Heat of Reaction in Aqueous Solution
Stephen J. MOSS
and Derek 1. Hill University of Hong Kong Hong Kong, B. C. C.
By Potentiometry and Calorimetry
I
11.
A r e d o x reaction
T h e instructional purpose of this experiment is the same as that described in Part I ( I ) : to measure the enthalpy change in a chemical reaction by determination of the temperature coefficient of emf of a suitable electrochemical cell, and to make an independent direct calorimetric check on the result. The version presented here was developed to further improve the agreement between approaches and the reproducibility within each approach by avoiding inevitable problems associated with solid metal-metal ion electrodes, and heterogeneous reactions at a metal surface. The same criteria have been applied in the choice of reaction: rapidity, simple stoichiometry, direct comparision between potentiometry and calorimetry, reasonably high temperature coefficient of electromotive force, and suitability for simple calorimetry. The reaction chosen is Fe(I1)
+ Ce(IV)
-
Fe(II1)
junction is made at the tip of each capillary. We have found that cells with agar salt bridges in contact with the solutions are unsatisfactory, as are a number of cell designs involving liquid junctions at ground glass surfaces. Reagent grade chemicals are used to prepare standardized stock solutions of approximately 0.40 N
+ Ce(II1)
in 1 N H2SOI solution. It is necessary to maintain a substantial concentration of sulfuric acid in the solutions, to avoid significant changes in the nature and proportions of the complex species during temperature variation. This reaction has been used with good results in our teaching laboratory for one term by 18 pairs of students. Again, with good organization of students' time, both parts can be completed in about three hours. The only disadvantage of this reaction compared with that described in Part I is the almost complete lack of pertinent literature data. The Experiment
A high grade potentiometer is used to measure the emf of the cell shown in Figure 1. I n this the liquid
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Figure 1. Cell. a, ground glarr plugs; b, 1 mm capillaries reduced to 0.5 mm at tip; c, thermometer pocket; d, 1 N H2SOli e, glass rod sup-
port; f, thermostot level.
TEMPERATURE Figure 2 .
TIME (minutes)
(TI
Cell emf versus temperature.
Figure 3.
Calorimetry.
A, heat of reaction AT.;
B, water equivalent
ATw
FeS04, 0.35 N Ce(SO&, 0.20 N Fez(S04)a,and 0.20 N Ce2(SOJl all in 1.00 N HISOa. I t is important to avoid heating Ce(SOa)zsolution a t any stage, otherwise precipitation may occur on mixing with Cez(S04)asolution. These solutions are mixed and diluted with 1.00 N H,SO, as appropriate just before each experiment to give 50 ml solution 0.200 N in Ce(1V) and 0.050 N in Ce(III), and 50 ml solution 0.200 N in Fe(I1) and 0.050 N in Fe(II1). Deterioration of the Fe(III), Fe(I1) solution occurs on standing and gives low values of cell emf. The stock solutions should be standardized regularly. The cell
is assembled with a bridge solution of 1.00 N H2S04. Clean P t electrodes are labeled and placed in the appropriate half cell solution for a few days before use, and always stored in the same solutions when not in use, to aid stability (9). The emf of the cell is measured at 25', 0°, and 40°C for periods of about 1.5 min after thermal equilibration. The calorimeter and heating circuit have been described in Part I. We use 0.200 N FeSOl and 0.200 N Ce(SO& solutions, both in 1.00 N HzSOa. Four hundred ml Fe(I1) solution (a large excess) are measured into the clean dry calorimeter. About 100 ml Ce(1V) solution are weighed into the tube A and the calorimeter is assembled. After equilibration, the solutions are mixed, taking care to flush and drain all solution from tube A before replacing the ground glass plug. The water equivalent is then determined as in Part I. From the weight of Ce(1V) solution and its density, the volume is obtained; and AH is calculated per mole Ce(1V). Results and Discussion
Data from 18 pairs of students have been collected and checked. Their results are summarized in the table. Student calorimetric checks average AH = -24.1 1
0.9 kcal/mole (mean deviation = 0.5), providing an over-all agreement withim 2.5%, for the whole class. I t should be noted that the students all performed the experiments with only normal supervision. After considerable experience with this method the authors repeated the experiment and obtained virtually the same values as the students had. A slightly larger value (-0.603 mv/deg) was obtained for a&/bTwhich gave AH = -23.60 kcal. Calorimetric results were similarly in agreement (AH = -23.75 *O.l5 kcal). Changing to 2 N H2S04or reversing the proportions of Fe(I1) and Ce(1V) had no effect on the calorimetric result. Data are illustrated in Figures 2 and 3. Unfortunately it is possible to make only one rough comparision with literature data. This concerns Ezss. In 1 N H2S04,E0 of [Ce(IV), Ce(III)] is given as 1.461 v (3) and &Of as 1.44 v (4, 5) although the latter figure may be an error for 1.46. EO'of [Fe(III),Fe(I1) 1 is given as 0.68 v (6, 7). Thus if activity coefficient ratios were fairly close to unity, Gzo8is calculated to be in the range 0.85 *0.02 v and AG = -19.6 1 0 . 5 kcal/mole. We have no direct evidence that liquid junction potentials are eliminated by 1N H2SOI,hut the addition of (NH&S04 gives results summarized in Figure 4. With 1 N H2S04alone in the bridge, slight streaming occurs through the plugs and capillaries which approximates a crude flowing junction. This may account for the greater stability of measured cell emf with this bridge, which we prefer. The use of such cells to obtain thermodynamic data has been discussed briefly in Part I (1). Student Results Average
&m ( v )
a&/W(rnv/deg)
0.8433 -0.596
Derived Quantities AS ( e a l i e g ) -13.74 AG (kcal) -19.45 AH (kcal) -23.54 Colorimetric Measurement AH (kcal) -24.1
Range
Mean Deviation
+0.003 10.036
0,0019 0.02
+0.80 +0. 25 10.30
... 0 1
+0.9
0.5
. ..
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We have considered the effect of heat of dilution of the excess Fe(I1) solution in the calorimetry. Unless the calorimeter is equilibrated for very much longer periods than is feasible in student work, the effect is not adequately measurable. Equilibration overnight shows however that a correction for this effect would he justified in a more refined calorimeter. A possible resolution of this is to use only a slight excess of Fe(I1) solution in a calorimeter of slightly smaller volume. Although rather more preparation is needed initially since several stock solutions must be prepared and standardized, from the student viewpoint the potentiometric measurements for this reaction are experimentally simpler than thosemade in Part I ( I ) , while the calorimetric measurements are closely comparable. In particular, apart from the conditioning indicated, the electrodes require no treatment before the experiment nor is an inert atmosphere necessary during it. Additionally, the measured emfs at each temperature are very stable and the combined measurements show a very close-to-linear relation to temperature. These features together with the good concordance between the two parts obtained by students makes for a satisfying experin,el>tof fundamental importance. Literature Cited (1) HILL,D. L., MOSS,S. J., AND STRONG, R. L., J. CHEM.EDUC., 42,541 (1965). (2) Ross, J. W., AND SHAIN,I., Anal. Chem., 28,548 (1956). (3) Kunz, A. H., J . Am. Chem. Soc., 53,98 (1931).
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0
50
100
ENH,,I,SO,] in 1N H,SO,, (% saturation at O°C) Figure 4. Variation of in bridge.
h&/hT and AH
with (NH4bS04concentration
(4) SMITH,G. F., AND GETZ,C. A,, Znd. Eng. Chem., Anal. Ed., 10,191 (1938). (5) LAITINEN, H. A., "Chemical Analysis," McGraw-Hill Book Co. Inc., New Yark, 1960, p. 378. (6) SMITH.G. F.. Anal. Chem.. 23.925 (1951). (7) LAITINEN,H' A,, Op. eit., 2'83
p.
