The Solubility of NiSOs-6H20 A General Chemistry Experiment Richard A. Pacer Indiana University-Purdue University, Fort Wayne Campus, Fort Wayne, IN 46805 One of the major topics covered in the typical general chemistry course is that of solutions and solubility. Yet there is often little emphasis on the measurement of solubility in the accompanying laboratory. Often, laboratory work in this area is deferred until the student takes a course in physical chemistry or analytical chemistry. Thus, an opportunity to reinforce in the lab those solubility principles introduced in the general chemistry lecture is lost. A survey of articles on solubility measurement appearing reveals relatively few which are appropriate in THE JOURNAL for a non-majors type of introductory general chemistry course. Experiments suitable for the physical chemistry laboratory include a study of the many different equilibria influencing the solubility of zinc oxalate (1) and an experimental determination of heats of solution based on solubility measurements (2). Experiments designed for the analytical laboratory, such as solubility measurements of calcium and strontium sulfates (3)or that of lead bromide in nitrate media ( 4 ) , emphasize such factors as ion pair formation, complex ion formation, and ionic strength on solubility. As Reynolds (5) has pointed out, however, most high school texts limit their discussion of solubility product experiments to silver acetate, the cost of which is likely to be prohibitive. Reynolds (5)goes on to describe a suitable solubility experiment using barium hydroxide. It is useful to emphasize in equilibrium studies, such as the measurement of solubility, the fact that the position of equilibrium may be approached from either side. That is, the position of equilibrium (a saturated solution) may be reached from the unsaturated side (noted below in Figure 1as Method I) or from the supersaturated side (noted below as Method 11). In Method I, an excess of solute is added to distilled water to produce a saturated solution a t the temperature in question, and the system is allowed to come to equilibrium (6).Novel methods are sometimes used to stir the solution. For example, Schmitt and Grove (7) suspended a polyethylene bottle from a wire into a constant temperature bath and rotated it by attaching the wire to a variable speed motor. Stirring was accomplished by means of stationary magnets, both in and below the bottle. Alternatively, water may he allowed to flow through the solute via a saturator column device, such as the systems described by Butter (8) or Pacer (9). When equilihrium is established, solubility measurements are made by a suitable analytical technique, such as gravimetry (a),acid-base titration ( 6 , 5 ) ,complexation (EDTA) titration ( 4 ) , UVandIor visible spectrophotometry (2,9), atomic absorption spectrophotometry (3), or radioactivity counting ( I ) . These measurements correspond to Point A in Figure l. Usually the student is told to assume that, after a certain period of time or a certain number of passes through a saturator column, he or she is indeed a t Point A. In Method 11, an unsaturated but moderately concentrated solution is prepared a t a temperature well above room temperature (10,lI). The solution is allowed to cool. When Point B on the solubility curve is reached, the solution is saturated with respect to that particular solute. Actually, in practice, the solubility curve is crossed, producing a supersaturated solution, with the "degree of crossing" related to the induction
10
20
40
30
50
Tw.PEII*IIIIIE.S
Figure 1. Approach to solubility equilibrium horn an unsaturated solution (I) and a solution (11) that becomes supersaturated.
0.0
1
,
550
100
410 WAVELENOIH
Figure 2. Visible absorption spectrum of
400
In4
NiS016H~0.
time associated with the crystallization process for that solute. Such a solution is unstable (or, a t best, metastable). The temperature of the onset of crvstallization is dulv noted. From thf,weixht ofsolute and the ieight (riwater, thisolubility (at the temperature in question) can he calculated readily. At this point, a weighed amount of additional solute and/or water can br added, the mixture stirred and warmed to dissolve all solute, and the cooling process repeated. After an appropriate number of points (about five of six) are obtained in this manner, the solubility versus temperature curve may be drawn. Often, this technique will give excellent results. Wolthuis e t al. (IO), for example, obtained solubility results that were very close to literature values for KBr, KC103, KN03, and KzCrz07. However, three other salts (KzSO4, NH&l, and AlPOJ gave erratic results, which they surmised to be due to either supersaturation or hydrate formation. In our lab, students make solubility measurements using both Methods I and 11. The solubility of KN03 as a function of temperature is found using Method 11, following a procedure similar to that described by Wolthuis e t al. (10). The solubility of NiS04.6Hz0 is measured using Method I, following the procedure described below. The novel feature of this procedure is that the students make measurements along Volume 61 Number 5
May 1984
467
Table 1.
the way, 90 that 1ht.y can folluu the approach to vquilibrrurn as wt4l a$ obtain data at eqwlibrium. 'I'hui, they can nmclude from their own data that tht.system has reached npilihrium. Fdluwing the experinwnt, studmts areasked tocontrast and comonrr the t u ~m , e t h d i by..~uintinzout - thea~lvantagesand disahvantages of each.
Student Data for Solubility of NISOa'6H20
Temperature ('C)
NiS04.6H20Solubility (9150.0 g H.0)
21
34.5
22
31.3
Procedure
All ohotometric measurements are made with a Bausch and Lomb Spect;onie 20 spectrophotometer.Students are advised to select fairly large crystals of NiSOcGH20 (-10 g) and to dissolve them in an equal mass of Hz0. After about 10 min of stirring, solution is transferred to a cuvette, and the absorbance at 543 nm is recorded. The solution is then returned to the vessel and the stirring-removal-measurementprocess rc~eateduntil three successive absorbance values agree to i O . O 1 unit. A set of standards is prepared by weighing (to *I%) 0.5 g, 1.0 g, 1.5 g, and 2.0g NiSOr6H20 and dissolving each in 50.0 ml HzO.2.00 ml of the saturated nickel sulfate solution is added to 48.01111HzO,and the absarhanee of this solution and that of the four standards is read at 395 nm. An appropriate standard curve is drawn, from which the concentration of NiSOcGHnO in the saturated solution may be determined.
Table 2.
Calculated Values for Solublllty of NISOoWi20 (Based on Data in Ref. (121)
Temperature 1%)
NiSOc6H20 Solubility (gl50.0 g H20)
0 5
23.5 25.6 27.8 30.1
10 15
Results
The results in Table 1were obtained hv students in a first semester general chemistry class. Over haif of the students in the class had not taken chemistrv. . ~reviouslvin either hizh school or college. These results may be compared with calculated values based on those published in Seidell (12). The data for NiSOa6H?O are reported in Seidell in terms of g NiS04/100 g H ~ O ,f r i m whi'h one may readily calculate g NiSOc 6Hz0/50.0 g H z 0 (see Tahle 2). Discussion
The students' results are seen to be in fairly good agreement with the published values. However, the instructor should emphasize several points to the class: (1)avoid use of powdered NiS046H20 because of its long settling time after stirring; (2) fill cuvettes to w3/4 full; and (3) stopper, but do not rinse, cuvettes after returning solution to the vessel. If stirring is done with a magnetic stirrer, placing the beaker on a double-watch-glass setup will minimize transfer of heat from the maenetic stirrer to the solution. T h e vis~hleiibwrption spcrlrum nf nickel(l1) sulfatr shows a well-defined maximum at 395 nm. which obeys Beer's Law over the concentration range in q"estion. his is the best wavelength to use for analvtical Dumoses. However, a saturated soiution of NiS04 absorhitod strongly a t this waveleneth. Conseauentlv. a wavelen&h of considerably less sensitiiity corresbondini to a shoulder in the spectrum, 543 nm, is used to follow the approach to equilibrium. In selecting NiS046HzO for this experiment, the following criteria were evaluated 1) the species must absorb in the visible portion of the spectrum with at least one well-defined maximum between 350 and 650 nm; 2) there should be no serious toxicity problems; 3) the species must be at least moderately soluble so its dissolution can be followed visually; 4) the species should be readily availahle commercially; 5) ~. the soecies . should be relativelv inexnensive; 6) there should he no problems with hydrolysis reactions, decomposition, deliquescence, efflorescence, or air oxidation. Other substances considered included FeS04.7Hz0, KaFe(CN)~, MnClz, MnS04Hz0, MnC1~4Hz0,Cu(CzH30z)z. CoCl2.6H20, and VzOb. None of these substances was considered as useful as NiS04.6Hz0, since they did not meet a t least one of the above criteria. A word about cost is in order. Although theinitial cost may be higher than desirable (-606 per student), NiS044H20 can be effectively recovered if large, wide-mouthed containers are ~
468
Journal of Chemical Education
made available t o the students for collection of the crystals and solutions (assuming that the distilled water used in the experiment is checkedand found to be of good quality). Assuming losses in the 10-20% range, the recurring cost would be only -106 per student. The NiSO&HzO solubility experiment as described takes about 2112 h, including a half-hour for an introductory lab I-. ~.r t~~m eThe . KN0- e x ~ e r i m e ncan t be done easilv in 2 h. If both methods are 6 be'studied, two lab lkriods woad provide a eenerous amount of time for a lablecture and would allow most, if not all, of the learning objectives described below to be met. The fact that students with no previous chemistry were able to obtain good results in the NiS04.6H20 solubility experiment suggests that it may he feasible as a high school laboratory experiment as well. If the amount of time available is inadequate, the standard solution could be provided for the students. [Note: A lab "handout" will be made available to those who request it.] ~
~
~~-~ ~
Learning Objectives
One might summarize the nature and scope of this experiment by listing the major learning objectives incorporated either directly or indirectly into it. Of course, several of these will require some degree of elaboration by the instructor, either as part of the laboratory introduction or in a n accompanying lecture period. These include (1) acquiring a greater understanding of the terms "unsaturated,"
"saturated," and "supersaturated;"
(2) learning the meaning of the expression "to approach equilib-
rium;" (3) experience using a magnetic stirrer; (4) learning how to use a pipet; (5) learning how to make simple dilution calculations; (6) acquiring experience with speetrophotometry as an analytical
tool; (7) acquiring an awareness that a cdlored solution absorbs in the
visible portion of the spectrum and that the extent of such absorption depends markedly on wavelength; (8) learning how to plot data and draw the "best straight line" througha set of points (foradvanced classes,onecould use this
as an opportunity to introduce the linear least squares method of selectine a "best straieht line"):.. (9)learning what a "standard curve" is; (10) acquiring a feeling for the uncertainty associated with a given measuremen$ and (11) learning that the rate of dissolving doesnot proceed linearly with time. (This lends itself to a discussion of such factors as surface area of solute, rate of stirring, and rate of diffusion on the overall rate of dissolution. For advanced classes, one might want to discuss the relationship described by Gleeson (13), dw -= KS(C. dB
- CB),
where W = weight of undissolved solute 8. = time S = surface area of undissolwd sdutr C. = ronrmrration oisolute ln iolution at saturatilm C,, = conrtntrarion of 1h1P in s~lulionat tlm? i t K = a constant, As a special project, advanced students might he asked to design an experiment to evaluate K for NiSOa4HnO at 25%) (12) learning that, while solubility is indeed dependent on temperature, the nature of such dependence cannot he predicted. (In ~~~~~
this context, one might like to point out the observation made by Bodner (14) that over 94% of the salts for which data are tabulated in the "Handbook of Chemistry and Physics" undergo an increase in solubility with increasing temperature. O v e r two-thirds of the exceptions involve salts of the oxyanions SO4=-, P O P - . However, those who wish SeO?, S 0 s 2 - , A s O ~ ~ -and , to correlate the soluhility dependence on temperature with AH(solution) should duly note the caution discussed by Bodner (141.)
Literature Cited 11) Lyndrup. M. h. Robinson, E A., and Spmeer, J. N., J. CHEM. EDUC. 49, 641
-..
11977, , -,.
(2) Cesaro, A, and Rurra, E., J. CHeM. EDuc., 65. 133 (1978). (3) Dyrssen, D., lvanavs,E., and Aren, K., J. CHEM.EDVC.,46,252 (19691. (4) Cwper, J. N., J. CHEM.EDUC.,49,282 (19721. (5) Reynolds,J. P.. J. CHEM. EDUC., 52,521 (1975). (6) Eddy, R. D.. J. CHEMEDUC..35,364 (19581. (7) SchmitGR. H.,sndGrove,E.L.,J.CHeM.EDUc.,31,150(19M)). (8) But*, S.A..J. CWEME ~ ~ c . , 5 1 , 7 0 ( 1 9 7 4 ) . ~ ,(19711. (9) Pacer, R A . . J. CHEME D U C . , ~225 P., J.CHeM. Eouc.,31,137 (10) Wolthuis,E.,Pruiksma,A.B.,andHeerema,R. (111 Pekr8on.B. H..J CHEM.EDUC.,34.612 (1957). (12) Seidell,Atherton,"SolubilitiesofInargsnieandMetalOlganicCompoundr: American Chemical Society, Wsshingtan. DC, 1965, Vol. 11, p. 1219. (131 Gieeson,G. W., J.CHEM.EDUC.,16,187 11938). ~ ~ ,(19801. (14) Bodner,G.M.,J. CHEM. E D U C . , 117
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