Solubility of Lead Sulfate in Water and in Sodium Sulfate Solutions An Experiment in Atomic Absorption Spectrophotometry Thomas A. Lehman Bethel College, No. Newton, KS 67117 and Wayne W. Everett Ouachita Baptist University, Arkadelphia, AR 71923 Measurement of Pb2+ concentrations hy atomic absorption spectruphotometry permits direct c~lculat~ion of the soluhility ~ r o d u c rcunsttlnt of lead sulfate in the presence of low conEentrations of added sulfate in an evp&iment that we have developed and used successt'ully for freshmen and advanced students. I t demonstrates the dependence of the solubility product on ionic strength in the presence of a slight excess of anion. Activity coefficients depart from unity when added ion concentrations reach about 0.001 M, ( l ) ,so that the solubility products increase quite markedly. Although atomic absorption has been popular for more than t,en "venrs - - - - (2). . ,, nuhlished exneriments have not usuallv dealt G t h chemical principles (3j;instead, they have been specific determinations (.4.. ) . or technical a d a ~ t a t i o n (5). s We wanted ~ - ~ an experiment that used atomic ahsorption spectrophotometry to illustrate a fundamental chemical concept.
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Students are told to put a small amount of PbS04into pure water and into solutions of NazSOl that range in concentration from lo-' to 2 X molar. Solutions are stirred vigorously, stirred again 20 min later, then set aside for 40 min while a lead calibration curve is prepared using solutions in the range of 1to 35 ppm Ph. Students usually find that their curve is not quite linear, which illustrates an important aspect of atomic absorption technique. As a precaution, the solutions are decanted and centrifuged before aspiration into the flame. The absorbance was read at 217.0 nm when using the Varian-Techtron Model 1200 atomic absorption spectrometer and at 283 nm on the UV setting when using the PerkinElmer 303 atomic ahsorption spectrometer. An air-acetylene flame was used in all eases. A copy of our instructions to students is available on request. Results and Discubslon Results obtained by eight general chemistry students appear in the figure. Error bars indicate f one standard deviation. By plotting (K,)'I2 as the ordinate and multiplying the scales along both axes hy the same power of 10, the students can see for the data in the firmre that the chance in square root of the apparent Ksp is lesi than one-tenth the vaiiation in added sulfate. This clear sense of the relative changes in the independent and dependent variables is hidden, a t least from freshman, if the ordinate is K.,. The data are treated as the square root of the K,, also, because of the lack of agreement of the data with anv activitv treatment. The m a t that can he said is that the silubilit; increases with increasing ionic strength. The data of Kolthoff and co-workers (6)on lead solubility in sodium sulfate solutions also cannot be explained solely on an activity basis. The point on the ordinate in the figure is a t 1.3 X 10-4 M, our calculated solubility of lead sulfate in pure water a t room
Variation of the square r o d 01 the solubilny prodvct constant wim Eoncematlon 01 added sulfate. Error bars indicate one standard deviation.
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temperature. (We do not use a temperature hath.) The agreement hetween the sauare of this value (1.7 X 10-8M2) &d the literature values (7) for the K., of lead sulfate is ex: cellent. Acknowledgment This experiment was developed in part during a workshop funded hv the National Science Foundation under grant No. HES 75-i4376 to the University of Arkansas f o r a mobile laboratory. We thank the NSF and Prof. T. D. Roberts, project director, for the opportunity to participate in the project. Literature Clted (1) The~oluhilityofPbSO~inMliouaaaiution.i~givsn by Links, Wm. F.,"Solubilitieaof Inorganic and M e t a l - W c Campunds," 4th Ed.. Vol. n, Arne,. Chem. Soe., 19ffi.
pp. 1320-1333. (2) Kahn. H. L.. J. CHBM. EDUC.,4J.A7 andA103 (1966). (3) An exmptim iepmvidd byDauiui, M. B.,uindLathhridpc.J. WWW JJC~.EDuC.,50, "n9,."m9>
CHBM. EDUC..53,165 11976). (5) Argauer, R. A,, and White C. E., J. C m . EDuC., 49.27 11972):Brut, J. C., J. C m . EoUc., 6% 396 (1973); Stock,M. G., Litfle, D. J..andDono-, R. J.. J. C m . EDUC., 61.51 (1974): Lieu, V. T., Cannon, *.,and Huddleston, W. E., J. CHBM.EDUC.,51, 752 (1971); Thistlewaite, P. J., and Trese, M.,J. CHBM. EDUC.,51. 687 11974); Haaking, J. W., Snell, N. B., and St-, B.T., J. CKBM.EDUC., 54,128 (1917). (6) Kolthoff. I. M.,Perlich, R.W., and Weiblcn, D . , J Phys. Chom.. 46,561 (1942). (7) T h e values found in eighttabla published in Ule h t d e d e range horn (1.3to 2) X loT8. can be caiev1at.d from data in NBS Tcehniesl Not. 270.3. A value of 1.34 x "selected Values of Chemical Thermodynamic Prowrtirs:' Weahingfon. D.C.
1968.
Volume 59
Number 9
September 1982
797
