Derek 1. Hill, Stephen J. Moss,
University of Hong Kong Hong Kong, 0. C. C. and Robert 1. Strong Rensselaer Polytechnic Institute Troy, New York
II
Heat of Reaction in Aqueous Solution By Potentiometry and Calorimetry 1.
A metal displacement reaction
Thermochemical measurements in the undergraduate physical chemistry laboratory play an important role in developing the students' appreciation of the laws of thermodynamics. Consequently, measurements of heats of reaction and experiments involving applications of the Gibbs-Helmholtz equation to both homogeneous and heterogeneous equilibria are common. I n addition, experiments are available in which temperature coefficients of electromotive force of galvanic cells are used to deduce via the GibbsHelmholtz equation the enthalpy change of the cell reaction, on the basis of the familiar relationships: a(aG) = -as aT
= AG - AH, and AG = -n5E,
so that
To the best of our knowledge however, no such experiment combined with a direct calorimetric check appears in current laboratory textbooks or manuals. Such a combined approach to one particular reaction offers distinct advantages, which have been realized in the experiments described in this paper. Many textbooks of physical chemistry (1-4) and electrochemistry (5-8) make use of the two approaches in discussing the Gibbs-Helmholtz equation, but generally the examples quoted are unsuitable for routine laboratory instructional purposes. These sources generally show agreement between the two methods for a particular reaction which are not better than 1% and sometimes considerably worse. The agreement we have obtained, using the relatively unsophisticated equipment and techniques described, compares most favorably. It is desirable that all equipment should be of a readily available kind. Each of the two experiment,^ should be completed easily within a normal three-hour laboratory period and should be capable of showing a good degree of concordance. The particular reaction to be chosen should be rapid, should have simple stoichiometry, and should be such that, as far as possible, conditions could be identical in the two separate experiments so that results would be directly comparable. It should provide an electrochemical cell with a comparatively high temperature coefficient of electromotive force and should also be suitable for simple calorimetry. The availability of complete literature data for comparison is advantageous. A reaction which satisfies these requirements is: Zn
-
+ 2Ag+ (aq)
Zn2+ (aq)
+ 2Ag
Silver nitrate is unsuitable for this experiment, because zinc slowly reacts with nitrate (9),giving a prolonged temperature rise in the calorimeter. Silver acetate may be used instead, but silver sulfate is preferred because it can be handled more easily. This reaction has been used in our teaching laboratories for about six years with mixed results, and the present communication describes recent improvements in technique which make it suitable for adoption as a routine undergraduate experiment. Students working in pairs, with good organization of their time, should be able to complete both parts of the experiment in about three hours. The Experiment
Any high grade potentiometer reading to 0.1 mv or better will suffice. Figure 1 shows a suitable design of cell, having an agar-KKOa salt bridge held firmly in place by fine porosity frits. The salt bridge is prepared by adding 3% agar to a solution of KXO3 which has been saturated at 0°C lo avoid crystallization of the salt during measurements in an ice bath. The bridge is protected when not in use by filling both compartments with saturated KN03 solution. Provision is made for more than one electrode in each compartment and for tank argon inlet and outlet. The gas is thermostated and presaturated by passiug it through a glass spiral and water bubblers immersed in the appropriate thermostat. Solutions of ZnSOl and Ag2S04 are prepared from reagent grade salts, standardized, and adjnsted to 0.0125 d l .
w Figure 1. Cell. o, orgon inlet; b, argon outlet; c, Rne porosity frih; d, agar-KNOS gel; e, thermometer; f, thermostot level.
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The provision of reproducible electrodes is the only difficulty in this experiment, but we have found that good results can he obtained with electrodes freshly prepared in the following way:
position. Because of poor thermal contact through the glass tube, only a trace of solution should be allowed to remain inside. The whole operation should take less than 20 sec. Temperature readings are continued for another 10 min. At this point, the calorimeter is Zinc. Take a cast stick of the metal which shows no outward cooled to about the initial temperature by placing a signs of having been previously bent, and insert it through a cork little solid COz in the tube. During this time the or bung which fits the cell. Dissolve all scale and surface layers dummy heater is energized and as soon as a suitably from the lower inch of the metal in 30% H2S04,rinse with dietilled water, and immediately place in saturated H ~ X ( N O ~ ) ~steady state is attained by the calorimeter, temperature solution for 2-3 minutes. Rinse with distilled water and then readings a t l-min intervals are resumed. After 10 with 0.0125 M ZnSOi. dace in the cell. and start the areon flow. such readings the current is switched to the calorimeter heater on a full minute, and current and voltage are torily for the duration of the experiment. It can be used again read at minute intervals until the initial temperature several times following immersion in 30% HBOc but does have rise is approximated. The heater is switched off on a a limited useful life. In our experience this is a more convenient full minute and temperature readings are continued method of amalgrtmating the surface than is usually described for another 10 min. in standard texts, and in addition gives more satisfactory results. The temperature rises are corrected graphically in The emf between pairs of freshly prepared electrodes immersed in ZnSOl solution is generally below 0.1 mv and often as little as the normal way as indicated in Figure 4. The water 0.01 mv, but does increase with time. equivalent of the calorimeter and contents is given by Silver. Insert a length of heavy gauge silver wire through a cork or bung which fits the cell. Anodize the lower inch of the metal for about 10 min in 10% KCN solution to obtain a clean surface. Place in a bath (10) containing 1% AgNOa in 90% methanol, and electroplate at low current density for about an hour. Rinse with distilled water and then with 0.0125 M Agp SOc place in the cell, and start the argon flow. These electrodes should be stripped in KCN and replated before each use. The emf between pairs of freshly prepared electrodes immersed in Ag2S04solution is usoslly less than 0.1 mv but increases slowly with time.
The cell
w =
V X I X t
4.184 X ATw
where V is the average potential across the heater in volts, I the average current through it in amps, and t the time of heating in seconds. The heat of reaction is given by
where n is the number of moles of Ag,SO, initially present in the calorimeter.
is placed in a 25'C thermostat and allowed to reach thermal equilibrium. The emf is then measured several times over a 15-20 min period, noting the reading of the thermometer in the agar bridge to within O.l°C. Corresponding measurements are then made with the cell, first in an ice-water bath and then at 40% Throughout the measurements a steady flow of tank argon is maintained through the solutions in each compartment. Even so, the electrodes deteriorate fairly rapidly, and it is advisable lie complete all measurements within a 2hour period. The calorimeter, shown in Figure 2, has a capacity of about 600 ml. Although designed for reactions between solutions, it has been found eauallv " satisfactorv for this experiment. The heating circuit is of simple type, consisting of a &volt battery supply, together with an ammeter and voltmeter calibrated to better than 0.2%. The nichrome wire heater of about 10-ohm resistance is enclosed in a deionized water-filled sheath and a duplicate heater is provided to act as a dummy load to stabilize the battery supply. Four hundred ml of standardized 0.0125 M AgzS04 solution are carefully measured into the clean dry calorimeter. About 20 g (a large excess) of finely powdered reagent grade zinc are placed above the plug in the glass tube A. A reasonably rapid and constant rate of stirring is maintained and the assembled calorimeter is allowed about 20 min to reach a suitably steady state. The temperature is then recorded at 1min intervals for about 10 min, and on a full minute the zinc is introduced by removing the plug and raising and lowering the tube A quickly to flush out all the powder. The tube is raised above the solution level to drain, the plug is replaced, and the tube lowered to its original
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J Figure 2. Calorimeter. A, tube with ground gloss plug and loose fltting rubber stopper; o, Dewor flask; b, cotton wool pocking; c, wooden case; d, asbestos lid; e, P e n p e i cover; f, Beekmonn thermometer; g, stirrer with Teflon bearing; h, heater.
heat effect. The major problem is the difficulty of obtaining a steady state in the calorimeter in the limited time available for a student experiment. I n view of the deviation from standard states, the difficulty of maintaining the electrodes in their initial conditions, the sensitivity of the measured hE/hT to very small errors in & a t the three temperatures, and the simple design of the calorimeter, we regard the agreement as very satisfactory. I n particular, concordance of the electrochemical result and the thermochemical check on the value of AH is very good. Much improved calorimetric techniques could be employed if desired; but because both experiments are intended to be operated in normal teaching laboratories and because of possible interference from side reactions,
TEMPERATURE ('C) Figure
3.
Cell emf versus temperature.
Results and Discussion
Potentiometric results are shown in Figure 3. We used four electrodes of each type in one cell in two separate runs, so that we have results from a total of 32 electrode pairs. Run 1 involved first use of the electrodes, and Run 2 subsequent use of the same electrodes freshly prepared. Average results (per mole Zn) from all 32 measurements fall within the ranges shown in the table. We have been unable to find the value of y* for Ag2S04in the literature, but using the same electrode preparation, and substituting 0.025 M AgN03 for 0.0125 M Ag2SOa,at 25°C we recorded 1.5322 1 0.001 volts in one run with 16 electrode pairs, and 1.5331 + 0.0002 volts in a subsequent run using 8 particularly concordant pairs. These values are in excellent agreement with the calculated cell emf which is 1.533 volts, using &Os.2+,z. = -0.763 and &oA,+,A. = 0.799 volts (12); y+ (ZnSOr) = 0.355 and y+ (AgNOJ) = 0.846 [interpolated from handbook data (IS)]. Typical calorimetric results are shown in Figure 4. Five such runs yielded values within the range AH = -86.7 zt 0.6 kcal per mole Zn. Figure 4 also shows that the addition of 20 g powdered Zn in 400 ml deionized water with the usual manipulation of the ground glass plug and tube A produces no significant
rimetry are the slow evolution of hydrogen from the solution subsequent to the displacement of silver, which is in fact clearly observable some hours later, and the possibility of alloy formation between zinc and silver (9). Our results suggest that neither effect is significant in comparison with the over-all experimental error. The use of electrochemical cells with liquid junctions is of course unsatisfactory for precise thermodynamic interpretation, but some recent comments accepting these limitations may be noted (1.4). We know of no reaction suitable both for a cell without liquid junction and for simple solution calorimetry. The real significance of the experiment to the student is the close concordance of the two independent approaches to a measurement of fundamental importance. The direct check provided by the calorimetry on the value of AH derived from the potentiometry, with average results deviating less than 1%, is better than most examples in the literature previously quoted ( 1-8). Literature Cited (1) DANIELS, F., AND ALBERTY, R. A,, "Physical Chemistry," 2nd ed., John Wiley and Sons, Inc., New York, 1961, D. 389.
Results
Gzes
(volts)
b&/bT (mv deg-') AS (cal deg-') AG (kcal) AH (kcal)
Run 1
Run 2
1.5298 (zt0.0005) -1.08
1.5292 (ztO.001) -1.15
-51.4 1 1.6 -70.5 zt 0 . 1 -85.9 0.5
*
Literature values for standmd states (. 1 1.) 1.562
-1.09
-50.34 -72.04 -87.05 TIME
(minutes)
Figure 4. Colorimetry: A, Zn addition to deionized water; reaction AT.; C, water equivalent AT,
B,
heat of
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(2) GLASSTONE, S., A N D LEWIS, D., "Elements of Physical Chemistry," Macmillan and Co. Ltd., London, 1960, p. 461. (3) MOELWYN-HUGHES, E. A,, "Physical Chemistry," 2nd ed., Pergamon Press, Oxford, 1961, p. 1091. (4) ROSE, J., "Dynamic Physical Chemistry," Pitman and Sons Ltd., London, 1961, p. 549. (5) GLASSTONE,S., "The Electrochemistry of Solutions," 2nd ed., Methuen and Co. Ltd., London, 1937, p. 288. (6) K O R T ~ ~G., MAND , BOCKRIS, J. O'M., "Textbook of Eleetrochemistry," Elsevier Publishing Company, New York, 1951, Vol. I, p. 244. (7) MACINNEE,D. A,, "The Principles of Electroohemistry," Dover Publications Inc., New York, 1961, p. 113. (8) POWER,E. C., "Ele~trochemiatry," Cleaver-Hume Press Lt,d., London, 1956, p. 79.
(9) MELLOR,J. W., "A Comprehensive Treatise on Inorganic and Theorc:tieal Chemistry," Longmsns, Green and Co. Ltd., London, 1923, Vol. 111. P. 319. (10) F I N D I ~ Y , A., '.'practical Phy&al Chemistry," 8th ed., Longmans, Green and Co. Ltd., London 1954, p. 254. (11) ROSSINI,F. D., ET AL., "Selected Values of Chemical Thermodynamic Properties," Circular of the National Bureau of Standards 500, Washington, 1952. (12) LATIMER,W. M., "The Oxidation Potentials of the Elements and their Values in Aqueous Solution," 2nd ed., Prentioe-Hall Inc., Englewood Cliffs, New Jersey, 1952. (13) PARSONS, R., "Handbook of Electrochemical Constants," Butterworths Scientific Publications, London, 1959. (14) IYES, D. J. G., AND JANZ, G. J., "Reference Electrodes," Academic Press, New York and London, 1951, p. 48.f.
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