An Experiment on the Molar Solubility and Solubility Product of Barium Nitrate Bern Wrudc and
Reinstein University of Wisconsin-Platteviile, Platteville. WI 53818
There are todav manv methods of determinine the K, of a slightly soluble &mpo;nd. These previous K., ixperiGents tend to be so involved in their chemistw that the students tend to lose sight of the final objective: The authors have long felt a direct gravimetric method of determining a K,, value would be desirable. T o accomplish this a salt with a relatively large K., value would he required. A literature search yielded an experiment by Moeller and Martin1 that used lead(I1) chloride. There were a few major drawhacks t o the experiment, namely lengthy cooling periods and evaporation time, and thecommon ioneffect wasonly investigated aualitativelv. experiment is designed to be easily completed The in a two-hour lab oeriod while keenine - the number of chemlcal principles to a minimum. Barium nitrate, whose molar soluhilitv is a ~ ~ r o x i m a t e 0.3 l v a t room temperature, is the salt of ~ k o i c ei.t~has good settling properti&, which allows for the solid remaining after dissolution to be weighed and the decantate to he discarded. In dilute solutions the molaritv of ions is approximately that of the activity. In more concentrated soluiions the activities may vary substantially from the mola~ities.~ In either case, this difference is frequently ignored in general chemistry textbooks. The authors feel that, in an introductory experiment such as this, the use of molarities in the equilihrium expression rather than activities does not detract from a n understanding of soluhility and soluhility product. ~
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Equipment and Reagents Solid Ba(NOs)2,0.50 M HN03, 250-mL beaker, stirring rod, thermometer, Bunsen burner, ring stand, wire gauze, large container (to hold barium nitrate solutions and solid). Recouery of waste. The barium nitrate solid can he recovered by evaporation of the water and/or HN03. Any harium that may he disposed of down the drain should first be precipitated as the sulfate. Caution: Barium nitrate is toxic if ingested. Procedure Part I Weigh a clean, dry 250-mL beaker to the nearest centigram. Place about 5 g of B a ( N 0 h solid in the beaker. and weigh t o the nearest centigram. Add 50.0 mL of distilled water. Stir for 10 min. Measure
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Moeller, T.; Mertin. D. Laboratory Chemistry; Heam: Boston. 1965. Reinstein. J. J. Chem. Educ. 1984, 81, 1009. . Meites, L.: Pode. J. S. F.: Thomas, H. C. J. Chem. ~ d u cI966,43. 667-672.
Table 1. Sample Results of the IZ, Determlnatlon 01 Ba(NO,), Initial weight of Ba(NO& (g)
Weight of Ba(N0.). in saturated soh (g)
1. 5.02 2. 5.03
3. 4. 5. 6.
4.20
5.06 5.00 (commn ion)
4.16 4.16 2.38 2.39
5.04 (common ion) gravityfiltrallon mean i standard deviation (n = 4) 7. sU*ion IllIralion mean f standard deviation (n= 4) Table 2.
Molar SolubilHy (moi/L)of Ba(N0.h
I&
of Ba(NO& (21 OC)
0.322 0.319 0.319 0.182 0.183
0.134 0.130 0.130 0.136 0.137 0.319 f 0.003 0.130 f 0.005 0.328 + 0.005 0.142 f 0.007
Sample Calculation of the IZ, 01 Ba(NO& with Common ion Effect
5.00 g lnnial weighl ot BaiN0.h 2.62 g WeigM ol Ba(NO& remaining a W dissolution 2.38 g Weight of Ba(NO& in 50.0 mL0f sahxated soh 9.12X 10-3moi Males of Ba(NO&in 50.0 mL of saturated soh Molarity of Ba(NOd2 0.182M 0.182 M Molarity of Be2+ Molarity of NO< (common ion) 0.50 Z(0.182) = 0.864 M [Ba2+][NOs-]2= 0.136 K,, of Ba(NO&
+
the temperatureofthesolution, and thrndecant asmuch as possible of the saturated solution of barium nitrate into the specified waste container. (The settling ropert ties of RnINO.,)?are such that less than 1mL saturated &t& remains, introducingan error of less than 2% which is satisfactory for K, determinations.) Heat the beaker and contents carefully over a low burner flame. The flame should be W in. from the bottom of the beaker. Evaporate to dryness (about 5 min), being careful to avoid splattering. Let cool to room temperature, and then weigh the beaker and content9 to the nearest centigram.
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Part If, Common Ion Effect Repeat the above procedure, except substitute 50.0 mL of 0.50 M HN03 for distilled water, and evaporate the contents in the hood.
Data and Calculations From the weight of the B a ( N 0 h that dissolved, one can calculate the molar soluhility of harium nitrate. Using the 1:2 mole ratio of BaZt t o NO3-, the solubility product can then be calculated. Sample data is given in Table 1. In the common ion experiment, the student must add to the molarity of Nos- coming from the dissolution of the salt Volume 66
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June 1989
515
the value of 0.50 M. I t should be emphasized that the [Baz+] to [NOS-] ratio in solution is no longer 12. A sample calculation is given in Table 2. An exercise that may be assigned after completion of the experiment could be as follows: What would be the K, value of Ba(N0s)z if one mixed 50.0 mL of 0.50 M BaC12 with 50.0 mL of 3.0 M HN03 and recovered 3.82 g of Ba(N03)2?
A problem of this type is useful for several reasons: It shows the student need not start with Ba(N03)z to determine the solubility product constant. (2) It shows. once azain. that under these circumstances. the ratio of [na2-'jto (N&] ;ernaining in solution is not 1:2.'~hisfact is difficult for students to understand and should be emphasized. (1)
Discussion An examination of Table 1 shows the K , values were reproducible both with and without the com&n ion effect. The molar soluhility is shown co decrease significantly with the common ion effect. Variations on the experimend procedure were tried, such as filtering the moist Ba(NO& on preweighed filter paper using suction or gravity. The solid was transferred and washed with 95% ethyl alcohol. The samples were dried by placing the filter paper on top of two watch glasses and heating over a water bath. If done carefully, the K,, value remains reproducible. Larger K,, values were obtained when 0.50 M KNOa was used in place of 0.50 M HN03. This would not he due to a
516
Journal of Chemical Education
residue of potassium nitrate remaining on the Ba(NO& crystals as this would have the opposite effect on the K., in this experiment. In order to account for the greater soluKility of barium nitrate in potassium nitrate, the conceotsof ion hydration and ion-pair formation may be utilized.- he solvation of the potassium ion may be less than that of the hydrogen ion. This could effectively increase the amount of water available for the solution p r o c e ~ sThe . ~ "free" nitrate ionconcentration may be reduced due toagreater number of ion pairs being formed by the putassium ion.5 Hoth of these factors would lead to an increase in the molar solubilitv of ~barium nitrate. The short duration of each ex~erimentaltrial has sueeested a study of theeffect of a change in temperatureon ti;;; K,, value. The simplicity of the experiment has allowed instructions to he kept at a minimum, and more importantly, it has resulted in a much greater understanding- of soluhilitv product by the student.
