indicated that manganese up to 100 p.p.m, was retained on 25-em. columns when the HC1 concentration was 10M or higher. Copper and iron were retained on the column under the same conditions. The Mn(1I) was eluted by using 25 ml. of 6 M HCl. This concentration and volume of acid gave quantitative recovery of manganese from 0 to 100 p.p.m. After Mn(I1) was eluted, a 2.5111 HC1 solution was used t o elute Cu(I1). The volume used was 70 ml. These conditions provided for quantitative removal of Cu(I1) with no indication of Fe(II1) in the eluate. Trials with 2.00 and 2.25M HCl were unsuccessful. The concentration and volume of HCl for this separation are more critical than those used for Mn(I1) and Fe(II1). After removal of copper, Fe(II1) was eluted with 40 ml. of 0.5M HC1. This also was quantitative and successful in the range 0 to 100 p.p.m. In all cases a total of 25 ml. of solution of concentrations up t o 100 p.p.m. were placed on the column. Thus, the total maximum amount of any ion on the resin column was 2.5 mg. Effect of HCI on Flame Emission. Since HCl concentrations in the eluate were high (from 2.5 t o 6.OM), the effects of HC1 on the emission intensities of the lines used for determination of manganese, copper, and iron were determined. In general, a decrease in emission intensity was noted for all three elements a t acid concentrations above 2M. Because of
Table IV. Average Concentrations of Manganese, Copper, and Iron Separated by Anion Exchange Procedures
-4ctual Concn., P.P.M.
Exptl. Standard Av. Deviation P.P.M. P.p.m. yo ’ Mn
100.0 80.0 40.0 10.0
99.6 79.5 39.5 10.0
100.0 80.0 40.0 10.0
98.2 79.0 39.9 9.6
1.38 1.47 0.95 0.22
1.38 1.84 2.37 2.20
3.00 1.87 1.10 0.40
3.00 2.34 2.76 4.00
5.40 3.92 1.49 0.63
5.40 4.90 3.73 6.30
cu
Fe 100.0 ~. 80.0 40.0 10.0
100.5 78.9 39.1 10.2
this effect all samples were evaporated just to dryness and then taken up in known volumes of dilute HC1 prior to aspiration into the flame. The ion exchange separation of Mn(II), Cu(II), and Fe(II1) before their flame photometric determination appears satisfactory. Schrenk, Graber, and Johnson (6) have shown that a similar procedure eliminates the extraneous cations such as those of alkali and alkaline earth metals and the interfering anions such as phosphate and sulfate. Hunter and Coleman (3) combined anion exchange separation and colori-
metric methods of analysis to determine trace quantities of various polyvalent metal ions in quantities of plant tissue. In their work, smaller diameter columns and a much lower flow rate were used. The proposed procedure seems well adapted to routine determinations. One operator can watch several columns simultaneously. The columca can be regenerated and used many times. A gradual decrease in flow rate with reuse was encountered. Eventually columns were repacked because of the change in flow rate. LITERATURE CITED
(I) Dean, J. A,, Beverly, M. L., “Role of Organic Solvents in Flame Photometry,” Pittsburgh Conference on Analytical Chemistry and Applied Spectroscopy, March 1959. (2) Dean, J. A., Lady, J. H., ANAL. CHEM.27, 1533 (1955). (3) Hunter, A. H., Coleman, N. T.,
Soil Sci. 90 214 (1960). d. A., Moore, G. E., J. Am. Chem. SOC.75, 1460 (1953). (6) Menis, Oscar, Rains, T. C., ANAL. CEEM.32, 1837 (1960). (6) Schrenk, W. G., Graber, Kenton, Johnson, Russell, Ibid., 33, 106 (1961). W. G. SCHRENK RUSSELL JOHNSON
(4) Kraus,
Department of Chemistry Kansas State University Manhattan, Kan. CONTRIBUTION No. 615, Department of Chemistry, Manhattan, Kan. Work supported in part by Public Health Service Grant No. RG-5827. Pittsburgh Conference on Analytical Chemistry and Applied Spectroscopy, February 1961.
Nucleation in Precipitation Reactions from Homogeneous Solution SIR: In an article in ANALYTICAL CHEMISTRY(e), Fischer discusses the nucleation prozess as it occurs in precipitation from homogeneous solution. We wish to call attention to certain aspects of his paper which we believe require further discussion. We will consider Fischer’s statement, “It is shown by calculations from the rates of reactions that the initial stage of precipitation, during which all nucleation occurs, is not one of homogeneous precipitation but rather is one of direct mixing of reactant solution.” While we admit the possibility of the “direct mixing of reactant solutions,” this is by no means the usual situation. To arrive a t this all-inclusive statement, Fischer has limited himself to two reagents-Le., thioacetamide and sulfamic acid. We shall indicate that, even with these two reagents, the conclusion is not necessarily valid and, further,
we shall show that it is not the case with other reagents. Fischer uses the precipitation of cadmium with sulfide produced by the hydrolysis of thioacetamide to prove his point. We shall use the same system with the requirements that we use a
freshly prepared pure thioacetamide solution, thus a solution theoretically free of or containing a negligible amount of sulfide, and that precipitation cannot occur until the solution is supersaturated with respect to cadmium sulfide. While Fischer’s conclusions are based on the tacit assumption that precipitation begins as soon as the solubility product is reached, the fact is that the initiation of a precipitation process requires considerable supersaturation. The latter can be defined (6) as the critical supersaturation ratio, Sorit.= (I.P./Ksp)l’z, where I.P. is the ion product and K., the equilibrium solubility product.
Table I shows the results of our calculations, in which we indicate the time required for the onset of precipitation, under a variety of assumptions. It can be seen from Table I, if it is assumed that phase separation occurs as soon as the solubility product (3.6 X 10-29) is reached, that 0.2 second is required for initial precipitation. Fischer used the solubility product value a t 18’ C. although the reaction he describes occurred a t 25” C. When we employed the value given a t 25’ C., we calculated a value of 0.6 second. Further, when we assumed a range of realistic values for the critical supersaturation ratio, we calculated time values up to 510 seconds. The presence of an impurity--e.g., sulfide in thioacetamide or sulfate in sulfamic acid-would have to be taken into account in the above calculations. Here, the use of pure reagents is a VOL. 33, NO. 12, NOVEMBER 1961
1801
Table I.
Time Required for Initial Precipitation of
CdS
([Cd++]= 0.01M, [Thioacetamide] = O.lM, [H+] = 0.01.W) Sorit.
Assumed 1 1
5
[S-I
3.6 X lodzB(18’ C.) ( 4 ) 1 . 2 x lo-** (250 C.) (7)
a
3 . 6 x 10-27 1 . 2 x 10-26 3 . 0 X 10-26
14
1.1 X 10-23
216 510
1 . 2 x 10-24 4 . 8 x 10-24
10
20 30
Time,b Seconds 0.2
0.6 55
Calculated from equilibrium expression used by Fischer in Table I11 of (2).
* Calculated from rate equation given by Fischer in Table I11 of (8).
sine qua non and in many instances the use of a freshly prepared solution is just as important-e.g., the hydrolysis of thioacetamide could produce sufficient sulfide to cause immediate precipitation of a metallic ion. Although we have had no experience with thioacetamide, we have never noted any visible initial precipitation when recrystallized sulfamic acid is used as a reagent to generate sulfate. In addition, in our work with 8-acetoxyquinoline, a reagent which could cause a troublesome hydrolysis problem, we have not encountered any difficulty. In fact, light-scattering experiments by Takiyama and Gordon (9) of the precipitation of thorium 8hydroxyquinolatc from homogeneous solution with pure 8-acetoxyquinoline have not shown any initial turbidity which could be ascribed to the presence of 8-quinolinol. Furthermore, the conclusion that precipitation from homogeneous solution is initially one of “direct mixing of reactant solutions” is not applicable to many methods of i n situ generation of
There is still another example where the precipitant is completely escluded from the parent reagent(s), as in the
SIR: The note by Haberman and Gordon and my earlier paper to which they refer are in agreement as to the possibility of nucleation by direct mixing in common precipitations from homogeneous solution. We differ as to the probability of this possibility. With respect to Haberman and Gordon’s Table I and associated discussion, I fully agree that some supersaturation is necessary to initiate precipitation, as I stated elsewhere in my paper, but values of the critical supersaturation ratio as high as 30 may not be very realistic. Furthermore, use of solutions as fresh as called for in this table is seldom called for in published procedures and, in fact, may enhance nucleation for other reasons [Fischer, R. B., Anal. Chim. Acta 22, 508 (1960)I
It should further be pointed out that calculations such as those of Table I can do no more than indicate a very rough order of magnitude of the reaction time needed to initiate precipitation because of uncertainty as to the actual mechanisms of sulfide precipitations once the sulfide and hydrosulfide ions are available and as to initial form of the precipitate. A freshly formed precipitate is not necessarily the same, in solubility and in other respects, as is the digested, more stable form (unpublished data from this laboratory substantiate this statement). With respect to the claim by Haberman and Gordon that no precipitate is initially visible in many precipitations, I would point out that precipitation nuclei are not necessarily visible. A
1802
ANALYTICAL CHEMISTRY
reagent. For example, in the precipitation of metal hydroxides with ammonia generated by urea (S), the reaction medium obviously contains hydroside ion; here, proper adjustment of the initial pH prevents precipitation from occurring. Two similar examples in which the presence of some precipitant in the reaction medium is also without effect are generation of hydrogen ion by the hydrolysis of ethylene chlorohydrin (8) or 0-hydrosyethyl acetate (S), and the release of a cation from an EDTA complex (3). Other examples in which the exclusion of the precipitant from the parent reagent is not totally impossible, but where we have not observed initial precipitation, are the generation of chromate ion by the reaction of chromium(II1) and bromate ion (S), and generation of iodate by the reaction of 8hydroxyethyl acetate and periodate (3).
precipitation of nickel dimethylglyoximate from homogeneous solution (6), where the dimethylglyoxime is produced by the reaction of biacetyl and hydroxylamine. Finally, we must add a comment relative to Fischer’s suggestion about speeding up precipitation process. Although hastening the process may have no effect upon the ultimate particle size, there is another estremely important effect which Fischer does not mention. Unfortunately, speeding up the growth stage of a typical analytical method will result in greater contamination of the precipitate by diverse ions ( I ) . The consequence of this is obvious. LITERATURE CITED
(1) Cohen, -4.E., Gordon, L., Talanta
7, 195 (1961).
(2) Fischer, R. B., ANAL. CHEW32, 1127
(1960).
(3) Gordon, L., Salutsky, M. L., Willard,
H. H., “Precipitation from Homogeneous Solution,” Wiley, New York, 1959. ( 4 ) Hodgman, C. D., “Handbook of Chemistry and Physics,” 42nd ed., p. 1746, Chemical Rubber Publishing Co., Cleveland, Ohio, 1960. (5) La Mer, V. K., Dinegar, R. H., J . Am. Chem. SOC.73,380 (1951). (6) Salesin, E. D., Gordon, L., Talanta 5,81(1960). (7) Seidell, A,, “Solubilities of Inorganic and Metal Organic Compounds,” 3rd ed., supplement, p. 158, Van Nostrand Co., New York, 1951. (8) Shaver, K. J., Ph.D. dissertation, Syracuse University, 1952. (9) TaFyama, K., Gordon, L., unpublished data. NORTONHABERMAX
LOUISGORDON
Case Institute of Technology Cleveland, Ohio
great body of literature has been built up on the scattering of light by molecules and by other small particles. Nucleation and some subsequent growth could almost certainly occur in an aqueous medium prior to the first visible appearance of precipitate, even in very refined light-scattering measurements. I agree with Haberman and Gordon that nucleation in precipitations from homogeneous solution is not necessarily by direct mixing of reactant solutions. However, I feel that the evidence indicates that it frequently must be.
ROBERT B. FISCHER Department of Chemistry Indiana University Bloomington, Ind.
