Chem 111 Lab Manual-2016 PDF

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Chem 111 Lab Manual-2016...

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Supplemental Laboratory Manual

CHEM 111 Physical Chemistry Laboratory Winter 2016

Department of Chemistry University of California, Riverside

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Chemistry 111

Report Format

Chemistry 111 Report Format

All reports should be typed, have a cover sheet, and be stapled. (For general ideas see Garland, Nibler, and Shoemaker (GNS), 8th Ed. pp. 10-28) General Format (Long Reports) 1.

Introduction - Explanation of what is to be measured and the mathematical and scientific ideas which must be employed in the experiment. Note: no "experimental" description in the usual sense is required.

2.

Calculations and Data - Algebraic and numerical calculations must be shown. This makes it possible for graders to trace errors. Where many calculations of the same type are required, a single sample calculation should be shown.

3.

Error Analysis - give a linear or R.M.S. propagation of errors treatment (see expressions in Ch. 2, GNS) and show percent error in final result algebraically and numerically. Mention precision vs. accuracy problems for experiments.

4.

Discussion - What has been learned about the molecules studied and how does this relate to chemistry?

5.

Critique - Discuss the strengths and weaknesses of the experimental technique which was used in relation to other ways of measuring the same quantity.

6.

Questions - Answer any and all questions in the write-up.

In addition to these, ten percent of each report score is based on the organization, appearance and general write up. The presentation of figures, graphs, and equations (numeration, labels, proper scale, etc.) is significant to the overall appearance. Comparison to Literature – Many values can be found in either the CRC Handbook of Chemistry and Physics or at webbook.nist.gov/chemistry/ Note: * No. 2-6 have most of the points. * Conciseness is urged. Page Limitation (not counting raw data pages): 1-period report - 5 pages 2-period report - 12 pages

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Chemistry 111

Report Format Phase Diagram

1. Introduction A. Purpose of experiment (one paragraph). B. Procedure (brief, no longer than one paragraph). C. Discussion - Theory, Phase rule phase diagram, Raoult’s Law, restriction to constant pressure. 2. Data and Calculations A. Tabulate results and clearly graph results. B. Barometer corrections, interpolation. 3. Discussion A. B. C. D.

Compare phase diagram to literature. Compare azeotrope to literature value Sources of errors. Deviation from Raoult's law, azeotropes, molecular interpretation. General comments.

4. References

Heat Capacity Ratio 1. Introduction A. B. C. of

Purpose of experiment (one paragraph). Procedure (brief, no longer than one paragraph). Discussion - Theory (brief summary of experimental and theoretical ideas), degrees freedom, equipartition of energy, etc.

2. Data and Calculations A. Do "CALCULATIONS" on p. 113 of the GNS textbook. Show one sample calculation and tabulate the rest. 3. Discussion A. Cite literature values and compare results. B. Discussion of errors. C. General comments. 4. Questions 5. References

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Chemistry 111

Report Format N2O4 Experiment

1. Introduction A. Purpose of experiment (one paragraph). B. Procedure (brief, no longer than one paragraph). C. Discussion - Theory (show a basic understanding of any derived formulas), derivations of Kp, H, S 2. Data and Calculations A. Show one sample calculation and tabulate the rest. Calculate Kp, G, H, and S at each temperature. B. Include plot of Kp vs. T and 1n Kp vs. 1/T. C. Do not forget barometer correction and thermometer calibration. 3. Discussion A. B. C. D.

Cite literature values and compare results. Discussion of errors. Discussion of underlying assumptions. General comments.

4. Questions 5. References H Vaporization 1. Introduction A. Purpose of experiment (one paragraph). B. Procedure (brief, no longer than one paragraph). C. Discussion - Theory, Clapeyron equation, assumptions, limitations of theory, temperature dependence of z and H. 2. Data and Calculations Tabulate results whenever possible and clearly graph results. 3. Error Analysis A. Use graphical and least-squares methods to determine error in H and propagation of errors for all other results. B. Show sample calculation for each value and tabulate the rest next to the results, i.e. 44.6  0.2 gms. 4. Discussion A. Cite literature values and compare results. B. Speculate on identity of unknown and give the reasons that lead you to this conclusion. C. Sources of errors. D. General comments 5. Questions

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Chemistry 111

Report Format

6. References Bomb Calorimeter 1. Introduction A. Purpose of experiment (one paragraph). B. Procedure (brief, no longer than one paragraph). C. Discussion - Theory (show and understanding of pertinent expressions). Derive expressions for w and Eg. 2. Data and Calculations A. Graphs - Calculate T. B. Complete sample calculation of w and g. C. Tabulate results neatly. 3. Error Analysis A. Complete example. B. Estimate errors in all parameters. C. Tabulate errors in same table as results. 4. Discussion A. Identify unknown and provide reasoning. Cite and compare data to literature values. B. Sources of errors, assumptions, etc. C. Critique - suggestions for improvement, accuracy and precision, comparison with other techniques. 5. Questions 6. References 2.

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Chemistry 111

(1 period)

Heat Capacity Ratio

Notes on Heat Capacity Ratio Experiment References: C. W. Garland, J. W. Nibler, and D. P. Shoemaker, Experiments in Physical Chemistry. 8th Edition. McGraw-Hill Publishing Co., N. Y., Experiment 3, pp. 106-114.

Notes:

The experiment is done according to GNS, with the following changes: 1. Take measurements on argon, nitrogen, and CO2. Helium gives rather poor data because of its high thermal conductivity. 3. Our setup uses valves instead of hose clamps, since the latter are inconvenient to open and close. We now use an electronic pressure transducer for the pressure measurements. We used to have a mercury rather than an oil manometer. Ignore the oil-to-mercury conversion mentioned in the text on p. 113. 4. In addition to the questions on p. 113 of GNS, estimate the effect on Cp/Cv of a very high heat conductivity, as is the case for helium.

5. Actual values for comparison and error analysis of the heat capacity ratios you have obtained can be found at http://www.engineeringtoolbox.com/specific-heat-ratiod_608.html. Be sure to cite this website as a reference if you use its data.

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Chemistry 111

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Heat of Dissociation

The Heat of Dissociation of N2O4 References: 1. Daniels, Mathews, Williams, Bender, Alberty, Experimental Physical Chemistry, 5th. Ed., p. 97-100. [4th. Ed. p. 110-112]. 2. Glasstone, Textbook of Physical Chemistry, 2nd. Ed. 3. Shoemaker and Garland, Second Edition, p. 46-69.

Object: In this experiment you will measure the apparent molecular weight of equilibrium mixtures of N2O4 and NO2, arriving at the dissociation constant, (T), and equilibrium constant, Kp(T), as a function of temperature. Measurements are made by a gas density method due to Dumas. G°, H°, and S°, the standard free energy, enthalpy and entropy changes for the dissociation reaction are obtained. Principle of the Experiment: A bulb terminated with a stopcock is evacuated and weighed. It is filled with dry air at a known pressure and temperature and reweighed. The volume is obtained from a knowledge of the density of the air (by using Mair = 29 g/mole and PV = (m/Mair)RT). The re-evacuated bulb is filled with gaseous N2O4 -NO2 and equilibrated in a thermostat at a known temperature and atmospheric pressure. The closed-off bulb is reweighed and the apparent molecular weight of the gas calculated. The experiment is repeated over a range of temperatures and the degree of dissociation and equilibrium constant are calculated over the temperature range. G° may be calculated at any temperature from a knowledge of the equilibrium constant, Kp(T). H° is obtained from the slope of the ln Kp(T) vs. l/T plot. S° may then also be calculated at any temperature. Procedure: Obtain a Dumas bulb with a capillary stopcock from the teaching assistant. If the stopcock does not turn freely, have your teaching assistant regrease it. Evacuate the Dumas bulb with the mechanical pump, warming the outside of the bulb gently with a bunsen burner flame if you observe moisture on the inside. Do not warm near the stopcock. Wear goggles whenever you handle an evacuated bulb. Because of the great length and small diameter of the capillary, it will take about 5 minutes to effectively evacuate the bulb. Weigh the cooled bulb. Fill the bulb with dry air from the drierite column in the hood. Do not forget to remove the cork and pinch clamp before filling and to replace them afterwards. Do not leave the drierite column open to the atmosphere. Note the temperature and measure the barometric pressure. Remember, whenever filling your bulb with gas, sufficient time must be allowed for pressure equilibrium to be obtained. Thirty seconds is safe. Weigh the bulb again. Weighings must

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Chemistry 111

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Heat of Dissociation

be made to the nearest 0.1 mg. Evacuate the bulb (5 minutes) on the mechanical pump and weigh again. If the weights of the empty bulb do not agree to within 0.2-0.3 mg, something is wrong with your procedure. Ask the teaching assistant to demonstrate the filling of the bulb with N2O4 from the cylinder in the hood. N2O4 is a very noxious gas. Significant quantities should never be allowed to escape in the lab. The procedure of filling is briefly as follows. Open main valve, crack the needle valve to dispel air from the tygon tube. Attach the bulb to the tygon tube so that glass touches metal. Open the needle valve wide (2-3 full turns). Open the stopcock valve gradually. When completely open, allow to remain open for 30 seconds. Be careful not to tilt the cylinder to avoid getting liquid N2O4 in the glass bulb. Close stopcock, close needle valve, close main valve, remove bulb from tubing. While still under the hood, open stopcock to dispel the overpressure of NO2 in bulb. When the gas has stopped leaking out, close stopcock and clamp the bulb in the large beaker of water. Do not immerse the stopcock. The bath should contain a thermometer and a large stirring rod. Warm the bath while stirring to about 30°C. During the warming process open stopcock for a few seconds occasionally to allow some of the built-up pressure to diminish. Maintain the bath at a fixed temperature in the vicinity of 30°C for 2-3 minutes, open the stopcock for 30 seconds while maintaining temperature to allow pressure to equilibrate. Record bath temperature. Remove bulb, dry it with a clean towel and allow it to equilibrate to room temperature. Meanwhile, remeasure the barometric pressure. Weigh and record the weight of the bulb to the nearest 0.1 mg. Repeat the process at about 40°, then 50° and finally at 60°. Never open the stopcock unless the temperature of the bulb is higher than it was at the time you previously opened the stopcock. Why? After you have finished your measurements, evacuate the bulb twice with the water aspirator. Never pull NO2 through the mechanical pump. When the yellow color of NO2 is absent or just barely perceptible, complete the evacuation on the mechanical pump, warming as before to dispel any condensed moisture. Make sure you let air back into the bulb before leaving. Why? Calibrate your bath thermometer using ice slush and boiling water. Calculations: In order to obtain the thermodynamic data, you must find the Kp for the reaction at each temperature using your experimental data. First, set up an ICE table for the dissociation of N2O4, assuming 1 mol of N2O4 is present initially and that it is completely undissociated. The equilibrium line of the table will give you the number of moles of each gas at equilibrium at a given temperature in terms of the variable xT. You will then obtain the reduced mass of the gases using the equation: (1) M = (m1 + m2) / (n1 + n2)

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Chemistry 111

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Heat of Dissociation

Where M is the reduced mass of the gas sample and m and n are the molecular weight and number of moles (derived from the last line of the ICE table) for the two gases, respectively. For the next step, write out the equilibrium expression for KP (equilibrium constant at various temperatures but at constant pressure) in terms of the partial pressures of each gas. Each partial pressure value must be divided by the standard pressure P° (1 atm in this case) in order to make KP unitless. Afterwards, find the partial pressure of each gas by multiplying P0 (the pressure of the mixed gas sample, which is always equal to room pressure in this experiment) by the moles of each gas divided by the total moles (moles N2O4 + moles NO2). This will give you a partial pressure for each gas in terms of P0 and xT. Now plug these partial pressures into the KP expression to obtain an equation for Kp in terms of xT, P° and P0. Finally, we must find an equation to solve for xT, which we can plug in to the KP equation to solve for the equilibrium constant. By assuming the reduced mass is the overall molecular weight of the gases in the sample, we can utilize the ideal gas law to obtain the equation: (2) P0V = (mT / M)*RT Where P0 is room pressure, V is the volume of the Dumas bulb (both are constant throughout this experiment), mT is the mass of the gas sample at a given temperature, M is the reduced mass in terms of xT, R is the universal gas constant, and T is the experimental temperature. Since P0V is held constant throughout the experiment, this equation can be set equal to the equation you used earlier to solve for the volume of the Dumas bulb: (3) P0V = (m/ Mair)*RT0 In which m is the mass of the bulb filled with air, Mair is 29 g/mol, and T0 is room temperature. The constant R will cancel out on both sides, after which you can rearrange the equation and solve for xT in terms of mair, mT, T, and T0. For each temperature, you can now plug in the appropriate mass of the gas sample and the experimental temperature to obtain an xT value, which can be plugged into the Kp expression to solve for Kp at each temperature. Calculate Kp and G° for each temperature of your measurements and tabulate. Plot Kp vs. T and ln Kp(T) vs. l/T and draw smooth curves through the experimental points. Calculate the mean H° for the experimental temperature range. If your plot of ln Kp(T) vs. l/T shows a decided curvature, calculate H° at 35°, 45°, and 55°. Calculate G° at 35°, 45°, and 55° from one of your plots. Calculate S° at 35°, 45°, and 55°. Tabulate G°, H°, and S° for these three temperatures. Analyze your experiment for errors (including use of the perfect gas law) and quantitatively estimate the accuracy of all your calculated results.

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Chemistry 111

(2 periods)

Heat of Vaporization

Heat of Vaporization by a Static Method References: 1. C. W. Garland, J. W. Nibler, and D. P. Shoemaker, Experiments in Physical Chemistry. 8th Edition. Experiment 13, pp. 199-207. 2. W. J. Moore, Physical Chemistry. Fourth Edition. Prentice-Hall, Englewood Cliffs, NJ, 1972. Chapter 7. 3. R. S. Livingston, Physico-Chemical Experiments. Third Edition. Macmillan, New York, 1957. pp. 67-70. 4. Smith and Menzies, J. Am. Chem. Soc. 32, 1413 (1910). 5. F. Daniels, J. H. Mathews, J. W. Williams, P. Bender and R. A. Alberty, Experimental Physical Chemistry. Fifth Edition. McGraw-Hill, New York, 1956. pp. 47-50, 370-373. 6. F. Daniels and R. A. Alberty, Physical Chemistry. Fourth Edition. Wiley, New York, 1975, pp. 136-141.

IMPORTANT: Make certain that you have discussed vacuum system techniques (see GNS pp. 587-600) with the T.A. and thoroughly understand the procedures. This experiment deals with equilibrium in a one component, two phase system. Application of the Gibbs phase rule shows that at a fixed temperature there is only one pressure at which the two phases co-exist in equilibrium. The enthalpy change, Hvap, for A(liq.) = A(gas) (p,T) will be measured. The method employed will use the “Isoteniscope” of Smith and Menzies. Procedure: Assemble the apparatus as shown in the diagram, placing the thermometer and isoteniscope bulbs in close proximity. The stirrer should operate freely. Fill the isoteniscope bulb about two-thirds full with the liquid sample, making sure that the liquid column in the U-tube is about two inches long. Grease the ball and socket joint with silicone grease. (N.B.: If grease gets into the liquid in the isoteniscope bulb, then your results may be seriously in error.) Stopcock 1 should be open to the ballast and stopcock 3 closed. (Stopcock 3 should always remain closed.) Stopcock 4 should be open to its capillary. Stopcock 2 should be open to the gauge. With the bath at about 70°C, establish your ability to hold the temperature constant to 0.1°C for several minutes. Slowly evacuate the isoteniscope until liquid in bulb boils, and a steady stream of vapor bubbles passes through the liquid in the U-tube. Do this by 11

Chemistry 111

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Heat of Vaporization

slowly opening stopcock 4 to the pump. Be careful not to let any liquid boil over into the vacuum manifold. Continue for 3 minutes, or until you believe all air trapped in the bulb has been driven out. Measure vapor pressure by closing off the pump and opening capillary leak at stopcock 4. Allow vapor pressure to build up slowly until the liquid levels in the two arms of the U-tube are equal. At this instant close stopcock 2 and record the pressure. Open stopcock 2, boil out for another minute, and keeping the temperature constant as possible, repeat the measurement to ascertain if complete air removal has occurred. If so, take a series of readings at approximately 5° intervals down to  30°C. Balance is most easily approached from the low system-pressure side by allowing air in through capillary. In this experiment never open the capillary to the manifold more than a crack. The bath must be held at constant temperature for at least a minute before a reading to allow liquid in the bulb to equilibrate. If liquid in the U-tube is drawn back into the bulb, or if it evaporates, air will get into the bulb and the degassing procedure must be repeated. Calculations: Tabulate original data, corrected values of pressure, P (Torr or mmHg) and T(°C), ln P/P0 , l/T (K-1). Plot on a suitable scale a) The vapor pressure vs. temperature b) ln P/P0 vs. 1/T Find the constants in ln P/P0 = - A/T + B from your graph. Take compressibility into account (GNS, pp. 200-201) using the following equation 9 P TC T2 (1  6 C2 ) Z  1 T 128 PC T and report the average value of Hvap over the temperature range encountered in the experiment. Estimate the normal boiling point from your graph, then calculate Svap, the "entropy of vaporization". Use Trouton's rule to estimate the normal boiling point and see how it compares with your data. Error treatment — Draw on plot (b) the greatest and least slopes that are determined by the data points. Use them to estimate error in vap. Also use the least-squares method to estimate the error. See GNS pp. 36-38 and 663-672. Questions: 1. Suppose there is 1% air left in the isoteniscope bulb at 70°C. What error is introduced into Hvap? 2. What is the significance of a slight curvature in plot b? 3. There is a temperature gradient established between the liquid in the U-tube and condenser. Will this affect the pressure reading?

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Chemistry 111

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