A Laboratory Experiment Yielding Entropy Changes The paramount importance of the entropy concept to thermodynamics requires that its application by our students he reinforced a t every opportunity, including, we believe, via laboratory measurements. Yet a survey of physical chemistry laboratory texts reveals that its calculation is only rarely called for. Even then, determination of entropy changes appears relatively incidental to the experiment and assumes that further material (e.g., chemical equilibria, emf) has already been covered. There appeared to be no experiment having both amajor emphasis on entropy ehanges'and suitable for the time when the entropy function is introduced. An experiment fulfilling these criteria is described below. Although it is quite simple, it yields A S from laboratory measurements directly and works well. Moreover, it capitalizes on the students' tendency to work on calculations for the laboratory report, whereas only half may do such an exercise as a homework assignment. In response to a questionnaire, students indicated that the experiment reinforced this lecture material for them. This eaperiment is offered, in part, to help fill a serious void and, in part, to provoke thought along these lines and to elicit still more useful laboratory experiments concerning entropy. The experiment involves miring two samples of water a t different temperatures and determining the component and net entropy changes for this irreversible thermal equilibration process. Water a t room temperature (Tc) is added to hot water (at TH) in a Dewar (a $2 Thermos refill) and the equilibrium temperature (TF) measured. The same thermometer, with 0.l0C divisions, is used throughout. TH and TF are obtained from temperature versus time plots, thereby correcting for heat losses. Thus the mixing carried out in the Dewar is made effectively adiabatic and so isenthalpic. Taking the heat capacities to he substantially temperature independent aver the 10-20°C interval
AH= O = CDA(TF- TH) + MHCH(TF- TH) + MCCC(TP- Tc) and AS, = CJn Tti..~ l"lUal
CDAis the heat capacity of the Dewar assembly (Dewar, thermometer and rubber stopper). MH(c) and CH(C,are the moles and molar heat capacities, respectively, of hot (cold) water. A typical run involving 293.1 g of hot water a t TH = 329.9 K and 299.4 g of ca!d water a t T c = 298.2 K resulted in TF = 314.4 K. Using average values of 75.265 and 75.227 J mole-' deg-' for CH and Cc,respectively, an average value of 82 J deg-I is ohtained for CDA,the calorimeter constant. CDAis especially sensitive to errors in temperature (reading or other) and so detects poor technique, provided TH - TC exceeds -20DC. The entropy changes are found to be -58.93 (J deg-') for the cooled water, -4.0 for the coaled Dewar assembly and +66.14 far the added room temperature water. Thus the net entropy change, ASN, is 3.2 J deg-I and is indeed positive for this irreversible process for the entire isolated system of 592.5 g water and Dewar assembly. Even if the student happens to obtain a negative CDA,A S N is still positive. Thus the student always sees the experiment work and our primary goals are achieved. (The only 'exception' occurs if a temperature is misread so thoughtlessly that, e.g., TF < TC (!I; this would
-- - n-
.
n .~ o ~ -. n n , .> A,%\, ., .a -.
The rxperimenr, mrrred out indrvidually, is ubviuusly easy 10 inaugurate and perform, inrxprnqivr and tnkrs only I h (of -5Odcrotrd w thermodynnmirrj of Inburamry time in this second year o>ursefor -15u srudrnrr. ( A ropy of thr wrilc-up prot,idrd the srodenrr i r wadnl,le m reque\t., Clearly, this could also be a dcmonsrration rxperimcnr, perhaps with n dig~ral ieadout displayed t o the class for theirobservation and recording. Selley, N. J., "Chemical Energetics,"Edward ArnoldLtd., London, 1971, p. 71; Selley, N. J., J. CHEM. EDUC., 49,212 119721. -, A. D. J o r d a n University of Alberta A. H. Kalantar Edmonton, Canada T6G 2G2
.--
Volume 55, Number4 March 1978 / 183
