CPT Lynn P. Beaulieu United States Military Academy West Point. New York 10996
A General Chemistry Thermodynamics Experiment
A common problem in general chemistry laboratory programs is that, except for Hess' Law experiments, simple yet accurate thermodvnamics exueriments are virtuallv. imoos. sihle to perform without resorting to bomb calorimeters or other expensive laboratory experiment ( 1 - 4 ) . This experiment provides the student with the opportunity to do experimental thermodynamics, and to calculate those thermodynamic values which usually cannot be determined with the simple equipment available in a aeneral chemistry laboratory. The experiment consists of meisuring the voltage of a battery a t different temperatures, and requires items common to most general chemistry laboratories: p H meter, glassware, metal foil electrodes, and solutions of readily available salts. The voltage measurements are made a t room temperature, in a cold water bath and in a hot water bath. The difference in voltage between the high and low temperature measurements can be related to the change in entropy, AS, by the relationship (5)
where C is the Gibbs free energy and T is the absolute temperature, a t constant pressure, P. This equation can be written
since individual data points will be measured, and the atmospheric pressure will b e assumed to be constant. Using the relationship between Gibhs free eneray and the voltage of an .~ electroche&cal cell AG = -nFb
(3)
substitution into eqn. (2) yields
A8
(51 AT where F is Faraday's constant and n is the number of moles of electrons transferred in the cell reaction. The battery consists of Ph/Ph2+ and Cu/Cu2+ half-cells connected hy a salt bridge. When equal concentrations of the Ph2+ and C U ~solutions + are used in hoth cells (0.50 MI the second part of the Nernst equation becomes zero. Assuming the concentrations (activities) of the Pb2+ and C U ~are + eaual ( f i ) , and assuming that roomtemperature is 25"C, the standard voltage of the cell will be equal to the measured voltage and will he 0.463 V (7). A temperature change of 55-60°C should produce a voltage change of 25-30 mV, a shift which is sufficiently large to be measured on the millivolt scale of astandardpH meter. More precise and extensive work, including other half-cell combinations which have considerably less change in voltage, could he accomplished on a p H meter with an expanded scale capahility; however, this equipment is not readily available in most general chemistry laboratories. If sufficient standard p H meters are available for each individual or team, many data points can be collected while the battery is slowly heated or cooled. The student can plot 6 versus T , and determine AS AS = n F -
from the slope of the line, X I A T , and eqn. (5). If, as is more likely, there i only one p H meter fure\.ery ru,oC.r three reams, the valur of C ran br dtwrmind initially;~~ rlrm ttmpt.r;iture ( 6 " if room temperature is 25°C) and then a t two widely separated values of T (e.g. 75 and 10°C). Procedure The half-cells are constructed by pouring 25-40 ml of CuSOl or Ph(NO& solution into a short 50-ml test tube and placing the appropriate electrode into the test tube. A third 50-ml test tube is filled with 25-40 mlof an electolytie solution of mixed PbiNOa12andKNO:,. The test tubes are placed in a 600-ml beaker, the salt bridges are added, and the pH meter leads are attached to the electrodes with alligator clips. Because the S042-ions will migrate from the Cu2+ solution, through the salt bridge and into the Pb"+ solution, an intermediate electrolytic solution of 0.25 M Pb(NO:J~/0.50 M KNO:, is used. The salt bridges connect the electrolytic solution to hoth the Pb/Pb2+and Cu/Cu2+half-cells. When the half-cells are heated, the rate of ion migration increases, hut the presence of Ph" in the electrolyticsolution precipitates the majority of the Sot" ions as PbS04 and prevents them from reaching the Ph/Pb'+ half-cell. After the room temperature measurement of C is taken and the temperature is recorded, ice and water are added to the 600-ml beaker to cool the cell. When the temperature has stabilized (5-10 m i d , measurements of C: and Tare recorded, and the solution is warmed to -70°C. The heating is stopped, and temperature and voltage measurements are made again. The value of 6 at room temperature is used in eqn. (31 to determine a value for AG for the reaction. The values of ii at high and low temperatures are used in eqn. (5)to determine a value for AS. Materials The solutions were prepared with distilled water. The CuS04, Pb(NO&, and KNOa were reagent grade chemicals, while the Pb and Cu electrodes were, respectively, '/-in. sheets of 99% Dure metal and 12 .eauee " electrical wire. The electrodes were cut to about 6 in. lengths. Bacteriological grade agar was used in the salt bridges. Voltage measurements were made with a Beckman Zeromatic SS-3 p H meter, with the zero set a t midscale ( p H 7) and the selector switch set to f mV. The salt bridges were prepared by dissolving agar (2%by weight) and KN03 (10% by weight) in water, heating until the mixture thinned, thickened, then thinned again. The hot solution was poured into wet glass U-tubes, and set up in 10-15 min. Results The values of 6, AS, AC, and AH for the Ph/Ph2+// K+.Ph2+, NO?-/lCu2+/Cu cell are shown in the table. These experimental values provide excellent agreement with the literature values. The ex~erimentalresults are most dependent upon accurate of the solutions of Table of Results Experimental
AH
-14.8
Value
kcallmole
Accepted Value'
-15.00
kcalimole (8. 9)
Literaturevaluer are for 8 '. I S ' . AGD. A@, The experimental values wmld be for swmrd states if fkrmm temperature measurementisat 25'C sothat i .= Eo. If me r w m temperature measurement is made at a temperature other than 2S°C. the value of 6' can be determined by using the experimental valuer of I s a n d 6 at room temperature and scaling to get the C at 25'C.
Volume 55, Number 1. January 1978 i 53
CuS04 and Pb(NOa)2.If these solutions provide a value of 6 a t room temperature between 0.458 and 0.468 V, the cxperiment will give accurate results. An infrequent problem was that poor electrical contact was sometimes made between the alligator clip and the electrode. This caused the voltage to drift steadily in one direction. A more serious problem was that the electrodes must have fresh metal exposed before each use. Steel wool worked well on the Cu hut not on the Pb, where more vigorous methods of physically removing the surface oxide were needed. The use of Z n P b , CuIAg, and PbIAg cells was investigated and they provided agreement to within 10-30% of literature values. However, none of these other systems gave the consistent results that were observed with the PhICu system. Discussion
As stated before, this experiment provides the student with the opportunity to do experimental thermodynamics and to calculate those thermodynamic values which usually cannot he determined with the simple equipment available in a general chemistry laboratory. The agreement between experimental and literature values should reinforce the concept that chemical thermodynamics is an experimental as well as a theoretical science. If time, talent, and equipment permit, the addition of a Zn/Zn2+ half-cell to the 600-ml beaker, with a salt bridge into the electrolytic solution, will make it possible to follow simultaneously Z n P h , ZnICu, and PbICu batteries. The sign of Ad is different for the ZnICu and ZnIPb cells, compared to the PbICu system. If the same clip is left on the
54 1 Journal of Chemical Education
P b electrode, both voltages can be quickly measured by moving the other clip from Cu to Zn. The G for the ZniPb is small enough that it will remain on scale with the p H meter prepared as previously described. The change in voltage for the ZnICu battery is extremely small (-8 mV) and an expanded scale capability is suggested for good reproducibility. Acknowledgment
I wish to thank the other members of the Department of Chemistry, USMA, for their suggestions and assistance in preparing this manuscript and, in particular, COL Wilford J. Hoff, Jr., Major Charles E. Figgins, Captains John S. Polles and Daniel E. Adams, Dr. W. David Loehle, and Mr. Frederick S. Rose. Literature Cited i l l Cn~ekf)xd.HL a n d N~we1l.d.\V.,"Labnratury Manualul'PhysiealChemistrs."d~ihn Wiieysnd L,nr, New Ycirk. 1956. p. 167. Chemistrv." Hsrcourt Brace luvsniwich. 121 Jollu. Wm. L.. "Enmunters in Exoerimentai ' lit, New Ywk. 1972.p. 78. I:li "Chemistry LahrafryMsnua1,"Vol I. 1976-1977. Department nf Chemistry. United States Military Academy, West Pi,int. New York. 1976.p. 3-A-I. 141 '"PhpicalChemistry Lahmatury Manual." 197b1976. Ikpartmmt i,fChemirtrr. United Staler Militar?Aeadem?.Wert Point. New York, 1975.p. 3.1. . - P ~ ~ S ~ C ~ d~ .~. is) M , X , , ~ ,W ~ I ~ ~ ~ . J . chprnirtry,"ard ~ ~ ~ i ~Ciiir~. ~ , N?W , ~ > d Jersey, 1962.p. 8fi. (61 Sksv~.I). A,. and West, U. M.. "Analvticai Chemistry," HcAt Rinehart and Winslon. NewYork, I961.p.27. (71 Latimer. W. M.."Oxidation Patent~als."2ndEd.. Prmtice-Hall. Inc.. Englew,md Clifis. New .Jersey. 1952. pp. 31(1-:115.
