A "Semimicro" Spectrophotometric Determination of the Kspof Silver Acetate at Various Temperatures John Lieberrnann, Jr.' and Ki J. Yun Thomas Jefferson High School for Science and Technology, Alexandria, VA 22312
We have had success in our Advanced Placement Chemistry course with a solubility study of silver acetate. Silver(1) concentrations in saturated solutions a t various temperatures are determined spectrophotometrically (I).The aualysis is based upon the quantitative reaction of Ag+ with equimolar mixtures of m-phenylenediamine and p-phenylenediamine a t p H 7. T h e reaction is illustrated below:
\
NH,
44'
+
3H'
+
E
-
s
N
NH? The product of this oxidative coupling of two aromatic amines has a molar absorptivity of 1.8 X 104Lmol-1 cm-1 a t 550 nm. T h e procedure is semimicro in chat a small volume (1 mL) of supernatant liquid is required for analysis. When choosing a salt for a solubility study, it is hest t o select one with a monovalent cation and anion, such a s silver acetate, for the following reasons: If one or both of the ions of the salt is divalent, the electrostatic forces between them may he so large a s t o produce appreciable ion-pair formation. T h e K,,then calculated from solubility data would not reflect the true ionic concentration in the saturated solution. Complicating side reactions such as hydrolysis would also be minimized. Finally, if we are to use concentrations in place of activities in equilibrium expressions, mean activity coefficients must beapproximntely equal toone. However, e\.en in dilutesolutions of salts with diwdent ions, the ionic strength may he such that this is not the case. Therrforr, if sulubilities are to be determined t h a t allow the calculation of K,,'s t h a t are reasonably close to t h e thermodynamic K,,'s given in handbooks and other reference works, salts like silver acetate should be employed. An excellent, in-depth discussion of these points has appeared in this Journal (2). Experimental Procedure Preparation of Saturated Solutions A stock saturated silver acetate solution is prepared by dissolving the solid in distilled water at 80 O C fsoluhilitv: . 2.5 el100 mL at 80 '0.(CautionY.)ort ti on soft his solution while still warm are olaeed in small irrew cap bvttles'. Thesp hotrles arc rhrn ri>ermostatedat Iempcraturps ronying from 5 OC trefn,o?ratur) ro 30 'C Iconatantowrnight. temperature baths, and allowed toequ~l~hrnrs ~
~.~
~~~~~
Preparation of Standard Curve A stock silver nitrate solution (1.00 X M) is prepared by dilution of a 0.100 M AgN03solution made with reagent-grade solid.
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Journal of Chemical Education
(Caution2.)Aliquots of thisstock solution (1.0 to 6.0 mL) are pipetted or measured by buret into 10-mL graduated cylinders (or volumetric~if available).To these aliquots are added by pipet 1.0 mL of pH 7 huffer (0.015 M NaH2P04 adjusted to pH 7 with 20% w/v NaOH), 1.0 mL of aqueous m-phenylenediamine (1giL) and 1.0 mL of aqueous p-phenylenediamine' (1 gL). ( C a u t i ~ n . ~These ) reagents should be added in the order given and the mixtures diluted to 10.0 mL with distilled water. Approximately 10 min after dilutionl absorbances are measured at 550 nm in 1-in. tubes with a Bausch and Lamb Speetronic 20 fitted with an adaptor for a 1-in. test tube cuvette. If a spectrophotometer with a long path length attachment is not available, aliquot volumes of stock solution and saturated solution can be douhled. The spectrophotometer is referenced with a blank made up in the same way as the standards. ~ are plotted against ~ ~ ~ from Agi concentrations calculated Absorbances dilution ratios to give a standard curve. Analysis of Saturated Solutions At least two samples from each saturated solution should he analyzed in the following way: Aliquots (1.0 mL) of saturated silver acetate solution at various temperatures are removed by pipet, taking care not to suck up any solid. These are transferred to 100-mL graduated cylinders (or volumetric flasks if availahle) containing distilled water (approximately 50 mL) which are then diluted to 100 mL with distilled water. Aliquots (1.0 mL) of this diluted saturated solution are pipetted into 10-mL graduated cylinders and treated as in the preparation of standards. Ahsorbances of the two or more samples from each saturated solution are averaged and compared to the itandud curve to determine the concentration of silver
' Author to whom corresoondence should be addressed.
-
Caution: Both silver sa'lts. ~ c-. and . ~mohenvlenediamine. ,~ ~.~and the solu1~~ns of these chemicals shoulo oe hano eo wh le wearing rJbber gloves. S lver n Irate is a corros ve solid and contact witn eyes and skin should be avoided. Solutions of silver salts can stain the skin. p phenylenediamine is a suspect carcinogen and produces skin sensitization. Both phenylenediamine and their solutions should be used in a hood. Bottiesshould be clear so that students can avoid sucking up solid when removing samples by pipet. However, the bottles should be wrapped in aluminum foil at other times to reduce photoreduction of Aat. = Since this solution became discolored in one to two hours, only enough was made for one class at a time. Fresh solution can be quickly prepared for a later class while students are working. T h e m phenylenediamine solution did not discolor rapidly enough to require this precaution. The blank solution also becomes discolored with time, but since the spectrophotometers are referenced with a fresh blank solution. this should not be a oroblem. preliminah havd~-~ shown that absorbances contin~~-~ ~,investiaatians ~Led10 ncrease rlowever. as long as approx mately tne same amount of time was al owed for co or development in 00th standaros and samples, reproducible results were obtained. Reproducibility was also obtained by timing from the addition of the pphenylenediamine solution. ~
~
~~
~
D~
~~
~~
~
~~~
~
~
~~
~
~~~
~~
~
~
Figure 2, van't Hofi plot for the K,'s
of silver acetate at various temperatures.
Reprerentatlve Student Data Temperature
[Ast]*
("C)
(M x lo2)
K,,
x
10'
*meaverage of Vlree determinations
Figure 1. Standard curve for the determination of the concentration of Ag+ in saturated solutions of silver acetate.
ion in the diluted saturated solution. The concentration of silver ion in the saturated solutions is then calculated from the dilution ratios. Discussion and Conclusion If students are sunnlied with a standard curve (Fia. 1)and . if there are sufficient rpectruphotometers lone for every two nair of students~, saturated solutiuns at three different trmperatures can be analyzed by each pair of students in a single 50-min lab period. A plan for coordinating the use of the spectrophotometers should be worked out ahead of time. This is especially true under these circumstances where the time at which an absorbance is measured is important. The typical student results in the table are an average of three determinations for each saturated solution. The K., values are calculated from the expression ~~~~~
~
.
It is assumed that the concentration of silver and acetate ion in the saturated solutions are equal (neglecting hydrolysis of acetate ion). Therefore, the K,,'s are calculated by squaring the silver ion concentration. The value obtained at 26 'C, 2.4 X 10-3. is in eood aereement with a literature (3) . . value of 2.3 X l0-3'(25 " Assumine that the chanees in e n t h a l ~ vand entropv for the dissolvrng of silver ace&e are independent of temperature over the range of temperatures employed, the van't Hoff equation is used to estimate the heat of solution of silver acetate. The form of the equation plotted is
'c).
data in the table is illustrated in Figure 2. The heat of solution calculated from the slope of the straight line obtained is 24 kJ mol-I. Analysis of the same saturated solutions by the Volhard method (4) for silver ion gave the same value for the heat of solution. Besides the eood results described above. there are other advantages t o t h i s particular experiment. The fact that it can be done in one period of average length and involves the use of a spectrophotometer is one of its chief advantages. In addition, laboratory exercises involvina silver salts can be cost prohibitive, but in this instance such small volumes of solutions are required that this should not be the case. Saturated iolutionsarr also easy to prepare as described and supersaturation of the stock solutiun ar it cools har never heen a prot,lem. The precipitation of rxcess silrcr acetate as well-formed needlelike cr).stals always catches the studenti' intereit. And,since somuch of thecourse time is spent in the discussion of various equilihria, the drtermination of an eauilibrium constant that agrees well u,ith a literature value isiery satisfying to studenis. If constant temperature baths are not available, the solubility could he determined at room temperature and a t refrigerator temperature. With solubility data a t two temperatures the heat of solution of silver acetate could still be estimated. The solubility of silver acetate in a series of sodium acetate solutions of varying concentration, but constant ionic strength, could also be determined. An experiment outlining how to use these solubilities to identify the species of silver in silver acetate solutions has appeared in this Journal (5). Llterature Clted V.
where K is the eouilibrium constant (K., in this case) at a particular temperature T (in degrees KeTvin) and AHb and ASo are the standard enthalpy and entropy changes, respectively, for the dissolving process. A plot constructed with the
1966.40.2053-2054.
1. Robrtus. R. L.; Levin, A n d Chem. 2. Meitel, L.;Pode,J. S. F.;Thomss. H.C. J . Chcm.Educ. 1966,43,667-672. 3. Margoiis. E. J. Chemical Principles in Colculationr o i rnnic Equilibria: Macmillsn: New York, 1966: p A. T a t b o o k ol Quontitolius Inorganic Anolvsis. 4th ed.; Longman: Esser,
1. voge1,
169.
England, 1978: BP 340-342. 5. Ramelto, R. W. J. Chem. Educ. 1966.13.299-302.
Volume 65 Number 8 August 1988
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