three equations, each with an acid dissociation equilibrium constant.

H3C6H5O7  H2C6H5O6− + H+

(Ka1)

H2C6H5O7−  HC6H5O62− + H+

(Ka2)

HC6H5O72−  C6H5O63− + H+

(Ka3)

The following graph shows how the concentration of each species changes when NaOH is added to a solution containing 100 mmoles of citric acid.

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What is the molecular weight of citric acid? 192.124 g/mol What is the mass of citric acid used in this titration? 19.2124 g After addition of how much base does [H2C6H5O7−] = [C6H5O73−]? 200 mmol After addition of how much base does [H2C6H5O7−] = [HC6H5O72−]? 150 mmol After addition of how much base does [H3C6H5O7] = [HC6H5O72−]? 100 mmol

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Post-Lab Key

Experiment 15 Post-Laboratory Question Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Are you able to determine all of the Ka values for citric acid? Why or why not?

Because the first two Ka values of citric acid are so closely spaced they are seldom seen as separate equivalence points. Normally only the first and the third are resolved.

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Experiment 16: Hydrogen Phosphate Buffer Systems Instructor Notes and Lab Preparation Chemicals and Equipment: pH probes and meters Bottles of buffer pH 4 Bottles of buffer pH 7 1 L bottle of 0.5 M NaOH 1 L bottle of 0.5 M HCl 1 bottle of KH2PO4 1 bottle of NaH2PO4*H2O 1 bottle of Na2HPO4*7H2O 1 bottle of Na2HPO4*12H2O Small weighing boats

Preparation: No preparation needed. 0.5 M standard solutions of NaOH and HCl can be purchased.

Instruction Notes: The buffer lab is designed to do three things: 1) give the students more experience preparing a solution, 2) give them more experience calibrating and using a pH meter, and 3) give them lots and lots of practice completing buffer calculations. The experiment itself is very simple to run and requires minimal setup for the instructor. Once the students prepare their buffer, they then test the buffering capacity of the solution by adding small increments of strong acid or base, measuring the pH after each addition. In the lab report, these measurements are compared to theoretical calculations using the same amount of acid or base. The main point of this lab is to give the students plenty of practice with ICE tables and the Henderson-Hasselbalch equation.

Points to Stress: The production of the buffer solution requires careful measurement of solid salts on the balance.

Stressing proper weighing technique here is important. Also, the proper use of stock chemicals should be reviewed before starting the experiment; otherwise contamination is probably going to occur. (Although, no matter how many times I say not to, someone always puts their excess back into the reagent bottle. Arghhh!)

Remind the students to make sure that after they create their buffer it is at room temperature before starting the test procedure. They should also be reminded to take an initial pH of the

.

buffer before starting each test run. This is important so they can get a baseline acid concentration and can thereby determine how well they made their buffer.

In the background it is mentioned that different pKa values should be used rather than the ones listed in most textbooks. This is because the activity coefficient of the buffer must be taken into account to get an accurate pH. We do not cover activity in our freshman classes so I did not go into detail about activity as part of the lab. But if you cover activity in your course it would be appropriate to stress this point as a practical application of activity coefficients.

Last, but not least, the calculations for the lab report require the student to calculate a theoretical pH for each addition of acid or base to compare to the experimental values determined by the pH meter; a percent error is calculated for each as well. Each acid and base addition normally requires 4–6 mL to change by 0.5 pH units for the full strength buffer, and 2–4 mL to change the diluted buffer. This is a lot of calculations and at least one should be demonstrated in lab so the students get a clear understanding of what is expected in their report. It is also important that the students calculate the initial HA and A- concentrations based on the actual (experimental) pH of the buffer. If they use the pH from their pre-lab, and the buffer pH is “off,” then every experimental pH will have a high percent error. Since the experiment, in part, is supposed to convince the students of the validity of Henderson-Hasselbalch, huge percent errors will be counterproductive.

Potential Problems: We have only run into a couple of problems when running this lab: 1) Trying to get the pH of the buffer to come out correctly, our pH values are always too low. At first we assumed student error, but later found that it was in fact our distilled water. Unfortunately, because of the climate in Florida our distilled water becomes fairly acidic (pH ~5.9) very quickly and thus our buffer pH values are always lower than expected. This is why I suggest that students back calculate their starting concentrations of acid and base from the experimental pH rather than just starting from the concentrations in their pre-lab. If they don’t do this calculation, then they should definitely consider the pH of the water they are using as a potential source of error in their comparisons of theoretical and experimental pH. 2) If the students have calculated their mass values incorrectly in their pre-lab this can cause error throughout the lab and report. I generally have my TAs run through the proper way to complete the calculations for at least one of the buffers. You can always select just one pH for an entire class to work with and go through the calculations for this pH. Rotate through different pH values from class to class to cut down on cheating/copying from class to class.

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Pre-Lab Key

Experiment 16 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) You need to prepare 1.000 L (in a volumetric flask) of 0.50 M phosphate buffer, pH 6.21. Use the Henderson-Hasselbalch equation with a value of 6.64 for pK2 to calculate the quantities of K2HPO4 and KH2PO4 you need to add to the flask.

Record the steps of these calculations and use them as a guide for the calculation you will have to make in the laboratory. (Calculate the intermediate values to at least one more significant figure than required.)

What is the [K2HPO4]/[KH2PO4] ratio you will need?

0.372 What concentrations of [K2HPO4] and [KH2PO4] will you need to make the total concentration 0.50 M? (Units required) [K2HPO4] = 0.135 M [KH2PO4] = 0.365 M How many moles of K2HPO4 and KH2PO4 will you need? (Units required) K2HPO4 = 0.135 mol KH2PO4 = 0.365 mol What mass of K2HPO4 and KH2PO4 will you need? (Units required) (FW of K2HPO4 = 174.2 g/mol. FW of KH2PO4 = 136.1 g/mol.) .

K2HPO4 = 23.6 g KH2PO4 = 49.6 g 2) Buffer capacity refers to the amount of acid or base a buffer can “absorb” without a significant pH change. It is governed by the concentrations of the conjugate acid and base forms of the buffer. A 0.5 M buffer will require five times as much acid or base as a 0.1 M buffer for a given pH change. In this problem you begin with a buffer of known pH and concentration, and calculate the new pH after a particular quantity of acid or base is added. In the laboratory you will carry out some stepwise additions of acid or base and measure the resulting pH values.

Starting with 60 mL of 0.50 M phosphate buffer, pH = 6.83, you add 1.7 mL of 1.00 M HCl. Using the Henderson-Hasselbalch equation with a pK2 for phosphate of 6.64, calculate the following values to complete the ICE table: What is the composition of the buffer to begin with, both in terms of the concentration and the molar quantity of the two major phosphate species?

(Units required) Concentration of HPO42−: 0.304 M Concentration of H2PO4−: 0.196 M Molar quantity of HPO42−: 18.2 mmol Molar quantity of H2PO4−: 11.8 mmol What is the molar quantity of H3O+ added as HCl, and the final molar quantity of HPO42− and H2PO4− at equilibrium?

Molar quantity of H3O+ added: 1.70 mmol Molar quantity of HPO42−: 16.5 mmol Molar quantity of H2PO4−: 13.5 mmol What is the new HPO42−/H2PO4− ratio, and the new pH of the solution? (Note: You can use the molar ratio rather than the concentration ratio because both species are in the same volume.)

HPO42−/H2PO4− ratio: 1.22 log(HPO42−/H2PO4−): 0.0872 pH: 6.73 Now take another 60 mL of the 0.50 M pH 6.83 buffer and add 3.7 mL of 1.00 M NaOH.

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Using steps similar to those above, calculate the new pH of the solution.

pH = 7.07

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Post-Lab Key

Experiment 16 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

Answer the following question: 1) Human blood contains a buffering system. What are the key molecules in this buffer and why is it necessary?

By far the most important buffer for maintaining acid-base balance in the blood is the carbonic-acid-bicarbonate buffer. The simultaneous equilibrium reactions of interest are: The pH of the buffered solution (i.e., the blood) is dependent only on the ratio of the amount of CO2 present in the blood to the amount of HCO3(bicarbonate ion) present in the blood (at a given temperature, so that pK remains constant). This ratio remains relatively constant because the concentrations of both buffer components (HCO3- and CO2) are very large, compared to the amount of H+ added to the blood during normal activities and moderate exercise. When H+ is added to the blood as a result of metabolic processes, the amount of HCO3- (relative to the amount of CO2) decreases; however, the amount of the change is tiny compared to the amount of HCO3present in the blood. This optimal buffering occurs when the pH is within approximately 1 pH unit from the pK value for the buffering system, i.e., when the pH is between 5.1 and 7.1.

All cells in the body continually exchange chemicals (e.g., nutrients, waste products, and ions) with the external fluid surrounding them. This external .

fluid, in turn, exchanges chemicals with the blood being pumped throughout the body. A dominant mode of exchange between these fluids (cellular fluid, external fluid, and blood) is diffusion through membrane channels, due to a concentration gradient associated with the contents of the fluids. Hence, the chemical composition of the blood (and therefore of the external fluid) is extremely important for the cell. If, for instance, the pH of the blood and external fluid is too low (too many H+ ions), then an excess of H+ ions will enter the cell. Maintaining the proper pH is critical for the chemical reactions that occur in the body. In order to maintain the proper chemical composition inside the cells, the chemical composition of the fluids outside the cells must be kept relatively constant. http://www.chemistry.wustl.edu/~edudev/LabTutorials/Buffer/Buffer.html 2) Using stoichiometry, explain why the capacity of a diluted buffer solution is lower than a more concentrated buffer.

Buffers act by reacting mole to mole (stoichiometrically) with added acid or base. At a diluted concentration the number of moles of buffer is reduced and therefore, the number of acid or base molecules capable of reacting with added acid or base is reduced as well. The buffer molecules are therefore, more quickly used up and the system is not as pH stable as a system with a buffer of higher concentration.

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Experiment 17: Entropy, Free Energy, and Chemical Equilibrium Instructor Notes and Lab Preparation Chemicals and Equipment: 50 mL burets Ring stands Buret clamps 1 L bottle of 0.05 M HCl 1 bottle of calcium hydroxide 1 rack of thermometers 2 dropping bottles of bromothymol blue Hotplates or Bunsen burners 1 pack of 0.1 nm filters (Whatman # 5 or 6, or equivalent) Corks Transfer pipettes Bulbs for transfer pipettes Buchner funnels Filter flasks 1 box of rock salt 1 cooler of ice Ice trays Waste container Setup: Place 1 hotplate on each bench.

The rest of the chemicals are placed on the front counter.

Place a waste container in the hood.

Preparation: To make 250 mL of Bromothymol blue: To a 250 mL volumetric flask, add 0.1 g of bromothymol blue powder in 16 mL of 0.01 M NaOH. Dilute to the mark w/ DI water.

Instruction Notes: The maximum in solubility for Ca(OH)2 is around 4 °C. When we first developed this lab we attempted to get students to find the temperature of maximum solubility, but this really required more data points at low temperature than the students were capable of collecting. Now the lab simply requires them to determine the trend in solubility and calculate the thermodynamic properties related to the solubility. .

Points to Stress: Students will be using the filtration flasks again. Because the determination of concentration is very dependent on good technique in this lab, it is extremely important that they filter their solutions carefully and that they don’t attempt to titrate cloudy solutions as this will throw off the results significantly.

Within the calculations they complete for their lab report, the dilution of the saturated solution before titration should be discussed. Because we are approximate in the volume of water added in the dilution, the students get confused as to how to determine the concentration. It should be stressed that we are titrating to determine the number of moles of hydroxide ion that were in the original 10 mL sample.

Potential Problems: As mentioned above, technique is very important. The most significant problem with this lab has come from students titrating cloudy solutions. Remind them that they only need a small amount of filtrate to titrate so they can filter very slowly and carefully without a significant loss of time.

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Pre-Lab Key

Experiment 17 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) You add an excess of Ca(OH)2 to water maintained at a particular temperature,

stir until the solution is saturated, filter, then determine the [OH−] in the solution by titration with acid. Titration of 10.0 mL of the calcium hydroxide solution to the endpoint requires 4.86 mL of 0.070 M HCl solution.

What is the molar quantity of OH− in the 10.0 mL of solution? 0.340 mmol

What are the concentrations of Ca2+ and OH−? 0.0170M and 0.0340 M, respectively.

What is the solubility of Ca(OH)2 under these conditions? 0.126 g/dL What is the Ksp for Ca(OH)2 under these conditions? 1.97E-05 2) You measure the solubility of a salt at four different temperatures and calculate the following Ksp values: Temperature (°C): Ksp:

ΔGosoln (kJ/mol)

15

31

51

70

2.89×10-2 4.11×10-2 6.08×10-2 8.45×10-2

8.5

8.1

7.5

7.0

ΔGosoln for the solubility process is given by the relationship ΔGosoln = -RTlnKsp where R = 8.314 J/(K*mol) and T is the absolute temperature in degrees K. Calculate ΔGosoln (in kJ/mol) at each of these temperatures. (In table)

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As the temperature increases, how does the solubility of the salt change?

Increases According to Le Châtelier's principle, dissolving this salt should therefore be which type of process? Endothermic In other words, what is the sign of ΔHosoln? Positive 3) It is possible to determine both ΔHosoln and ΔSosoln for the solution process from this data. One or both of two graphical methods can be used: Method 1: From the equation: Method 2: Alternatively, rearranging the o o o ΔG soln = ΔH soln-TΔS soln equation relating Ksp and ΔGosoln: a plot of ΔGosoln versus T will give ΔSosoln as the slope, and ΔHosoln as the intercept. and plotting lnKsp versus will give as the slope, and

as the

intercept.

Using one or both of these graphical methods, determine ΔHosoln (in kJ/mol) and ΔSosoln (in J/(K*mol)) for this salt. (Be sure to pay close attention to your units and signs in both your plots and your calculations.) o H soln = 16 kJ/mol o S soln = 26 J/(K*mol)

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Post-Lab Key

Experiment 17 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Why do you have to filter the Ca(OH)2 solution before titrating it?

As you are making a saturated solution, there will be non-dissolved particles of salt left over. Removal of these particles is essential to get an accurate concentration of the dissolved salt.

2) Why don't you have to accurately record the amount of Ca(OH)2(s) used in the saturated solution?

Since you remove all of the non-dissolved salt and titrate the saturated solution to determine its concentration, accurate measurement of the amount of solid salt placed in the solution is a waste of time.

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Experiment 18: Electrochemistry: The Nernst Equation Instructor Notes and Lab Preparation Chemicals and Equipment: Bottle of 0.1 M CuNO3 solution Bottle of 0.1 M Mg(NO3)2 solution Bottle of 0.1 M Fe(NO3)3 solution Bottle of 0.1 M Pb(NO3)2 solution Bottle of 0.1 M Al(NO3)3 solution Bottle of 0.1 M KNO3 1 bottle of each of the metal strips: copper, lead, aluminum, iron, magnesium 1 container of agar powder (inexpensive type, Fishersci #AC40039-5000 or equivalent) 1 container of cotton 1 sheet of sand paper, cut into about 12 x 15 cm pieces 1 cooler with ice U-tubes Thick-stemmed disposable pipettes Hotplates – 2 per lab bench Multi-meters with alligator clips Setup: Place all chemicals and supplies on the front counter.

Place a metals waste container in the hood.

Place 2 hotplates on each lab bench.

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Preparation: The salt solutions should be prepared as nitrates in a 0.1 M nitric acid solution.

Instruction Notes: Experiment 19 is all about electrochemistry. Each student builds his or her own unit cell and measures the potential. There are actually only five metals: Fe3+, Cu+, Mg2+, Al3+, Pb2+. You should assign each combination to at least two students (if possible) so that the data is reproduced when pooled together. I normally create a table on the front board so that as the students complete the measurements they can report their findings to the class. They only pool the room temperature data. We have recently dropped the concentration to 0.1 M to save on chemical cost. This appears to work just as well as the 1.0 M solutions. We use small pieces (.3 cm x 1 cm strips) of each metal cut from either thin sheet or ~14 gauge wire. You have to really watch the students so they don’t throw them away and that they are returned to the correct container at the end of the lab.

Points to Stress: The students should have calculated the theoretical potentials for all of the cells before coming into class. The experiment allows them to confirm their calculations and should also help them decide which metal should act as the cathode and which should act as the anode. In other words, if they have prepped their lab properly, the potentials should be positive and similar to the ones they have calculated.

The inclusion of the Nernst equation and temperature dependence of cell potential allows this experiment to bridge the gap between chapter 17 and 18, making another use of the thermodynamic properties enthalpy, entropy, and free energy. The importance and relevance of the relationship of these properties to yet another chemical process should be discussed.

Potential Problems: When setting up the u-tubes, make sure the students soak the cotton in the respective solutions before placing them in the tubes. If the cotton is not soaked through they will get bad readings.

We use multi-meters (set to VDC) for this experiment. Occasionally we have had problems with students not setting the meters to VDC. This error is the first thing to check when students report anomalous readings.

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Pre-Lab Key

Experiment 18 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Using literature values for the half-cell reactions of the following metals,

calculate the total cell potentials for each of the 10 possible cells. Assume complete reduction of the metals. Show all your work please. Copy the resulting potentials into your lab notebook.

Metals: Fe3+, Cu+, Mg2+, Al3+, Pb2+ 3Cu+ + Fe ↔ Fe3+ + 3Cu Ecell = 0.521 V + 0.04 V = 0.561 V 2Fe3+ + 3Mg ↔ 2Fe + 3Mg2+ Ecell = - 0.04 V + 2.37 V = 2.33 V Fe3+ + Al ↔ Fe + Al3+ Ecell = - 0.04 V + 1.66 V = 1.62 V 2Fe3+ + 3Pb ↔ 2Fe + 3Pb2+ Ecell = - 0.04 V + 0.126 V = 0.086 V 2Cu+ + Mg ↔ 2Cu + Mg2+ Ecell = 0.521 V + 2.37 V = 2.89 V 3Cu+ + Al ↔ 3Cu + Al3+ Ecell = 0.521 V + 1.66 V = 2.18 V

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2Cu+ + Pb ↔ 2Cu + Pb2+ Ecell = 0.521 V + 0.126 V = 0.647 V 3Mg + 2Al3+ ↔ 3Mg2+ + 2Al Ecell = -1.66 + 2.37 V = 0.71 V Mg + Pb2+ ↔ Mg2+ + Pb Ecell = -0.126 + 2.37 V = 2.24V 2Al + 3Pb2+ ↔ 2Al3+ + 3Pb Ecell = -0.126 + 1.66 V = 1.53 V 2) The standard potential of a galvanic cell (Eocell) can be determined from thermochemical data such as the free energy of formation (ΔGof) of each of the species in the reaction.

For example, given the reaction: Fe3+(aq) + Al(s)  Al3+(aq) + Fe(s)

Species ΔGof (kJ/mol)

Fe3+(aq) Fe(s) Al3+(aq) Al(s) -10.5

0

-481.2

First use the data to the right to calculate the ΔGorxn. -470.7 kJ/mol Then calculate Eocell from ΔGorxn. 1.626 V 3) Following is the shorthand notation for a galvanic cell: Al(s)|Al3+(1M)||Cu+(1M)|Cu(s) Write the overall balanced equation occurring in this cell. (Indicate the physical states of reactants and products, i.e., Al(s), Cu+(aq), etc.)

3Cu+(aq) + Al(s)  Al3+(aq) + 3Cu(s) Identify the anode and cathode of this cell.

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0

Cu(s) Cathode; Al(s) Anode You make the following voltage measurements at four different temperatures: Temperature (°C):

10

35

55

74

Eocell (V):

2.190

2.176

2.165

2.154

ΔGorxn (kJ/mol) -633.9 -629.9 -626.7 -623.5 Calculate the ΔGorxn at each of these temperatures. (In table) Now graph this data using the relationship ΔG = ΔH - TΔS, and determine ΔHorxn (in kJ/mol) and ΔSorxn (in J/(K*mol)) for this process. (This assumes that ΔHorxn and ΔSorxn do not vary with temperature. Remember to use absolute temperature, and pay attention to units and signs in your plot and your calculations.)

ΔHorxn = -679.8 kJ/mol and ΔSorxn = -162.2 J/(K*mol)

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Post-Lab Key

Experiment 18 Post-Laboratory Question Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) What is a Watt? How is it related to the potential of a battery?

A Watt is a unit of energy defined as a Joule per second. Battery power is defined as the amount of energy that can be delivered by a unit of time. The greater the potential of the battery the more energy that can be delivered in a unit of time, thus the Watt is a unit of energy that can be used to describe the potential (power) of a battery.

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Experiment 19: Analysis of a Hydrate Instructor Notes and Lab Preparation Chemicals and Equipment:          

Barium chloride dihydrate Copper(II) sulfate pentahydrate Magnesium sulfate heptahydrate Bunsen burners Ring stands Ring clamps Clay triangle Crucibles + lids Crucible tongs Analytical balances

Setup: Place hydrates in small plastic bottles and mark as unknowns A – C. Set out crucibles, ring stands, clay triangles, and Bunsen burners for each student.

Instruction Notes: This is a simple dehydration experiment. The hydrates given are sufficiently different in water percentage so the identification should be easy barring experimental error.

Potential Problems: From experience, be warned that the hydrate salts tend to get spilled a lot onto the balances.

Students should be monitored closely to make sure they do not leave a mess.

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Pre-Lab Key

Experiment 19 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

Use the information to answer the questions. Show work, include units, and put your answers in the blanks.

1) Geoffrey weighs an empty beaker and finds it to have a mass of 95.83 g. After putting a spoonful of an unknown hydrate into the beaker, he finds that the mass has increased slightly to 99.87 g. He heats the beaker and its contents twice, and finds that the mass has dropped to 97.22 g. Geoffrey is told by his teacher that the molar mass of the anhydrous salt is 74.10 g. a) What mass of hydrate did Geoffrey start with? _____4.04 g_____________________ b) How much water was driven off from the hydrate during the heating process in units of…

1. grams?

____2.65 g____________________ 2. moles?

___ 0.147 mol_________________ c) How much anhydrous salt remained in the beaker in units of: 1. grams?

____ 1.39 g__________________ 2. moles?

_____0.0188 mol______________

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d)

Write down the mole ratio as decimal numbers:

__0.0188____ moles anhydrous salt : __0.147____ moles water Write down the mole ratio as whole numbers:

__1____ moles anhydrous salt : __~8____ moles water e) What is the formula of the hydrate? (use MN to symbolize the anhydrous salt)

_______MN.8H2O_________________ f) Based on Geoffrey’s data, calculate the percentage of water in the sample of hydrate.

_______65.6%_________________ g) Why must Geoffrey heat the sample twice instead of just once?

To completely rid the sample of water.

2) In an experiment, a student dehydrated a hydrate unknown and found that the mass of hydrate was initially 4.13g. After heating, the mass of the desiccant was 3.52g. (a) Calculate the mass percent of water in the hydrate. 14.7% (b)

The table below lists the possible hydrates that could be your unknown.

Complete the table.

Hydrate

Mass Salt (g) Mass Water (g) Mass Percent (%)

CuSO45H2O

249.68

90.05

36

BaCl22H2O

208.236

36.02

17

MgSO47H2O

246.47

126.07

51

(c)

Compare the mass percent values to determine the identity of your unknown.

Mass percent of the unknown is closest to BaCl22H2O

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Post-Lab Key

Experiment 19 Post-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Give some reasons why this experiment for finding the formula of hydrates might not work?

Underheating or overheating, loss of salt in transfer, etc. will give bad mass information that can skew results.

2) If the Bunsen burner left some black soot on the bottom of your crucible, how would this change your answer?

This answer will depend on the answer to the first, but if the student did not already mention a weighing issue it should be discussed here.

3) What are some unavoidable sources of error (like procedural) to explain why you could not achieve 100% accuracy?

One of the most unavoidable sources of error that is not accounted for in this procedure is the humidity in the atmosphere. If the procedure is completed in the Southeast where humidity is high, the results will differ from those completed in the Midwest where humidity is much lower since the hydrates can reabsorb water from the atmosphere as they cool.

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Experiment 20: Qualitative Analysis: Cations, Anions, and Complex Ions Instructor Notes and Lab Preparation Chemicals and equipment:                   

Well plates Wax pencils Bunsen burners Hoses Wire pieces 2 flame tests (in the hood) Metals waste jar Kimwipes Multiple sets of the reagents: 6 M NaOH 6 M NH3 6 M HCL 6 M nitric acid Concentrated sulfuric acid 0.1 M nitrates: Na, Mg, Ni, Cr, Zn, Ag, Pb 0.1 M NaNO3, NaCI, NaCL, NaSO4 0.5 M NaI (acidified) 0.5 M FeSO4*7H2O Unknown containing 2 metals

Setup:    

Set up 2 flame tests in the hood Place waste jar in the hood Place 1 set of reagents on each lab bench Place all other supplies and chemicals on front lab bench

Preparation: Prepare 3 or 4 unknowns, each containing 2 different metals. Label as unknown 1–7. Place each chemical solution in a dropping bottle. Note: If students will be sharing chemicals, we have found it helpful to mark the paper covering the counter with the names of each chemical so that students return the bottles back to where they find them.

Set up a hot water bath in the hood. .

Set out small weighing boats.

Instruction Notes: This is a shortened, one day, all-inclusive qualscheme lab. We have cut way down on both waste and chemical costs by microscaling the tests used to confirm the unknown content.

Points to Stress: Good note taking and careful observations are very important in this lab. Make sure that students know to not only write about the color of the precipitates but also their texture, etc.

Make sure that the students understand that having a plan of attack when starting on the unknown is essential to complete the lab in a timely fashion, and to be confident about their results. This is supposed to be a qualscheme activity so they should proceed through a standard set of reactions to separate the various ions, and to avoid false positives.

Potential Problems: We have white well plates which can sometimes make seeing the detail of white precipitates difficult. We tell students that they can also run the tests in a test tube if they have trouble making good observations. Students also often forget about the anions in the mixture and, since they should be finding both, must be reminded of this aspect of the lab.

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Pre-Lab Key

Experiment 20 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) The qualitative analysis scheme depends on the relative solubilities of the salts of the cations with various anions. The aqueous solubility of a salt can be estimated from Ksp, the solubility product constant, describing the equilibrium reaction between the solid salt and the ions in solution: MxAy(s)  xMy+(aq) + yAx-(aq),

Ksp = [My+]x[Ax-]y

(MxAy does not appear in the expression because the pure solid is considered to be in its standard state with an activity of 1.)

The actual solubility may differ slightly from the calculated value, though, if the undissociated salt itself is slightly soluble, and if there are intermediate ion-pair structures formed in solution. Nevertheless, the calculated value is still very useful in comparing solubility relationships of various ions, and predicting whether or not a precipitate will form when two ionic solutions are mixed.

For the salt BaSO4, Ksp = 1.1×10-10.

What is [Ba2+] for a saturated solution of BaSO4? 1.0e-5 M What is the calculated solubility for BaSO4? 2.4e-4 g/dL Addition of a soluble sulfate, such as Na2SO4, will suppress the solubility of the Ba2+ ion because of the common ion effect.

What will [Ba2+] be in the presence of 0.15 M Na2SO4? 4.4e-10M

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2) For the salt SrF2, Ksp = 4.3×10-9.

What is [Sr2+] for a saturated solution of SrF2? 1.0e-3M What is the calculated solubility for SrF2? 1.3e-2 g/dL What will [Sr2+] be in the presence of 0.10 M NaF? 4.3e-7M 3) To estimate whether a salt will precipitate under a particular set of conditions, a reaction quotient, Qsp, can be calculated from given concentrations of the cation and anion from the equation: Qsp = [My+]x[Ax-]y If Qsp is less than or equal to Ksp, no precipitate will form, but if Qsp is greater than Ksp, a precipitate will form and the ion concentrations will decrease until they reach values that satisfy the Ksp expression.

In the case of metal hydroxides, the anion is the hydroxide ion, and its concentration depends on the pH of the solution. Whether a precipitate occurs will therefore depend both on the Ksp of the hydroxide and on the pH. For the reaction: Mg(OH)2  Mg2+ + 2OH−, Ksp = 5.6×10-12.

You prepare a solution of 0.20 M Mg2+ and adjust the pH to 6.0.

What is the value of the reaction quotient, Qsp, under these conditions? 2.0e-17 Will a precipitate form? No What will be the equilibrium concentration of the Mg2+ ion? 2.0e-1M

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4) Insoluble salts can sometimes be brought into solution by the formation of a soluble complex ion. For example AgCl(s) can be dissolved in ammonia solutions by the formation of the soluble Ag(NH3)2+ complex ion. The ammonia competes with the Cl− ion for reaction with Ag+ as follows: AgCl(s)  Ag+(aq) + Cl−(aq) Ag+(aq) + 2NH3(aq)  Ag(NH3)2+(aq)

Sum:

AgCl(s) + 2NH3(aq)  Ag(NH3)2+(aq) + Cl−(aq)

Assuming all the Cl− comes from the dissociation of AgCl, then [Ag(NH3)2+] = [Cl−], and or What is the maximum [Ag(NH3)2+] that can be obtained by dissolving AgCl in 1.5M ammonia? 8.0e-2M 5) For another silver salt, AgBrO3, Ksp = 5.4×10-5.

What is Knet for the reaction AgBrO3 + 2NH3  Ag(NH3)2+ + BrO3−? 8.6e2 What is the maximum [Ag(NH3)2+] that can be obtained by dissolving AgBrO3 in 4.1 M ammonia? 1.4e2M

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Post-Lab Key

Experiment 20 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Which cation forms a 2-ligand complex with NH3?

Silver 2) Which anion does not form a precipitate with lead?

Nitrate 3) Write the balanced equation for the dissociation of the compound in question 2 in water.

Pb(NO3)2(aq) Pb2+(aq) + 2NO3-(aq)

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Experiment 21: Isomerism in Coordination Chemistry Preparation of a Coordination Compound: Tetraamminecopper(II) Sulfate Monohydrate

Instructor Notes and Lab Preparation Chemicals and Equipment: 2 50 mL beakers 1 400 mL beaker 1 glass stir rod 1 25 mL filter funnel 1 125 mL side-arm flask for filtration 1 ring stand 1 Pasteur pipette and bulb 1 watch glass 1 vacuum trap 5 mL of 15 M ammonia solution 5 mL of deionized water ~0.400 g copper sulfate pentahydrate 15 mL of 95% ethanol 10 mL of 1:1 (by volume) ethanol/ammonia solution Ice

Setup: Place copper sulfate pentahydrate on the front counter along with weigh boats.

Place ammonia and ethanol solutions in 1 L containers in the hood.

Instruction Notes: Tetraamminecopper(II) sulfate is readily formed in high yield provided that the “mother liquor” is not too dilute with respect to the reactants. Students should be reminded not to use more solvent than described in the procedures. Although the drying time can be reduced with heating, the use of a drying oven is discouraged because of the volatility of ammonia.

Points to Stress: The procedure given is able to determine the stoichiometry of the compounds within 1–5% without rigorous analytical technique. Students should be reminded however that the more meticulous their technique, the better both their yield and their data will be.

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Potential Problems: The only issues occur if students add too much liquid during the synthesis of the crystals.

Students should be warned to be cautious not to overly “wet” the salt.

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Pre-Lab Key

Experiment 21 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

Use the information to answer the questions. Show work, include units, and put your answers in the blanks.

1) Write a balanced chemical equation for the reaction of CuSO4 · 5H2O with ammonia to form [Cu(NH3)4] SO4 · H2O.

CuSO45H2O + 4NH3  [Cu(NH3)4] SO4 · H2O 2) How many moles of ammonia are present in 5 mL of a 15 M NH3 solution?

0.075 mol 3) How many moles of Cu2+ are present in 1.5 grams of CuSO4 · 5H2O?

0.006 mol 4) Calculate the molar mass of [Cu(NH3)4] SO4 · H2O.

MW = 245.75 g/mol

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Post-Lab Key

Experiment 21 Post-Laboratory Question Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Based on the color of the filtrate, is it possible to obtain a 100% yield on this reaction?

Explain why/why not?

This answer will depend on the student’s results but in general they should respond that a 100% yield is not possible. Reasons may include inadequate mixing, purification, etc.

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Experiment 22: Light Emitting Diode Instructor Notes and Lab Preparation Chemicals and Equipment: Ruthenium(III) chloride trihydrate Sodium tetrafluoroborate Gallium indium eutectic Hypophosphorous acid, 50 wt % in water 2,2’-Dipyridyl Sodium hydroxide Acetonitrile Tin oxide coated glass. This comes with the ICE Nanocrystalline Solar Cell Kit; the kit has ten pieces and this experiment only requires one Voltmeter (with leads) 6-volt power supply, such as a lantern battery Cotton swab

Preparation: “RuCl3” from commercial RuCl3·3H2O Place the commercial RuCl3·3H2O (~3 g) in a mortar and pestle. Place this in an oven (120 to 130 °C) for 3 hours, allow it to cool, and grind it. Place the ground material back in the oven overnight prior to the laboratory experiment. There is most likely some water remaining, as anhydrous RuCl3 is insoluble. The product should be dark brown when dry.

NaH2PO2 preparation from H3PO2 Hypophosphorous acid is monoprotic (two of the hydrogens are attached directly to phosphorus).

It does not fume, but with a pKa of 1.1, it is stronger than phosphoric acid. Use caution when handling, and keep it away from oxidizers (do not store it in a cabinet where the usual mineral acids are kept!). Measure 10 mL of 50 % H3PO2 into a beaker with a stirbar, add 6 mL of water, and begin stirring. Slowly add pellets of NaOH until the pH is about 6–8 (wait until the previous pellet dissolves before adding the next). This will require about 3.9 g of NaOH. The final solution is about 6 M in NaH2PO2. If the solution becomes too basic, additional drops of H3PO2 can be added to bring the pH in the correct range.

Synthesis Dissolve 0.083 g (0.40 mmol) “RuCl3” in 8 mL of water in a 30-mL beaker equipped with a Teflon-coated stirbar. Add 0.188 g (1.204 mmol) 2,2’-dipyridyl and, using a syringe or Mohr pipette, 0.44 mL of NaH2PO2. The liquid will be deep greyish-green. Mark the beaker to indicate the level of liquid, place a watch glass on top of the beaker, and heat so it boils for 30 min. Add

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water as necessary to bring the level to the mark. The liquid will now be deep orange. Allow the solution to cool to room temperature. Dissolve 0.333 g of NaBF4 in 1.5 mL of water and add the resulting solution to the orange solution. Solid should form. Heat the mixture until the solid dissolves, adding water as needed. Cool in ice; crystals should form. Collect these by filtration and wash with ethanol. Draw air through the orange solid to dry.

Instruction Notes: This experiment is fairly sophisticated and uses chemicals more suited to an organic laboratory.

However, if Parts I–III are completed for the student, the production of the light emitting diode is fast and simple enough to complete in the general chemistry lab.

Points to Stress: The relevance of this experiment to what is currently being researched is important. Thin organic materials such as the diode created here have only been produced in a meaningful way within the last 10 years or so and the production of better “diodes” is an area of very active research. The ability to make a light from a thin, easily produced chemical compound has use for flat-screen displays from cars to watches, to safety striping on streets. Understanding the chemical physics behind the functionality of these compounds is therefore important and relevant to any chemist.

Potential Problems: None noted to date.

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Pre-Lab KEY

Experiment 22 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

Use the information to answer the questions. Show work, include units, and put your answers in the blanks.

1) Calculate the theoretical yield of [Ru(bpy)3](BF4)2 (MM = 743.29 g) recovered. Use the RuCl3 as the limiting reactant.

2) Identify the central atom, ligand(s), and coordination number in the coordination compound [Ru(bpy)3](BF4) 2. Draw a sketch of the complex ion.

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3) Would the photoluminescence in this reaction be classified as fluorescence or phosphorescence? Why?

It is phosphorescence because it took several seconds for the light to appear whereas if it were fluorescence it would happen almost instantly.

4) When the negative and positive leads were switched, no light is emitted. Why is this and what does this prove?

This shows that the circuit only allows current to flow one way, and proves that the device is a diode.

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Post-Lab Key

Experiment 22 Post-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) Calculate the percent yield of [Ru(bpy)3](BF4)2 (MM = 743.29 g) recovered. Use the RuCl3 as the limiting reactant.

The answer to this question will depend on the student’s data and should use the theoretical yield calculated in their pre-lab problem 1.

2) Does the [Ru(bpy)3]2+ complex exhibit a small or large crystal-field splitting energy? Explain. (Hint: Think about the color of the pure compound.)

It has a large crystal- field splitting energy. The color of a compound is the wavelength of light that an object reflects. Since it is orange, the compound reflects a relatively longer wavelength of light. The longer the wavelength reflected, the shorter the wavelength absorbed, or the larger the crystal-field splitting energy.

Copyright © 2014 Pearson Eduaction, Inc.

Copyright © 2014 Pearson Eduaction, Inc.

Experiment 23: Amino Acid Analysis by Ninhydrin Instructor Notes and Lab Preparation Chemicals and Equipment:        

Pencils Rulers TLC Plates (1/student). Avicel microcrystalline cellulose 20x20cm 250microns. Cat# 05011 (Analtech uniplate, analtech.com) TLC Reservoirs: placed in hood.

Hot Plates/Heat Gun/oven (set to 110oC) Glass capillaries (Kimax-51 EF3773B, size 0.8-1.1x100mm) Bunson burners and strikers Waste container for used solvent (maybe recycled)

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3 Different Amino Acid Standards (Glutamic Acid, Lysine, & Serine) o To obtain 12 mg/mL stock solutions, weigh out 1.2 g of each amino acid and dissolve it in 100 mL of ddH2O. Each of the three standards goes in its own well-labeled reagent bottle with the concentration and name of amino acid. Note: For Glutamic acid (which has a solubility of 0.864g/100ml (8mg/ml)), make an 8mg/ml solution.

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5 Different Unknowns (5 combinations with each containing two of the standard amino acids and supplied in a reagent bottle labeled as “Unknown A”, “Unknown B”, etc.). o Production of these “unknowns” is nearly identical to the procedure for the standard solutions given previously. The only difference is that in these solutions there will be 1.2 g of EACH amino acid that will be in the unknowns.  Unknown A = 0.8 g of glutamic acid dissolved in 100-mL ddH2O.  Unknown B = 1.2 g of lysine dissolved in 100-mL ddH2O.  Unknown C = 1.2 g of serine dissolved in 100-mL ddH2O.

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Chromatography Solvent (1-Propanol:di water 1:1 volume ratio) o Prepare enough for 200 mls/student (~4L/class), place it in 1 L bottles, and keep in the hood.

(Can be recycled.)

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0.2% Ninhydrin Solution (w/v): 0.2 g in 100 mls ethanol. The ninhydrin is sprayed onto the TLC plate using a Spra-tool spray gun. Pour the ninhydrin solution into the the polyjar and attach to the spraygun. Place in hood along with cardboard box spray shield and tray.

Setup:  

Standard solutions, unknown solutions, pencils, rulers, TLC plates, and heat gun should all be placed on the front counter.

The TLC reservoirs should be placed in the hood, along with a waste container for used solvent so that it may be recycled.

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4 functional hot plates should be placed at each bench.

Instruction Notes: Each student will have one of three unknowns to identify: (1) Unknown A = Glutamic Acid; (2) Unknown B = Lysine; and (3) Unknown C = Serine. The two essential values which must be calculated and discussed are the unknown’s molecular weights and isoelectric point. The molecular weight can be deduced from the equivalence volume obtained, and the weight of unknown titrated, while the isoelectric point can be determined from the student’s generated curve, which must be turned in showing the pertinent data as outlined in the graphs on the following pages.

Unknown A (Glutamic Acid)

Unknown C (Serine) Unknown B (Lysine)

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Pre-Lab Key

Experiment 23 Pre-Laboratory Assignment Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

1) How would you expect the amino acids glycine, alanine, and phenylalanine to travel relative to each other on the silica plates using a mixture of chloroform/methanol/ammonia? Explain your answer.

As glycine, alanine, and phenylalanine all possess side chains that are essentially nonpolar (glycine is an exception in some texts), these amino acids will be separated only by their molecular weight as the mobile phase in this experiment is extremely polar.

Essentially, the amino acids will “spot” from greatest molecular weight to smallest (the heaviest amino acid will be towards the bottom of the plate, while the lightest will be towards the top). Thus, there will be three spots moving up the silica plate, the first spot will be phenylalanine, the second alanine, and the third glycine.

2) You spot an unknown sample on a TLC plate and elute it with chloroform.

After visualizing the plate, you detect only one spot with an Rf of 0.97.

Does this indicate that the unknown material is a pure compound? What could be done to verify the purity?

Yes and no! The unknown material could be a pure compound if it truly “spotted” immediately before the solvent front. However, the unknown material may also not be a pure compound on the basis that components making up the unknown were nearly identical to those used in the solvent, and were unable to efficiently separate. To truly validate the purity of the compound a different solvent could be utilized to run the separation again, or a sensitive analytical technique such as GC-MS could be utilized.

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3) You spot a sample of an unknown pure compound on a TLC plate and elute it with hexane. It’s Rf is 0.25. Would you expect its Rf to increase, decrease, or remain the same if you elute with acetone? Why? Hexane is non-polar whereas acetone is polar and since the unknown had a small Rf, its affinity for the non-polar solvent was relatively low and did not want to move up the plate with the solvent. On the other hand, if acetone was utilized as the solvent of choice, the Rf of the unknown would most likely increase as its affinity for the solvent would now be greater than what it was for the non-polar hexane.

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Post-Lab Key

Experiment 23 Post-Laboratory Questions Name: ___________________________ Date: _________________ Instructor: ____________________________ Sec. #: ___________

Show all work for full credit.

Answer the following questions: 1) Why is it important that the spots you made on the TLC plate not be more than 1 mm in diameter? What might have happened if the spots had been larger?

Large spots cause bleedover between lanes.

2) We saw in the mechanism for the ninhydrin reaction that the molecule reacts with amino acids via their amine group. What other molecules might give a “false positive” reaction with ninhydrin?

Cleaning products like bleach or ammonia can react with ninhydrin. Any other molecule that contains an amine group is possible.

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