ON THE SPECIES PRESENT IN AQUEOUS SOLUTIONS OF "SALTSff OF POLYVALENT METALS'
0
I.
History and Scope o f Discussion
LEWIS POKRAS Polytechnic Institute of Brooklyn, Brooklyn, New York factory. However, in many equally important cases, such thinking can lead only to erroneous conclusions the past 20 years it has hecome increasingly concerning the properties of solutions. Even in the evident that the species present in aqueous solutions of compounds of polyvalent metals with simple anions simple cases, precise physicochemical measurements he explained on the basis of equation (1). are much more complex than was formerly believed. This Paper will outline most recent thinking in this The concept of hydrolysis, or, in the BrBnsted terminalfield and present part of the evidence for the conclusions ogy, acidity of hydrated metal ions, is now include$ even in elementary texts,as is a discussion of complex drawn. It will also demonstrate that equations like ( 1 ) represent the exceptional case rather than the rule. ions. TO generalize, we will consider compounds2 of the Yet, these concepts are too often introduced as some where M is a metal in oxidation what special topics with the implication that they apply type: M~+"X~-'~CHZO, only to a. few special or rarely encountered situations. State +Y ( y > l ) and X is (usually) a simple inorganic Many students retain the impression that these proc- anion like nitrate, chloride, or sulfate, of charge -2. the &em- The c molecules of water per empirical formula unit of esses may usually be ignored in ical and physical properties of electrolyte solutions. the compound may, in the solid state, be bound in in water, several different ways which are not of interest in this ~h~~ visualize the dissolution of A~CI,.~H,O for example, to be described completely by the following discussion. The primary process which occurs on dissolution of equation: the solid compound in water is often represented by H,O A1Cls~6H,0(solid)F=+ IAl(H,O).l +a(aq) 3 ~ 1 - ( a q ) (1) equation (2) : H,O I n many cases such equations may be quite satisM.+uXs'.cH10 ===2 a[M(HIO)a] +u (aq) + bX--(aq) (2) Each metal atom becomes the center of a "relatively The first in a series of three papers on this topic to be lished in consecutive issues of T H E JOURNAL OF CHEMICAL EDUCA- free" solvated ion of charge y. The "complex ion" TION. is enclosed here by brackets, as are all other complex ions in this paper. The number of waters per ion, represented by the symbol d, cannot be assumed equal to the maximum INTRODUCTION
+
+
6.5
6.0
5.5
5.0
0.2 1.
P~ 0f an A
0.4
a6 08 10 I chloride~ sol on ~~ d d i t of i ~~vuiou. ~ ~ salts
Several of the reviewers of this paper have objected to the use of the term "salt" for the metal compounds of interest here. Their objections are well founded on several grounds; yet there is no alternative word or simple phrase which can be substituted. The phrase "metal compound" is too general since species such as borides, nitrides, carbides, ete., are of no interest here. Reactions between anions (such as carbonate) and water are of secondary interest, exoept where the products of suoh reactions may be involved in further reactions with the metal-ion species. "Electrolyte" is also too broad, since it includes HNOa, NaOH, etc., which are excluded from this paper. Conversely, the phrase "metal-ion" or aquo-ion is too narrow, since it is precisely the interactions between such ions, water, and common anions which constitute the subject. We Veh to inquire as to the various complex species which may form when a single pure solid substance (suoh as written generally in equation (2)) is dissolved in water. The properties of such solutions are not the summations of properties of simple cations and anions, but reflect interactions between the simple ions. We me, therefore, forced~ to use the word "salt" for this class of solutes ~ for want of a better, generally accepted term.
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VOLUME 33, NO. 4, APRIL, 19%
coordination number of the metal. Evidence from transference data, as compared with other sources, yields highly discordant values for the extent of solvation of such ions (1). The anions are also solvated. Little attention will be devoted to the processes of equation (2). The "secondary" reactions, which in many cases are more important than equation (2) in determining the physicochemical properties of solutions, are our subject.
Composition of Aluminum Sols Illustrated i n Figure 2 (Concentrations Are i n Milliequivalents Pep Liter) Anion i n Sol
I-
Br-
Cl-
Alt++ Anion
60.9 3.27
70.9 3.50
80.2 3.54
C2H30218.01 2.53
Professor A. W. Thomas of Columbia University. Professor Thomas and co-workers have explained the HISTORICAL BACKGROUND properties of many colloidal and "true" solutions in In the classical theories of Arrhenius and Ostwald on terms of ideas originally proposed by Werner and solutions of electrolytes, the salts as well as acids and Stiasny (2). The Thomas papers assume that the sols of Al, Be, bases were considered to be only partially ionized in aqueous solutions. The concept of "degree of dis- Th, Ti, Zr, Cr, and Fe are polymeric systems with the sociation" was applied t o the salts in exactly the same metal atoms bound by "diol" (dihydroxy) bridges. sense as it is now applied t o solutions of weak acids, Reactions leading to this polymerization are described - as 'Lolation"; these reactions are shown t o be affected like acetic. Primarily because so many of the "strong" electrolyte greatly by the particular anions present, which may also solutions did not obev mass law ex~ressious.Arrhenius- be involved in complex formation. Under appropriate Ostwald concepts ha;e been replaced by the interionic circumstances-generally those of high temperature, attraction theory of Debye and Hiickel for such cases. prolonged aging, and/or high pH-reactions described We now treat electrolyte solutions as though complete as "oxolation," which involve the essentially irreversible bonds, are assumed t o occur. dissociation occurs-in accordance with equations (1) formation of M-0-M Some of the evidence from the Thomas papers is or (2). Discrepancies in thermodynamic and other properties are explained primarily in terms of the effect abstracted below. All of the data are then interpreted of the ionic atmosphere-quantitatively expressed by in terms of the Thomas concepts. Figure 1 illustrates the effect of various anions added the ionic strength--on the activity coefficients of the ions in question. Specific salt effects must still be as potassium salts on the pH of a typical aluminum invoked to explain physicochemical data quantitatively. sol (3). The sol was prepared from AlCla solution by Very few inorganic substances are considered adding aqueous NH,, washing the precipitated hydrous exceptional. Only a few lead and mercury compounds, alumina, and then peptizing with minimum quantities of and several halides, were classically described as dilute HC1 a t 80'. The suspended matter was centridissociating by processes other than that of equation fuged off, and the clear sol was aged for 45 days. (2). However, several hundreds of papers in the Analysis showed (Al(II1)) = 44, (Cl-) = 1.22, and equivalents per liter. (Note recent chemical literature on complex formation make (NH,+) = 1. X it appear that the behavior of the alkali and heavier that braces ( ) are employed throughout this paper t o alkaline earth salts represent the exceptional rather denote concentration of the enclosed species.) This sol, like those of Table 1, contains a basic than the typical case. aluminum compound. In all these sols there was less PHASES OF THE PROBLEM than one equivalent of anion (other than hydroxide or For purposes of this discussion, the problem of aqueous species may be subdivided into five major phases? (1) hydrolysis of salts (2) polymerization of aquo bases, (3) formation of complexes in solution, (4) composition of solid phases a t equilibrium, and (5) the late of achievement of equilibrium. LFFECT OF K& ON VARlW5 Unfortunately, it is not always easy t o "unscramble" ALUMINUM HYDI(OYI0L 5OL5 the various a s ~ e c t sof the ~roblem. The difficulties DATA OF T Y O M A ~a?Al .--. . -~--. involved are illustrated by the following studies. ~~~
~~~
~~
A
WORK O F A. W. THOMAS
ALLTATL
To the best of this author's knowledge, the first American chemist t o interpret the data of aqueous chemistry consistently in the present manner was a The Introduction and Phase (1) constitute the present paper. -. . The second paper will consist oi Phase (2) of the'problem, and the third paper will contain the remaining material.
o
4.5
o
a1
0,2
0.3
0.4
0.5
NORMALITY AWLD K250+
ri,..
2.
mfect of KSO,
on the PH of SOB
vsriou. muminurnx y & ~ ~ i d .
JOURNAL OF CHEMICAL EDUCATION
Fi-
3. Conductance of Aluminum Hydroxide Sola eft*. Trratment with Various &lid Silver %Its
oxide) per equivalent of Al(II1). Such species are sometimes called, f or example, "aluminum oxychloride." The potassium salt solutions added (all were one normal) were initially a t pH values from 5.07 t o 7.00. Therefore, a slight decrease in pH (if any change) should have been anticipated on addition of salts. Actually, the acidity of this sol was little affected by dilution with water, but rose sharply as various anions
0
2
.4
.6
.8
1.0
NORMALITY ANION Fimm 4. EBwt of Anions on the pH of Maimurn Precipitation of "AI(OH)8"
were added. Oxalate was about ten times as effective as other anions studied, producing a definite pH change a t one-tenth the concentration of the other anions shown. Ammonium acetate was about as effective as oxalate. The pH with KBr was 0.114.27 pH unit less than for KC1 of equal concentration; with KI, 0.190.34 pH unit less than for I E 1 . The range is for increasing concentration of halide. Figure 2 illustrates the converse effect of addition of the same anion to sols originally possessing different anions (4). Four sols with compositions shown in Table 1 were all treated with K2S04and the pH rose with increasing (KzSO1]in every case. Figure 3 illustrates the effect of various anions on the conductance of a basic aluminum iodide sol (4). The studies were made by equilibrating 100 ml. of the stock sols with weighed amounts of solid silver salts (various anions) in a thermostat for several days and then measuring the conductance of the systems. Figure 4 summarizes results of a study (6) of the precipitation of aluminum by NaOH from A1C13 solutions containing added salts. The abscissas are the "equilibrium" pH values a t which maximum recovery of A1 (as ignited "A1203") was achieved. The ordinates are normalities of added anions. The term "equilibrium" is quoted since the pH's of the solutions were reported t o drift noticeably with time. I n one case Thomas reports that the p H changed from 6.18 to 6.81 in seven hours, increased to 6.82 in 0.05 pH two more days, and stayed constant within unit (glass electrode) for two weeks. Obviously equilibrium is achieved rather slowly in these systems. he "pH of maximum precipitation" rises with {anion) for F- and CzOr-, but drops with (lactate) and (SO Br- > Cl- > C2HaOz-. properties of sols other than those of chromium. To The reverse sequence is, therefore, the order of stability explain these effects, he proposed various processes which lead to the formation of polymeric metal-ion of the corresponding aluminum complexes. The conductance titration of Figure 3 illustrates the species. The following generalized interpretation has been extended (by the present author) somewhat same point. If the reaction were simply: beyond the concepts of Thomas. Ag+ W a d AgI C It is proposed that the hydroxides in the aquo-bases substituting the anion of the solid silver salt for conjugate to the aquo-ion acids (see equations (3) and iodide in solution,then the conductances should depend (4)) can act as "bridging" groups. The simplest on the ionic mobilities of the anions substituted for I-. possible reaction of this type is written as equation (9): The mobilities are Sod-- > I- > NOa- > CzHaOz- > tartrate. The conductance should rise as the more mobile SO,-- replaces I-; instead, it decreases. The SO,-- cannot merely replace I- in solution; it However, there is some evidence that dimeric species must enter the coordination sphere of the Al. involving only one hydroxide are relatively unstable The effect of salts on the pH of maximum precipita- with respect to dissociation into monomers. The simtion in Figure 4 is far too great for explanation by plest reaction leading t o species which are of sufficient medium or buffering effects. The retention of A1 in stability t o exist in appreciable concentrations is: solution by citrate and tartrate a t high pH is rather conclusive evidence'of complexing by these ions; the other cases differ only in degree. The different slopes of the curves are difficult t o explain, even qualitatively. This dimerization reaction yields an ion with the They undoubtedly reflect variations in such factors as proposed structure: (1) stability of various complexes, (2) effect of external acidity on the relative stability of the complexes, (3) solubility of the various species, and (4) relative rates of exchange between complexed and "free" ligands. Both the mass of the precipitates (Figure 5) and their Structure I sulfate content (Figure 6) indicate that SO4-- is rather tightly bound t o Al. The decrease in (SO,--) in the Dissociation of additional urotons bv the dimer precipitate with increasing p H indicates an equilibrium 4 Prof. M. Kilpatrick points out, in a private communication, involving OH-. that the rates probably depend on physical transfer and the conThe data of Figure 7 on the rate of solution of centration gradient, which in turn depend on the complexing prohydrous Al,OJ in acids are an excellent confirmation of posed here.
+
-
VOLUME 33, No. 4, APRIL, 1996
(Structure I) would provide additional hydroxides which could then form more bridges, as for example by equations (11) and (12):
Such reactions, which Thomas collectively calls "olation," lead to higher polymers whose properties would probably differ from those of the simple aquoions. The olation reactions might be expected to be reasonably rapid and reversible. Thomas therefore proposes the further reactions of "oxolation" to account for the slow return of aged sols to their properties prior to aging: H
\
--M(>\
/
Rate 1 \ /O\ / + Ma a -M\ /M+ 2H.Of Rate 2 / 0 \
(13)
H
For simplicity, coordinated groups other than those in the bridge are omitted from equation (13). Since there are some doubtful points in the "oxolation" argument, it is presented here without further comment.
MI. of 0.1001 M (in Mo) (NH&MorO14 solution Figure 8.
CONFIRMATORY EVIDENCE OF EYRING
C0nductomd.i~Titmtion of Cn(NO& by Ammonivm Puamolybdmts
The concepts proposed by Thomas have been conWhen fresh Cr(NOt)~or CrCI3 solutions are aged firmed in an interesting paper by H. T. Hall and H. under reflux, the authors report that, on addition of Eyring (7). The research cited is a rather novel ammonium paramolybdate, the Mo/Cr ratio drops approach to the study of hydrolytic equilibria in significantly. Their data are given in Table 2. solutions, in this case, of Cr(II1) salts. To explain the lower Mo/Cr ratios found on titration Figure 8 illustrates a typical conductometric titration of fresh Cr(N03)%by ammonium paramolybdate its of the aged solutions of Table 2, the authors propose carried out by these workers. They interpret the that OH- and 0-- if bound by Cr(II1) are not replaced by (MoOIH-), a t least in the time of the experiments. curves in terms of the reaction: For the polymeric species of Structures 11, 111, and IV, they predict ratios of Mo/Cr of 5.0, 4.0, and 14/3 respectively. +5
H Hall and Eyring cite the fact that R. D. Hall (8) in structure 11 [(H,O),CR-0 - - crw,oJ 1907 prepared a solid complex ammonium molybdoH chromate heutahydrate corresponding to the complex in equation (14): They then propose that this species form; in solution on replacement of each coordinated HtO by an (Mo0,H-) group. During formation of the complex, the N&NOa and HN03 formed are said to TABLE 2 produce an increase in conductance, due to the high mobility of H30+. However, past the equivalence Mo/Cr Ratios as a Measure of Aging of Chromium (111) Salt S o l u t i o ~o n Reflux point they suggest that excess (MoOlH-) reacts with H30+with a decrease in conductance." On this basis, four studies yielded an average value of 6.000 0.005 for Mo/Cr for fresh Cr(N03)3solutions.
-
*
6 L. C. W. Baker and co-workers (9-lo), in a recent study of the heteropolymolyhdates, show that this formula for the complex in equation (14) cannot he correct. However, the exact formula of the complex species is not of extreme importance; the cogent point is the Mo/Cr ratio. In this respect Baker reports (9, p. 2137) the preparation of solid samples (by the Hall method) with Ma: Cr ratios of 6:0.98.
158
JOURNAL OF CHEMICAL EDUCATION
-pH OF HYDROLYZING 'ALT 5OLUTION5 THL
DATA OF ZLTTLRWAU
LOB
r i p r e 9.
I
c (WMALITY)
The pH of Hy&olg=in. Sllt Solutions
I
THF. pH OF HYDROLYZIUG CaClr UPPER CURM-AMLRlCW LOMR CVPVE -GERMAN
DATA OF ZLTlERWI\LL
1
7
k
-2
-3
-1
0
LOG C (NORMALITY)
Fi-
-.
10.
The pH of C&I1 S o l u t i o ~
-
~
Structure I V
Table 2 demonstrates that the initial "olation" or formation of hydroxyl links is quite rapid, as evidenced by the rapid decrease in Mo/Cr ratio. The further, slow decrease in Mo/Cr ratio on aging is consistent either with the "oxolation" concept of Thomas, or possibly with slow loss of H N 0 3 or HC1 on reflux, resulting in additional "olation." The lower values for C1- relative t o NOa- indicate chloro-complexing, which was checked by an independent technique t o yield the data of the right-hand column of Table 2. HYDROLYSIS OF SALTS
Salt hydrolysis has been studied for a t least the past 60 years, with the results of various investigators often in serious disagreement. The type of data reported is exemplified by Figure 9, taken from the data of ZetterpH Value. for AlCL Solutions Normalilv AlClr =
0.1&5
0.06M
Data of Zetterwall(l1) Data of Thomas and Whitehead (3)
3.26 4.14
3.44 4.36
wall (11) on the pH of salt solutions. The figure illustrates the dependence of pH on salt concentration, and in the case of aluminum, on the identity of the anion. The data of two different investigators on solutions of AlCl, are compared in Table 3. These differences, amounting to 0.9 pH unit, are not unusual. However, they may not be due to experimental error but may instead reflect actual differences in the stoichiometry of the materials studied. Figure 10, also derived from the Zetterwall data, compares the pH of two samples of reagent grade CaCl*. The lower curve is for Kahlbahn material, from a European manufacturer whose reagents are regarded as excellent. The upper curve is for a sample, prepared by an American reagent chemical supplier, which meets all A. C. S. specifications. The differences in pH can be due only to differences in Cl/Ca ratios in the solid samples studied; in many cases the formulas on the labels of reagent chemicals can only he idealizations. They cannot be interpreted as giving the exuct stoichiometry of the materials enclosed in the package. For this reason, as well as others, little reference to the older literature on "hydrolysis" is made in this paper. THE BR~NSTED-LOWRYAPPROACH TO HYDROLYSIS I n the middle 1920's Lowry and Bronsted independently arrived a t the concepts of aquo-ion acidity which have now largely replaced the "hydrolysis" concept. It is instructive t o consider part of the evidence presented by Bronsted. Figure 11 on the hydrogen-ion concentration (determined by a kinetic method) versus concentration of Cr(C10& is typical of the data of Bronsted and King (13). The observed relationships could be explained by (1) the effect of salt concentration on K , of water, (2) the presence in the Cr(C10& of free HClO. as an impurity, or (3) accepting the behavior of the aquo-chromium ion as typically that of a weak acid. Applying the data of Harned and Owen (13) on the effectof electrolytes on K, and assuming the Cr(C10& to be inert as far as acid-base properties are concerned, [H30+)a t 15°C. should increase from 6.95 to 8.30 X 10-8 as ( [Cr(H20)o](C104)3] increases over the range studied. Actually, ( H 3 0 + )varies from 1.00 to 8.20 X over this range. Not only is the order of magnitude of ( H 3 0 + ) in disagreement by a factor of lo4, but also A (H30+I is about five times that predicted for an exclusively K , effect. Thus, possibility (1) may be eliminated from further consideration. The ratio (HC104]/[Cr(C10~)a)should be fixed if the acidity is due merely t o an impurity. Again assuming Cr(II1) to be inert, the {H30+)ie given by the two dashed curves of the figure. The upper dashed curve is derived from the datum a t lowest (Cr+3J,the lower curve from the datum a t highest (Cr+8]. Neither fits the observed relationship between (HaO+) and (Cr+81within experimental error. The chromium curve is therefore compared with the
153
VOLUME 33, NO. 4, APRIL, 1956
acetic acid curve, derived from the data of Shedlovsky and MacInnes (14). From the similarity of the two curves, as well as other data, one must interpret the behavior of the aquo-chromium ion as being that of a typical weak acid. Data obtained bv Bronsted and Volovartz (15) ~, for several aquo-ion acids are given in Table 4. Values given are the thermodynamic dissociation constants, p = 0, a t 15"C., corresponding to the equilibrium of equation (3). An mieht he ex~ected.for a eiven metal ion K.increases as the number of coordinated molecules of water increases. This trend is clearly exhibited by the four cobalt aquammines. The magnitude of these and the other dissociation constants indicates that these acids are frequently at least as strong as acetic. In fact, the aquoferric ion is a stronger acid than H F and is almost as strong as phos(K, = 7.2 X phoric acid ( K , = 7.5 X lo-$) in water. ~
~~
~~
~
~
-
HOAc
(H&)'-] ----- [C.IMPURITY 0
-
INFLUENCE OF THE ANION ON AQUO-ION ACIDITY
4 s the data of Thomas indicate, the quantitative study of aquo-ion acidity (hydrolysis) may be greatly complicated by formation of complex species with anions. The magnitude of the effect of complexing may be illustrated with data for the rare-earth ions which, as recently as 1930, were said to form no complexes. Figures 12 and 13 present part of the data of Moeller and Kremers (16) on the pH titration of rare-earth nitrate and sulfate solutions, respectively. All solutions mere 0.1 molar in rare-earth(II1) ion, and all were titrated with 0.1 normal NaOH a t 25 0.5'C. The marked difference in slopes of each rare-earth nitrate curve as compared with the corresponding sulfate curve indicates considerable interaction between cation and anion. The nature and extent of the interactions are not clear from these data; however, relevant information is found in the literature. Spedding and co-workers (17) have studied the conductance of ten different rare-earth sulfates and found that for nine, the ''instability constants" for the process of equation (15) range from 2.2 to 2.6 X 10W4.
*
I
-
I
2
0
4 10'
Fig"-
11.
6 [ACID]
8
10
I2
I
(MOLAR)
Th. Acidity of th. Aquockolnium Ion
for HS0,-equilibria at the ionic strengths they studied, one may derive: [NpSOd ++
Npi'
+ SO,--;
K'w = 3.14 X
(18)
Apparently, the instability constants for both lanthanon and actinon rare-earth sulfates are of the order of low4. If this value is applied to the sulfate solutions titrated by Moeller and Kremers, then 99.4 mole per cent of total rare earths were present in those solutions as snlfato-complexes, a t least a t the beginning of the titrations. It is not surprising that the sulfate titration curves differed from the nitrate curves, for entirely different chemical species were being titrated. I n the nitrate case, the rare-earth species were probably essentially simple aquo-ions (since other evidence indicates that nitrate possesses relatively little tendency to form complexes). Similar differences are found on extending the comparison to other salts. Table 6 presents data on the pH a t incidence of precipitation, and also when 0.4 moles OH-/mole rare earth have been added to various TABLE 4 Thermodynamic Dissociation Constants a t 15'C. Aauo-Ion Acids
Acid
for
K,
These low When R is yttrium, K I,,,,,,., = 3.4 X values of K,.,.,., correspond to rather stable complexes. Sullivan and Hindman (18) have studied the sulfate complexes of neptunium and reported that the equilibria of equations (16) and (17) are important in this TABLE 5 system. They also report values of "formation constants" (for the stepwise association of Np+' and Values of Formation Constants of Neptunium Sulfate Complexes a t Various Temperatures HS04-) reproduced in Table 5. NpC4f HSOI-
= [NpSO.] + H + KL ++
(16)
From the K values given by Sullivan and Hindman
JOURNAL OF CHEMICAL EDUCATION -
1,
NITRATE SOLUTIONS 10
DATA OF MOELLeR
.*
salts. The data on perchlorates are by Moeller and Fogel (19); a11 other are from (16). The rather regular differences between the several salts of the same rare earth ions are striking. However, it is unwise to attempt to elucidate too much from data of this type. Even though they may reflect the utmost care in obtainine the data., thev Dose an nnresolvable problem common in the current chemical literature: they offer no evidence as to the initial stoichiometry of the salts employed for the studies. While it is true that the acidity constants cannot be very big for these aquo-ions, this fact offers no real assurance as to the degree of hydrolysis of the solid sample employed in the
-
"
&
DH Values
study. The aquocalcium ion studied by Zettenvall (11) is probably also a weak acid, yet note the variation in pH found for several samples. ~ ~ U E ~ ~ s ~ ~ ~ M ~ ~ ANIONS , " , " IN y "STUDIES ~ T E OF
To avoid difficulties of the kind cited above. it has become common practice in studying the aqueous chemistry of metal ions to employ the perchlorate of the metal. As far as can be ascertained from the literature, the C10,- ion does not form complexes. However, Sutton (20) has recently proposed that C10,-, too, forms complexes in some systems.
During Precipitation of Rare Earth "Hydroxides" pH Values
7 -
At incidence of precipitation
-
At 0.4 OH-/rare earth
