The FERRIC THIOCYANATE EQUILIBRIA* ROBERT DuBOIS Fresna State College, Fresno, California
F
OR many years, both in this country and abroad, lecturers on chemistry and textbook writers have illustrated their discussions of the important subjects of mass action and chemical equilibrium by showing the reversibility of the reaction between fenic chloride and ammonium thiocyanate in solution, usually represented by the equation The red color produced is attributed to the formation of undissociated, colored molecules of ferric thiocyanate, Fe(CNS)s. In the lecture demonstrations these two solutions are mixed, and the resulting dark red solution is poured into four separate beakers. The law of mass action is then illustrated by adding to one of the beakers an excessof ferric chloride, to the second an excess of ammonium thiocyanate, and to the third an excess of ammonium chloride, the fourth beaker being kept for reference to compare the color changes produced in the others. In the first two beakers (addition of ferric chloride and ammonium thiocyanate) the concentration of the colored product is increased because of the increased concentration of the reactants, and the red color deepens. In the third case the color fades somewhat as the ammonium chloride is added, and this effect is interpreted as a reversal of the reaction represented above. Now ferric chloride and ammonium thiocyanate are both strong electrolytes, and their solutions contain principally the ions Fern, C1-, NH4+, and CNS-. The reaction which occurs when these two solutions are mixed is therefore primarily an ionic reaction which would be more accurately represented by the equation
Considerable evidence has been provided in favor of the view that the redcolored product in the ferric thiocyanate reaction is a complex ion, presumably the hexathiocyanatoferriate ion, Fe(CNS)6-. The very convincing evidence offered by Schlesinger and Van Valkenburgh' was based on molecular weight determinations of ferric thiocyanate in non-aqueous solvents, on the similarity in the absorption spectra of fenic thiocya-
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* Presented before the Division of Chemical Education at the ninetieth meeting of the A. C. S., San Francisco, California, August 21, 1935. 1 H. I. SCALESINGER AND H. B. VAN VALKENBURGH, 3. Am. C h m . Soc., 53, 1212 (1931).
nate in water and in ether, and on the electronegative migrationof theredcomponent of aqueousferric thiocyanate solutions. In further support of the complex ion postulation may be mentioned (1) the large tendency of thiocyanate ion to coordinate with metallic ions and (2) the tendency of the transition elements with incomplete electron snbshells, and of iron in particular, to form coiirdination complexes, usually with a coordination number of six. Table 1gives a partial list of some of the known thiocyanate complexes.
We shall therefore represent the ferric thiocyanate reaction by the equation
There now arises the question, why should ammonium chloride have the effectof reversing the formation of ferric thiocyanate, which involves neither ammonium nor chloride ions? It is o&us that the explanation of the decolorizing action of this salt will have to be based on something other than the principle of mass action, which it is traditionally supposed to illustrate. THE
DECOLORIZING ACTION NOT DUE SOLUTION ADDED
TO ACIDITY OF
In looking for indirect ways in which ammonium chloride might d e c t the ferric thiocyanate equilibrium, the tendency of ammonium salts to hydrolyze suggests that the acidity of the ammonium chloride solution might, in part at least, account for the effect observed. This suggestion must be abandoned, however, for two reasons (1) the conductivity data show that thiocyanic acid is a strong electrolyte, and there is very little tendency for the added hydrogen ion to remove thiocyanate ion as undissociated HCNS; (2) as a matter of f a c t i ~ . M. Vernon,=de Brouckke and Gillet,s and others, have found that the first effectof adding acid to a solution of
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' H. M. VERNON, Chem. News. 66, 214 (1892). L.
DE
BROUCKBRE AND A. E. GILLET, Bull. sac. chim. Bdg.,
42, 281 (1933).
324
femc thiocyanate is to increase the color. The effect is due to conversion of thecolloidal ferric hydroxide present (formed by hydrolysis of the femc salt) into free femc ion available for reaction with thiocyanate ion
ions and a change in the activities of these ions. The effect of an added salt on the velocity of a given ionic reaction will be measured by the effect on the activity coefficients of the reacting ions. In the case of the ferric thiocyanate reaction
it will he noted that the interionic attraction effect of added salt will tend to alter the activity coefficient of the colored complex product as well as those of the femc and thiocyanate ions. It is not immediately obTHE DECOLORIZING ACTION NOT DUE TO SPECIFIC EFFECT vious that the salt effect should he a decolori~ingeffect, OP NX%+ OR CL- IONS that is, that i t should hinder the forward reaction more than the reverse. However, that this should he the That the decolorizing action of ammonium chloride direction of the effect can be shown in the following way. on ferric thiocyanate solution is not due to any specific We shall first introduce the activity coefficients,f,of effect of either ammonium ion or chloride ion can be the reactants into the expression for the equilibrium shown by comparing the effect of adding equally conconstant of the reaction: centrated solutions of a series of salts such as ammonium chloride, ammonium nitrate, sodium chloride, andsodiumnitrate. When the experiment is performed, one observes about the same decrease in color in all which rearranges to cases. THE DECOLORIZWG ACTION PRIMARILY DUE TO ADDITION OF CHARGED IONS (SALT EFFECT)
In the experiment just described it has become apparent that we can effect the apparent reversal of the femc thiocyanate reaction by adding anelectrolytewhich has no part a t all in the chemical reaction as usually written and which has, indeed, no ion in commonwith any of the reactants or products. We cannot, therefore, escape the conclusion that any salt (not causing precipitation) might have the same effect on the ferric thiocyanate equilibrium as does ammonium chloride. We find support for this prediction in the fad that many examples have already been found of the influence of neutral salts on reaction velocities and on chemical equilibria between substances with which the added salts do not react directly in a strict chemical sense. Of these examples may he mentioned (a) the large increase in the dissociation of weak acids in the presence of salts such as sodium chloride, (b) the increased solubility of slightly soluble salts brought about hy the addition of foreign electrolytes to the solvent, and (c) the effect of neutral salts on theinversionof sucrosecatalyzed by acids and on the catalyzed hydrolysis of esters. These salt effects have been discussed in detail by B r ~ n s t e d . The ~ effect on the velocity of a reaction hetween uncharged molecules is proportional to the concentration of the added electrolyte, and this has been called a "linear salt effect." In the case of a reaction involving one or more ions the principal salt effect is proportional to a higher power of the salt concentration than the first and this is, therefore, called an "exponential salt effect." This exponential effect (referred to hereafter simply as "salt effect") is due to a large increase in the ionic atmospheres surrounding the reacting
We see that the effect on the apparent dissociation of the colored Fe(CNS)s--- complex is measured by the change in the ratio of the activity coefficients. If the solutions are sufficiently dilute, we may represent the dependence of the activity coefficients on the salt concentration and the valence type of added salt by the use of the Debye and Hiickel limiting law, and the predictions thus made will he a t least qualitatively valid for more concentrated solutions. The relation just mentioned is log f j = -0.5 zip.\/;
or fj =
~O-~~""G
wheref, and z, are the activity coefficient and the valence of the ion i, and p, the ionic strength, is '/& c&. When the concentrations of the ionic reactants are small, as is usually the case in the femc thiocyanate experiments, the ionic strength of the solution is mainly determined by the concentration and valence type of the added salt. Substitution in the equilihrium expression gives
The salt effectis therefore fiositiue,a t least a t low concentrations. That is, with increasing salt concentration, the apparent dissociation of the colored complex is increased and the color of the solution decreases. We shall expect the salt effect to he (a)important a t low concentrations (because of the exponential nature of the function) and (b) more pronounced the higher the valence type. Both of these exoectations are ahundantlv fulfilled bv the results of experiments which have been made under
the author's direction in the chemical laboratory a t Stanford University. Mr. C. L. Flimdt has used a photoelectric calorimeter to measure the effects of increasing concentrations of various electrolytes on 0.00375 Msolutions of ferric thiocyanate. The relative magnitudes of the effects for equal molar concentrations of potassium chloride, potassium sulfate, magnesium chloride, aluminum chloride, and phosphoric acid were just in the order of the ionic strengths of the solutions. The decolorizing action of the higher valence type salts was pronounced a t 0.01 M concentration and increased rapidly. The effects of hydrochloric acid and sulfuric
acid were smaller than those of the corresponding potassium salts, and the difference in each case can he attributed to the effect of hydrogen ion in reversing the hydrolysis of the ferric ion. Experiments made in the same laboratory by Mr. John Lyman and Mr. Bryce L. Crawford show that chloride solutions have a slightly greater decolorizing action than the correspondmg nitrates, presumably because of the formation of the complex hexachloroferriate ion, FeCls---. Other compleu-formingions, such as oxalate ion, have even more pronounced decolorizing effect on ferric thiocyanate solutions.
