July 20, 1958 DISPLACEMENT RE.wrxon-s AT TIIE SULFUR .ITOM 3565

sion constant of the polymer. For a typical value of R of 0.1 cm. and taking tD = r , we find for the Dow PSt. in tetralin with = 1.5 X lo6 a value of D of 4 x lo-’ cm.2/sec. This compares, in order of magnitude a t least, with a diffusion con- stant of 1.06 X lo-’ ~m.~,’sec,reported~~ for a PSt. fraction in toluene with a ilIw,r7 of 1.1 X 106.24 Such a diffusion process is caused not by the pres- ence of an initial concentration gradient but by a

(23) H A Stuart , ref 21, p 474 (24) The M dependence of I (since Y a 1/D) 15 not inconsistent with

t tie above.

gradient in chemical potential of the polymer in the drop.

Acknowledgments.-’e are grateful to Dr. E. Fat t and his associates a t the California Re- search Center for allowing us to study their pendant drop apparatus. Finally we wish to thank Prof. K. J. bfysels for many valuable suggestions and clarifying discussions in connection with the ex- perimental work as well as for his interest in this problem. LOS ANGELBS, CAI..

[CONTRIBUTION FROM THE CONVERSE MEMORIAL LABORATORY O F HARVARD USIVERSITY]

Displacement Reactions at the Sulfur Atom. I. An Interpretation of the Decomposition of Acidified Thiosulfate

BY ROBERT EARL DAVIS’ RECEIVED JAR’UARY 30, 1958

The kinetic results of La Mer and co-workers on the decomposition of acidified sodium thiosulfate in dilute aqueous The initial rate of production of sulfur is given by the rate k(Sa&-

This rate expression has been explained using a series of nucleophilic displacement reactions at the sulfur The production of various poly-

solution have been interpreted on a mechanistic basis. 08)8/z(HCl)1/z. atom. thionic acids and hydrogen sulfide during the decomposition has been explained.

The proposed reactions fit the experimental data and the rate laws are derived.

Introduction The decomposition of acidified sodium thiosulfate

solutions has received intensive study2-I2 and has been important in the development of the physical chemistry of colloidal solutions.6-10 The products of the reaction have been determined4-6s12 and various kinetic investigations have been made,4-6~8~11s12 but the results have remained es- sentially unexplained and difficult to interpret on a mechanistic basis. l1 Any proposed mechanism must be able to explain the products of the reaction, and the results of the kinetic studies: the reaction order and the effect of salts upon the rate.

Products.-The products of reaction 1 are sulfur dioxide, colloidal sulfur, Sa, Sa, various polythionic acids of the formula H&06 and hydrogen sulfide.

S203- I- H + + complex mixture (1) The nature of the products has been determined by Bassett and Durrant5 and confirmed by other workers. Several forms of sulfur are produced in the reaction; much of the sulfur occurs as Sa and some as insoluble sulfur which is a polymeric form having long sulfur chains. Hexatomic sulfur, &,, is produced in small yield from the decomposition of concentrated solutions of sodium thiosulf ate and

(1) h-ational Institutes of Health Predoctoral Fellorv, 1937-1998. (2) H. Landolt, Ber., 16B, 2958 (1883). (3) G. Foussereau, A n n . chim. p h y s . , [ U ] 15, 533 (1888) (4) J . Scheffer and F. Bohm, Z. anorg. Chem., 183, 151 (1929). (5) H. Bassett and R . G. Durrant, J . Chem. SOC., 1401 (1927). (6) C . K. Jablczynski and Z. Warszawska-Rytel, Bull . SOC. chim.

(7) G. Oster, J . Cviioid Sci., 2, 291 (1947). (8) E. Ll. Zaiser and V . K. La Mer, ibid., 3, 571-588 t.1948). (9) V . K. La Mer and A. S. Kenyon, ibid. , 2, 251 (1947). (IO) I. Johnson and V. K. La Mer, THIS JOURNAL, 69, 1184 (1947). (11) R. H. Dinegar, R. H. Smellie and V. K. La Mer, ib id . , 73, 2050

(12) F. Prakke and E. Stiasny, Rec. f r a u . chim., 54, 015 (1933).

FTnnce, 141 39, 409 (1926).

(1951).

hydrochloric acid in the This form has been characterized by molecular weight deter- minations, by its greater reac t i~ i ty , ’~ by its ultra- violet s p e c t r ~ m ~ ~ J ~ and by X-ray diffraction pat- terns.l6 The production of the polythionic acids can be increased by the addition of varying amounts of arsenic tri0xide~~t18 to the reaction mix- ture. The presence of hydrogen sulfide formed during the reaction escaped the attention of many of the early workers but has been detected in small amount^.^,^^ The yield of hydrogen sulfide has always been low; this would be expected because sulfur dioxide and hydrogen sulfide react in water, producing among other products sulfur. 18,20 The results of the radiochemical study of reaction 1 demonstrate that the two sulfur atoms of thiosul- fate are not equiva1ent2l and that during the prep- aration of thiosulfate and during the decomposi- tion in acid the sulfur atoms maintain their identi- ties and individualities.

(13) 9. H. W. Aten, Z. physik . Chem., 88, 321 (1914). (14) P. D. Bartlett and G. hleguerian, T H I S JOURNAL, 78, 3110

(1956). (15) R. E. Whitfield, Thesis, Harvard University. (16) C. Frondel and R. E. Whitfield, Acta Crys t . , 3, 242 (1950). (17) A. Kurtenacker and K. Matejka, Z . aizorg. Chem., 229, 19

(1936). (IS) D. XI, Yost and H. Russell, “Systematic Inorganic Chemistry,”

Prentice-Hall, New York, N. Y . , 1946, pp, 387-399. (19) J. Janickis, Z. anorg. Chem., 234, 193 (1937). (20) T h e reaction of hydrogen sulfide with sulfur dioxide in water

produces a solution known as Wackenroder ‘s Solution. IXistorically , the polythionic acids were Erst observed in this reaction mixture. Piumerous investigations on this reaction have been made since 1850; however, the chemistry is much in doubt and the mechanism of t h e reaction is not known. T h e same solution can be prepared by thc hydrolysis of sulfur chlorides in water. Wackenroder’s Solution w11 cold vulcanize gum rubber.

(21) H. H. Voge and W. F. Libby, THIS JOURNAL, 69, 2474 (1937) E. B. Andersen, Z. p h y s i k . Chern., B32, 237 (1936).

Vol. SO 3566 ROBERT EARL DAVIS

(2) ( 3 )

SO,‘ + “S --f “S-SO,‘ J6S-S03’ + 2H+ --+ 35S +

Kinetics.-The most recent and reliable kinetic investigation of the reaction has been that of La Mer and co-workers.8-11 The decompwition was studied spectrophotometrically in dilute aqueous solutionsJ1 a t constant ionic strength. Measure- ments were made a t 400 mp, and the time of ap- pearance of discrete sulfur particles was determined when the Tyndall beam appeared. The amount of sulfur produced a t the time of Tyndall beam ap- pearance was estimated to be of the order of 3.9 X lo-’ molar (based on s8). The results were found to be expressed in terms of the rate of sulfur ap- pearance as

rate = k’ ( Na2S203) 3/2( HCl)’/x

Such a rate expression was difficult for the investi- gators to interpret.ll The reaction was followed further by titration8 l 1 and the rate of production of sulfur dioxide was iound to obey the relationship

’The addition of salts demonstrated a positive salt effect.

Interpretation. Production of Sulfur.-The re- action scheme which can explain the experimental data is based upon the steps

( 4 )

rate = k ” ( ST&03)?(HCl)1 ( 5 )

A-9

s2o3-2 + H + 11S-o~- (0) ki

k?

ki

f ir

f i ,

HSzOy- f SL&-’ --f EISaOj- -t ( 7 )

11s303- + sio9-* + EISiOa- so,-’ (x )

HS4Oj- f S201-2---+ H S s 0 3 t SO$-‘ ( 9 )

€IS,Oa- 1 S201-2 -+ I I s , o ~ - + so3-’ (10)

FIS~O?- t S - C ) ~ - ~ -+ I IS ,O~- t so,-7 ( I 1 1 kti

€ISTO,- I S203-’ -3 HSyO, -1 SO? (13)

I I S s 0 , - + SiOp --+ IISoO?- + (13)

( 1 k )

k ,

k I - l S 0 0 3 - + s s + IISO3-

Derivation of Rate Equation.-The feature of equations 6-14 is that a series of bimolecular reac- tions produce the precursor (HS903-) of sulfur. This precursor then reacts by a unimolecular reac- tion. The nature of the intermediates d l be considered in the following c2iscussion.

Stage I.-The estent of redction when the first turbidity of sulfur appears is only 0.01%. The initial concentrations of thiosulfate and hydrogen ion are about molar and the concentration of sulfur produced is of the order of lo-’ molar based on S8.*’l Thus (HS20?-) and (S203-?) are con- stant. This fact allows the complete solution of the rate equations derived from the reaction scheme. Let x = (HS903-), B = (HSsOa-), C = (HS703-) . . . G = (HS303-), S = and T = (HSz03-), S and T constant, then

(ID d t k4ES - k6DS

CE = k3FS - kaES

df = k2GS - k3FS

(20 )

(21)

( 2 2 )

d t

d t tl G tit

= k l T S - kiGS

Several methods can be used to solve the system of equations 15 to 22. One treatment would consider the system as a sequence of first-order processes because S, the thiosulfate concentration, and T , the bithiosulfate concentration, are constant during the initial stage of the reaction. However, equa- tion 22 is a linear first-order differential equation in G ( T and S constant) which can be integrated from t = 0 to t and evaluated using the boundary conditions that G = 0 a t t = 0. The result can be expressed as

k,GS = klTS(1 - e--k2SL)

This solution is then put into equation 21 which is another linear first-order differential equation in F. Integration and evaluation with the appropriate boundary conditions give kJ7.S = klTS (1 - ~ ~ 3 ~ ~ ) +

(23)

(24) k k? - kz

Substitution of equation 2.1 into 20 and integration gives k4ES = kiTS ( I - e k 4 S t j -t

k l r S --< [ e - k d t - e--kzSt]

e - k h S t e - k a S t ksk4 kiTSkc ___ - -1 + klTS- k3 - kz [ k r - k0 k , -- kn

The form of the equations can be seen to be

This operation is repeated until an expression for k7BS is obtained in terms of KITS and k , ( j = 2 , ~ . 7 ) . Equation 1 6 is then solved for x in the same nianner and k x introduced into equation 15

<.’’ = k l T S ( l -- e d t

k , Z S = klTS(1 - c -k iS t j + klTSf(kz , k3, . k l , e - k J S f )

’ t ) + k lTSf (k , , S, t ) 126)

were f (k,,S,t) is a composite function like the terms of equation 25. Integration of (26) with the boundary conditions that Ss = 0 a t t = 0 gives

f - k t __ 1 .SS = klT.S[! - - - k --] + klT.SF(k,, S, t ) ( 2 7 )

where the function F is a complex summation of terms. La Mer” obtained the first term of equa- tion 27 by considering only one bimolecular reac- tion followed by a unimolecular decomposition. The data were then treated by plotting the con- centration of sulfur versus

e - k t - 1 [ t + T I

The resultant linear plot had a slope of klTS and an intercept, which in their figure (cf. Fig. 3)” is not identically zero, but very close to zero. The inter-

July 20, 1958 DISPLACEMENT REACTIONS AT THE SULFUR ATOM 3567

cept represents the second term of equation 27, which is now seen to be small. No assumptions have been made in the derivation concerning the values of the various rate constants. However, many of the rate constants ought to be nearly equal because of the similarity of the chemical inter- mediates. If kz = k~ = k4 the complex terms of equation 25 are zero, leaving only the first term. Expressing the concentrations of bithiosulfate ( T ) and the thiosullate ions (S) in terms of the initial concentrations of sodium thiosulfate and acid, the following equations are needed

I(H+)o - (HSzOa-)I [(s~o3-~)0 - (HSzOa-)I [HSzOa-] (28 )

where (H+)o is the initial acid concentration, 6%- 0 3 - 2 ) o the initial thiosulfate concentration, and Kz is the second ionization constant of thiosulfuric acid. Neglecting terms in (HS~OS-)~

and

Substituting (29) and (30) into (27) gives (sz03-') = (S~oa-~)o - (HSzOa-) (30)

Further approximation of (31) can be obtained by expanding the exponential, e-kt , into the first three terms of a power series (in Stage I of the reaction both k and t are small). The first term of (31) becomes

(32) can be transformed into

(33) 1 kl/z(H +)o'/z[Kz + (Sz03-2)o] ' 1 2 - = _ _ . ~ . t (2Ss ) ' /~[Kz + (H+)o + (SzO3-')01

Expanding [IC2 + (Sz03-2)0]1/z in power series and neglecting terms above the square gives

(34) 1 - a ( H + ) O ~ / ~ ( S Z ~ ~ - ~ ) O " n > 1 t

As the amount of sulfur produced a t the time of Tyndall beam appearance is constant," l / t is pro- portional to the rate. The higher terms of (31) become of the form

A =(H+)'/s (k,S + k,2S2 + kq3S3 + . . . . ) (35) 1

where the constants k,, k,, k , are functions of the rate constants kz, . . , . . , k7. The individual rate constants of these intermediate reactions are not known; however, they would probably be of the same order of magnitude. Therefore, only the first few terms of (35) would be important. An estimation of the values of k l , . . . , k7 and the con- vergence of the series could bound n within the range -2 > n > 1. The form of the rate expres- sion during the initial part of the decomposition was found experimentallysoll to be

rate 0: (II + ) ' / z ( S ~ O ~ - ~ ) ~ / P

compared to the predicted (equation 34) rate a(H+)'/~(S20~-Z)n 2 > n > 1

Thus the suggested reactions 6 to 14 do predict the rate equation for the initial production of sulfur.

Stage 11.-When the extent of reaction pro- ceeds further, S and T can no longer be considered constant. An analytic solution to the system of competitive, bimolecular reactions cannot be ob- tained in closed form. However, the steady-state approximation can be made: dG/dt = G' = 0

kz GS = kl BS. Then = F' = E' = D' = C' = B' = 0, then klTS =

@ = klTS - k x d t

This pair of equations can be solved in closed form by introduction of dimensionless parameters, K and r , and the results expressed in terms of IC, r and the exponential integral

However, in the present discussion only the order of the rate expression is required. As the time be- comes large, e-kr terms become small and the rate reduces to KITS. This result also can be obtained from the steady-state treatment assuming dx/dt = 0 in the pair of preceding equations. Express- ing T and S in terms of the initial concentrations, the rate of production of sulfur (which is propor- tional to the production of sulfur dioxide) becomes

rate 0: (H) (SZOS-~)~

which was found experimentally. Discussion

Having analyzed the kinetic situation and fitted the series of reactions to the observed rate data, the nature of the intermediates and chemistry in- volved should be considered, The equations ex- plain the positive salt effect, indicative that ionic species of the same charge type are reacting, and the observed dependence upon the acid concentra- tion is explained in terms of a bimolecular reaction between bithiosulfate and thiosulfate. The thio- sulfate can be visualized as undergoing a nucleo- philic reaction displacing sulfite ion which then forms sulfur dioxide upon protonation. The sys- tem of equations 6 to 14 can be complicated further by including all the various protonated forms. The sulfite and thiosulfate groups have been ob- served to undergo displacement reactions with one another22-26 and several reactions have been stud- ied kinetically using radiosulfur. Equation 7 can be described in terms of the displacement reaction

ki ~ ~ 4 0 3 - + s-so~-~ ----f HS-S-SO~- + S*oa-2 (34a) and the following reactions displacing the -SOI- group on the end of the sulfur chain. The poly- sulfur intermediates

(22) D. P. Ames and J. E. Willard, THIS JOURNAL, 76, 3267 (1953); 73, 164 (1951). (23) J. A. Christiansen and W. Drost-Hansen, Nature, 164, 759

(1949). (24) A. Fava, Gasa. chi,. i fal . , 83, 87 (1953). (25) A. Fava and G. Pajaro, THIS JOURNAL, 78, 5203 (1956).

3 3 3 ROBERT EARL DAVIS Vol. so * ka

H-S--S--S-SOI- + S-SOI -2

* H-S-S-S-S-S0a- + 503-2 (35a)

would resemble the hydrogen polysulfides (poly- sulfanes) in many respects and would be expected to have a weakly acidic hydrogen (H-S-S,-). There is no a priori reason to exclude the attack of the thiosulfate in the middle of a sulfur chain; no doubt such reactions do occur

II-S,-S-S,-S03- + s203-2 --+

The formation of the S8 ring in a unimolecular reac- tion could be formulated as

S I - 5 ' k S

H-Sz-S-S203- + -S,-SOI- ( 3 G )

/S--SH Sos- p5-S- SOxH

,s t- S I 4 - i

,S-s \

\s--s

S s 1 + HSOI- ( 3 7 ) S 9

The precursor of sg would be HS703- which then would undergo an internal displacement forming the S6 ring. The yield of SS would be predicted to be low because of the instability of the SO ring com- pared to the Ss ring14,z6; thus, the frequency factor of HS703- forming Ss would be lower than the chance of producing longer sulfur chains by dis- placement reactions with thiosulfate. The longer sulfur chains of insoluble sulfur would be built up by further displacement reactions. The true nature of these chains and the nature of termination are not known. The formation of macro rings or/and chains of H-S-SZ-S-S03H ( x large) are possibilities but the current experimental data give no indication of the character of these materials.

The formation of the polythionic acids can also be explained using thiophilic, nucleophilic displace- ment reactions. The sulfo-polysulfaneZ7 inter- mediates could produce the polythionates by using the sulfide sulfur as the displacement reagent. Thus the production of pentathionate could be for- mulated

I ~

! so, - SOIH

SH S- ---+ S-SH

s I ' s + 60,- --+

S-S + -SH I I s SO,- 1

SOiH

( 3 8 )

Fragmentation of the sulfur chains by attack in the middle of the sulfur chains forming polj7sulfides and polythionates are likely reactions also. The species of (38) could be in various stages of proto- nation. Hydrogen sulfide is never produced in

(21;) L Patiling, Pvoc. Y u l . A c a d . Sci. , 35, 455 (1549). ( 2 i ) Cuniyounds of the type H-S2-S0,H have been recently rc-

ported and s<,me reactions discussed. I n many cases the reactions are explicable on the thesis of nucleophilic displacement reactions a t the sulfur atum ( M . Schmidt, Z . aizovg. Cheii;., 289, 158 (1557)).

yields equivalent to that of the polythionic acids; however, hydrogen sulfide and the polysulfanes re- act with sulfur dioxide producing sulfur and other products. *O Addition of arsenic trioxide increases the yield of the polythionic Arsenic (111) would shift the equilibria of reactions like (38) by removing the hydrogen sulfide as arsenic sulfide. In dilute aqueous solution La Mer8," estimated that sulfur is the main product of the reaction with only small amounts of polythionate formation. Under controlled conditions, using arsenic trioxide, rather good yields of the poly- thionic acids can be obtained.".ls

The Decomposition in Concentrated Solutions.- The kinetic results were obtained in dilute aqueous solution of the order of molar in the reactants. The effect of large amounts of acid has been meas- ured by many investigators. If cold, concen- trated hydrochloric acid is used in large excess, there is an immediate precipitation oi sodium chlo- ride, and the solution remains clear and colorless for eight hours.19 Rassett and Durrant5 obtained minima in their ctirves of sulfur appearance time wrszis the concen tration of hydrochloric acid added. The miniilia occur in the region of one molar acid. The facts can be described by the effects of hydrogen ion upon the equilibrium 6 and the \-arious other reactions which produce decom- position. Thiosulfuric acid in very strong hydro- chloric acid appears to be somewhat stableIg and much less stable in phosphoric a c d 5 FossZY pre- dicted that pure thiosulfuric acid should decompose into hydrogen sulfide and sulfur trioxide

liiSro:3 -+ 11:s -t so.

Recently, has prepared solutions of thio- sulfuric acid by the decomposition of sodium thio- sulfate with HCl in ether a t -78". A dietherate of thiosulfuric acid was obtained which decomposed a t -5" by reaction 39. Likewise, the reaction of hydrogen sulfide and sulfur trioxide in ether a t low temperatures produced thiosulfuric acid. Reac- tion of the polysulfanes (H.S.?) with sulfur trioxide in ether a t -178" was found to produce either the polysulfane monosulionic acids, H-S,-S03H, or the polythionic acids, HS03-S,-S03H. The poly- sulfane rnoiiosulfonic acids have been inferred to be prcsent in the decomposition analyzed in this re- port.

The stepwise build up of a sulfur chain by dis- placement reactions is consistent with the kinetic results of a series of reactions in which a sulfur chain is degraded atom by atom. The reactions of triarylphosphines,'4,3u cyanide ion,31 sulfide and sulfite32 with sulfur h a w been studied kinetically in homogeneous solutioii. In each case the sulfur ring is opened and then degraded by the thiophilic reagent one atom of sulfur a t a time. The sug- gested mechanism of formation of the polythionates is consistent with the radiocheniical data of Brodskii

(3'31

( 2 3 ) 0 P,ISS, 7 ' i . i ~ . Kjiri;i H c i y n p s r i i .\It*/.. 6, 3 (194tj), Chiin. , 289, 141, 158, 175 ( 1 l ) ; i ) a n d R. 12 Davis. unpublished data

) 1' 1) U u r t l c t t : d m ! I<. 1; Davis, THIS J O U R X A L , 80, 2.518 (19,S).

( 3 2 ) R. E. Davis, unpublished da ta .

July 20, 1958 STUDIES OF ZIRCONIUM MANDELATES 3569

and E r e ~ n e n k o . ~ ~ It was found that the deconi- position of *SS03-2 in the presence of acid and sodium arsenite gave -O&3-*SZ-SO3-. The gen- eral over-all decomposition was found to be

- l o o

arsenite *S-s03-* + H+-+

-03s - *S, - SOX- + HzS* + sulfur

Estimation of the Rate Constants.-The values reported for the dissociation constants of thiosulfuric acid have been near 0.45 (K1) and in the range 0.01 to 0.OG2 for the second dissociation (Kz) in dilute solutions. l1 The previous investigators1' used the value of 0.01 for K1 and calculated the rate con- stants to be k = 0.14 * 0.02 min.-' and kl = 0.31 1. mole-' min.-' during the initial stages of the reaction and as 0.29 1. mole-' min.-l from the titrametric data from the latter stages of the reac-

(33) A I Brodsku and R K Eremenko, J Gcx Chem ( U S S R ), 24, 1137 (19541, 26, 1188 (1955)

tion based upon the gram atom of sulfur. The values of the rate constants are dependent upon the value of Kz but the values calculated are correct within the order of magnitude of the data. A reasonable assumption can be made concerning the other bimolecular reactions. These reactions are probably somewhat faster than kl and probably k3 - k b - k5 - . , . because of the very similar nature of the intermediates. Assuming k~ = hd = ks . . . makes many of the higher terms of (26) and (27) zero, thus simplifying the expression and increasing the importance of the first term. The rates of formation of the various polythionates as equation 38 cannot be obtained from the data but in dilute solution these reactions are much slower and the sulfur is the main product of the decom- position.

The author wishes to thank Professor P. D. Bartlett for helpful discussion and permission to publish this interpretation. CAMBRIDGE, MASSACHCSETTS

[CONTRIBUTIOS FROM THE CHEMICAL LABORATORIES, UNIVERSITY OF LUCKNOIV]

Organic Compounds of Zirconium. V. Studies of Zirconium Mandelates BY R. N. KAPOOR AND R. C. XEHROTRA

RECEIVED DECEMBER 4, 1957

The reaction of zircoiiyl chloride with mandelic acid in aqueous solution yields mainly zircoiiiuni IiionoIiiatidelate ivhicii dissolves in one equivalent of alkali. The tetramandelate is formed only in the presence of higher concentrations of hydro- chloric acid. The reaction between zirconium isopropoxide and mandelic acid in benzene shows that the carboxyl group is more reactive than the hydroxyl group in replacing isopropyl alcohol; the latter becomes active only nhen a sufficient con- centration of carboxyl groups is not available. The following new compounds of zirconium have been isolatcd : Zr(C.HeOaj- (OPr-iso)z, Zr(C8H603) (CsHi03) (OPr-iso), Zr(CsH?03)~(0Pr-iso)z and Zr(CsHsO3) (CsHiOs) (OBu).

The preparation of zirconium alkoxides'.? and tricarbo~ylates~ has been reported recently. The alkoxides and tricarboxylates are immediately hy- drolyzed by water, and even the zirconium salts of the strong acids are appreciably hydrolyzed in solu- t i ~ n . ~ Therefore, the precipitation of pure zirco- nium tetramandelate from solutions of zirconyl chloride in concentrated hydrochloric acid5 is note- worthy in view of the weakly acidic nature of mandelic acid (dissociation constant at 25' = 4.28 X To explain this, FeigP postulated the formation of a chelate type structure

C~H~-CH-C=IO

€IO 0 I I I /

Zr/4

This structure explains the solubility of the tetramandelate in dilute alkaline solution as due to the enhanced acidity of the weakly acidic hydroxyl group through coordination.

Hahn and Weber,' in a conductometric titration (1) D. C . Bradley, R. C. Mehrotra and W. Wardlaw, J . Chenz. SOL.,

( 2 ) R . C. Mehrotra, THIS J O U R N A L , 76, 1266 (1984). (3) R. N. Kapoor and R . C . Mehrotra, Chern. & I n d . , 08 (1958). (4) B. A. J. Lister and L. A. Donald, J . Chem. SOC., 4315 (1952). ( 5 ) C. A. Kumins, A n a / . Chem., 19, 376 (1947). ( 0 ) F. Feigl, "Chemistry of Specific, Selective and Sensitive Re-

(7) R B. Nahn and L. Weber, 'lkrs JOURNAL, 77, 4777 (1955).

2027, 4204, 5020 (1952); 1ti34 (1053).

actions," Academic Press, Inc., New York, N. Y., 1949, pp. 213-215.

of zirconyl chloride solution with sodium mandelate, observed a sharp maximum when the ratio of mandelate to zirconium was 2.3: 1. To explain their observations, they assumed that the following reactions occur in solution

ZrO'+ + 2HCsH;O3 --+ ZrO(CsHiOa)2 f 2H- (soluble)

ZrO(C6H;03)2 + 2HCsH70d --f Zr(CbH;03)1 + H 2 0 However, it has been reported recently"9 that

zirconium monomandelate preferably is precipitated when zirconyl chloride is treated with mandelic acid or sodium mandelate solution. The nionoinan- delate could have either structure

I I1 The easy hydrolysis of the alkoxy oxygen -zirco- nium linkage suggests that I1 is probably correct. Furthermore, in the titration of zirconium inono- mandelate with sodium hydroxide (Fig. l), we found that that monomandelate dissolves when about 1 mole of alkali has been added; this can be explained by the neutralization of the hydroxyl hydrogen atom which has been made more acidic (dissociation con-

(8 ) R. B. Hal in and IS. S. Baginshi, A i i o / ( ' h i r i r A<Lc, 14 , 42 (105t;) (9) R. h' Kapoor and R . C. bIel i rutra , J . Sei, i i idusii . . Rt.s , lCB,

300, 304 (1937) .