The stabilities of some transition metal 1-hydroxycyclopentanecarboxylic acid complexes

J. inorg,nuel.Chem.,1967,Vol.29, pp. 1729to 1733. PergamonPressLtd. Printedin NorthernIreland

THE STABILITIES OF SOME TRANSITION METAL 1-HYDROXYCYCLOPENTANECARBOXYLIC ACID COMPLEXES* J. E. POWELL a n d D. L. G. ROWLAr~DS Institute for Atomic Research and Department of Chemistry, Iowa State University, Ames, Iowa

(First received 26 September 1966; in revised form 9 January 1967) A I n ~ c t - - T h e formation constants of the first two species formed by some dipositive transition metal cations with the 1-hydroxycyclopentanecarboxylate(HCPC) ligand have been determined. K1 and Ks, the first and second successive stability constants for the systems investigated, are respectively: Cu, 630 4- 6, 60 q- 3; Ni, 66 4- 2, 20 4. 3; Co, 37 4- 1, 10 4- 2.7; Zn, 77 4- 2, 20 4. 3; Cd, 28 4. 1, 8.5 -4- 3-9. Measurement was by a potentiometric method at 25"C, and a constant ionic medium == 0.1) was maintained by the addition of the supporting electrolyte, sodium perchlorate. The proper acid dissociation constants (Ks) ranging from 1-13 to 1.09 × 10-4 were interpolated from a plot of Ks against the anion concentration [A-]. A discussion of the use of a variable acid dissociation parameter is included. Dehydration studies on the solid species were also undertaken to establish the hydrate numbers.

INTRODUCTION THE SUCCESSIVE stabilities o f the r a r e - e a r t h H C P C chelates have been r e p o r t e d , tl~ T h e results i n d i c a t e d a c o o r d i n a t i o n n u m b e r o f nine f o r the r a r e - e a r t h series a n d a c h a n g e f r o m t r i d e n t a t e to b i d e n t a t e c h e l a t i o n in m o v i n g f r o m the light to h e a v i e r l a n t h a n i d e s . T h e p r e s e n t investigation is a n a t t e m p t to m e a s u r e a n d correlate the stabilities o f s o m e d i v a l e n t t r a n s i t i o n m e t a l c a t i o n (Co, Ni, Cu, Zn, Cd) complexes w i t h the H C P C ligand, a n d to c o m m e n t o n the a p p a r e n t n a t u r e o f the chemical b o n d s f o r m e d . A l i t e r a t u r e survey i n d i c a t e d n o p r e v i o u s m e a s u r e m e n t s h a d been m a d e w i t h these systems. EXPERIMENTAL Preparations. Sodium perctflorate and HCPC were prepared as described previoasly.C1~ The transition metal perchlorates were prepared using an ion--exchange technique. A weak acid cation exchange resin (Amberlite IRC-50) was loaded in a 1-in. Pyrex brand glass column which was 5 ft in length. A sodium hydroxide solution was pumped up through the column to convert the resin bed to the Na + form, and the column was washed thoroughly to remove excess Na + ions. Next, an excess of 0.5 M transition metal sulfate was passed through the column to saturate the resin with the divalent cation. Further washing with distilled water was done to ensure that the system was SO, 2- free. Then a solution of,--.O.3 M HCIO4 was passed through the column at 10 ml/min, and 1500 ml of the central portion of the eluate was collected. The molarity of the resulting solution was determined by passing aliquots through a cation exchange resin in H + form. The acidic eluate and washings were combined and titrated with standard carbonate-free potassium hydroxide, and these results were confirmed by direct titration of the metal salt solution with standard EDTA. t'~ The anion and the cation concentrations determined by these means agreed well. Procedure. A buffer solution (,--0.2 M) was prepared from HCPC acid by half neutralizing with a measured volume of standard carbonate-free sodium hydroxide and diluting to volume. Then the * Work was performed in the Ames Laboratory of the U.S. Atomic Energy Commission. Contribution No. 1841. ta> j. E. POW~LLand D. L G. ROWLA~,rOS,lnorff. Chem. 5, 819 (1966). ~ J. E. FRITZ, J. E. AnBr~K and M. A. PAYNe, Analyt. Chem. 33, 1381 (1961). 12 1729

J . E . POWELLand D. L. G. ROWLA~S

1730

uuneutralized acid remaining was determined precisely by titrating aliquots with standard KOH. Between 17 and 20 different volumes of buffer were added to a series of 100-ml volumetric flasks together with 4 ml of 0.1000 M transition metal ion solution. A constant ionic strength of 0.100 was maintained throughout by adding calculated volumes of standard sodium perchlorate solution. Solid compounds. Stoicheiometric quantities of the transition metal chlorides and ammonium salt of HCPC were stirred until precipitation was complete. The resulting pH was about 6, except in the case of Zn where excess HCPC was added to prevent the formation of basic zincate. The solids were washed free of C1- ion and air-dried before thermal decomposition studies were made. Thermal decomposition of individual salts was followed for 14 hr over a temperature range from 25 to 500°C, using a thermobalanc¢. Knowing the initial weight of each sample and the final oxide composition and weight, hydration numbers were calculated from recorded weight losses, with the tacit assumption that a metal ion retained three units of undecomposed ligand until the compound had lost its water of hydration. In each case a level corresponding very closely to the true formula weight of the anhydrous compound was observed. Measurement. The stability data were obtained potentiometrically at 25.00 =t=0"05°C using a Beckman Research pH meter with glass and saturated calomel electrodes. Activity coefficient corrections were unnecessary as the meter was standardized on a concentration basis with a/~ = 0.100 HC1-NaCIO~ solution at pHc = 3.712 =t=0.001, and checked for response within the pH range used, vs. similar solutions containing known amounts of mineral acid. That is to say, the H + concentration was known as a function of scale readings to 1 part in 500. The observed pile values are reported in Table 1. TABLE l:---OBSERVEDpile VALUBSOF TRANSITIONMETALHCPC BUFFERSOLUTIONS M1. buffer 100 ml sol. 0.50 1.00 1-50 1.75 2.00 2-25 2.50 2.75 3.00 3.25 3.50 3.75 4.00 4.25 4.50 4.75 5.00 6"00 7.00 8.00 9.00 10-00 12"50 15.00 17.50 20.00

Cu I+*

Ni *+*

Metal Co2+t

3.657 3.627

3.961 3.923

4.069 4.046

3.614

3"926

3.614

Zn2+.+

Cd2+:~

4.010 3.912

4.061 3.999 3.971

4.024

3"902

3.956

3.922

4.016

3.896

3.953

3.617 3.628 3.628

3.915 3.908 3.908

4.012 4.010 4.010

3.890 3.886 3.886

3.948 3.944 3.944

3.641

3.902

4.002

3.882

3.941

4.002

3.880

3.939

4.001 3'873 3"993 3"988 3"988

3.878 3.930 3'872 3"869 3.871 3.800 3.883

3"936

3'663 3"685 3.702 3.721 3"736 3"750 3'779 3"806 3.824 3"847

3"898 3.895 3.893 3.898 3.905 3.906 3.908 3.910 3.911

3"992 3"992 3"992 3"996

* Buffer solution, 0.1011 M salt-0"0963 acid; Cu~+, Ni s+, 0"1000 M. ? Buffer solution, 0.0920 M salt-0.0750 acid; Co s+, 0"I000 M. Buffer solution, 0.1011 M salt-0.0989 acid; Zn 2+, CdI+, 0"1000 M.

3'920 3"917 3'918 3.920 3"923 3"926 3"928 3'930

The stabilities of some transition metal 1-hydroxycyclopentanecarboxylic acid complexes 1731

Calculation. The formation constants 81 = KI ----[MA]/[M][A] and [~s = K1Kz = [MAI]/[M][A] 2 were calculated using a two parameter least squares curve fitting program ~*)on an IBM 7074 computer. These constants are equilibrium concentration ratios at a specific ionic strength (0-100M) rather than thermodynamic equilibrium constants. Computations were based on the treatment of L~D~ ~4) and FRONAEUS,(s) which has been reviewed by S o ~ s o N (e) and ROSSOT~ and RossoTn. (7) qC) X

~

t.l,'---:

Z

~ooO%~.. '

o ""~*

-* %

-

0 1.03

o

-

001 0.0 ~; 0.03 0.04 0.05 LIGAND ANION CONCENTRATION, [A-I, (moles/liter)

F[~. 1.--Ionization constants (Ka) vs. anion concentration for HCPC at 25.0°C and p = 0.100. To calculate stability constants it is necessary to set up material balance equations for total ligand, metal and available hydrogen, including all pertinent species thought to exist in the system. Let A = total concentration of chelating agent available; a = concentration of free (uncomplexed) ligand [A-]; B = total concentration of metal ion available; b = concentration of free (uncomplexed) metal [13-+]; H = total hydrogen ion (acid) available; h = concentration of hydrogen ion [I-I+]; Kt ----ionization constant of the acid. Equating the pertinent ligand, metal and hydrogen containing species results in the expressions:

A = b(flxa + 2f12as) + a(1 + h/Ka) B = b(1 +

flxa + fl2a2)

H = h(l + aIKa) From these quantities a value for fi, the average number of ligands attached to the central metal ion, can be calculated. t] = [A -- a(1 + h/Ka)]/B In practice H, A, and B axe known, h is measured directly, and a is computed using a supposed constant Ka. It was invariably observed that a plot of Ka vs. a (see Fig. 1) was essentially linear, but that Ka increased slightly as a approached zero. This trend has been observed for twelve different organic acids studied in this laboratory under conditions of constant ionic strength from p = 0.01 to (2) W. R. STAC~ and J. E. POWELL, USAEC Report IS-727 (1963). (4) L. Lm3EN,Potentiometrisk unders~kning ov nagra kadmiumsalters komplexitet (Diss.) Lund (1943). (5) S. FRONmUS, Komplex system hos kopper (Diss.) Lurid (1948). (G~A. SON~eSSON,Acta chem. scand. 12, 165 (1958). (~) F. J. C. ROSSOTn and H. RossoTri, Determination of Stability Constants. McGraw-Hill, New York (1961).

1732

J.E. POWELLand D. L. G. ROWLANDS

2"0. Perusal of a recent article by THUNet aL (8) reveals that they too encountered a comparable phenomenon in the ease of mandelic and atrolaetic acids. In the present method a buffer solution of known neutralization was diluted along with calculated amounts of NaCIO~ to provide 25 samples whose pH's were then measured. The whole procedure was repeated independently and all fifty datapoints(allat25"0°Candp = 0.100) are illustrated in Fig. 1. The systematic drift can be attributed, at least in part, to a form of salt effect dependent on the relative sizes of the ClOt- and HCPCionsJ °~ We have found it impossible to control this systematic variation of K~ in practice and have introduced statements in the computer program which provide for a selection of the proper Ks value for each individual data point. RESULTS AND DISCUSSION The first and second successive stability constants, their ratios, and the waters of hydration observed in the solid compounds are presented in Table 2. TABLE 2.-----STABILITY CONSTANTS AND THERMAL DECOMPOSITION DATA FOP. SOME TRANSITION

METALHCPC COMPLEXES

Transition metal

K~

162

CAt Ni Co Zn Cd

630 + 6 66 -4- 2 37 4- 1 77 + 2 28 4-1

60 5:3 20 5:3 10 4- 2.7 20 4- 3 8.5 5:9

KJK~

Hydrate Nos.

10.5 4- 0.6 3.3 4- 0.6 3.7 4- 1.1 3.8 5:0.7 3.3 4- 1.6 - -

Dehydration achieved (°C)

0 2 0 2 0

113 99

Onset of degradation (°C) 202 288 269 322 295

The hydration data lead to n o c o m m o n coordination number for the solid species, This is probably due to a change in structure during crystallization. WELLS(1°~ implies that, while these cations appear in six-coordinate oetahedral arrangements mostly, distortion can yield an essentially square planar structure with two additional sites far removed along the axis through the plane. Let us in this instance consider a conventional oetahedral model and apply BJEl~lJl's statistical ratio theoryC11~ for all possible modes of complex formation. The theoretical ratios are listed in Table 3. T~LE

3.~STATISTICAL RATIOS OF STABILITY CONSTANTS FOR AN OCTAHEDRAL CONFIGURATION OF COORDINATION SITES

Ligand type

Kx/K~

Monodentate Bidentate Tridentate

2.4 4'8 16

In the case of the rare earths and H C P C c1~we pictured a tripositive cation amply surrounded by water dipoles, and considered the metal-ligand bonds to be electrostatic in nature. The chelation process involved merely the replacement of several water dipoles from the coordination sphere by a uninegative but polydentate ligand. ts~ H. Tram, F. Vra~EEKand W. VANDEm.~EN,J. inorg, nucL Chem. 28, 1949 (1966). ~0~H. S. HARNEDand B. B. OWEN, ThePhysieal Chemistry of Electolytie Solutions (3rd Edn), p. 64. Reinhold, New York (1958). ~1o~A. F. WELLS, Structural Inorganic Chemistry. Clarendon Press, Oxford (1962). ~ix~j. BJEP.RUM,Metal Ammlne Formation in Aqueous Solution. P. Haase, Copenhagen (1941).

The stabilities of some transition metal 1-hydroxycyclopentanecarboxylicadd complexes 1733 Electrostatic bonding was supported by excellent agreement between the theoretical and experimental K1/K2 ratios observed with the rare-earth HCPC chelates; ~1~ however, in the present case, results vary widely from the purely statistical values. Admittedly, one cannot draw concrete conclusions in the cases of Ni, Co, Zn, and Cd where the maximum ti values achieved were only 0.80, 0.43, 0-55 and 0"38, respectively. Nevertheless, from the fact that the HCPC complexes formed by these four cations have such a low degree of stability, one is inclined to conclude that K1/K2 ratios falling between 2.4 and 4.8 are merely the manifestation of a failure to form a bidenate linkage in every event, i.e. coordination probably involves a mixture of mondentate and bidentate linkages. Cu ~+, on the other hand, forms relatively strong complexes with the HCPC ligand; so that one might postulate a combination of bidentate and tridentate attachments by the ligand to account for the K1/lc2z ratio falling between 4.8 and 16.0. Alternatively, one could assume purely bidentate behavior of the ligand and covalent rather than electrostatic bonding (due to a more than ten-fold increase in K1). Covalent bonding might introduce an electrostatic factor of approx, two by effectively neutralizing half the cationic charge of the cupric ion. Note that 10.5 ~ 0.6 = approx. 2 × (K1/K2)stat" = 9"6.