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An alternative mechanism for the formation of the cobalt(III) molybdate cation, Co(NH3)5MoO4+
Polyhedron Vol. IO, No. 19, pp. 2317-2329. 1991 Printed in Great Britain
0
0277-5387/91 $3.00+.00 1991 Pergamon Press plc
AN ALTERNATIVE MECHANISM FOR THE FORMATION OF THE COBALT(III) MOLYBDATE CATION, Co(NH,),MoO,+ MICHAEL
R. GRACE*? and PETER A. TREGLOAN
Inorganic Chemistry Section, School of Chemistry, University of Melbourne, Parkville, Victoria 3052, Australia (Received 7 March 199 1; accepted 5 June 1991)
Abstract-An alternative mechanism is proposed for the formation of CO(NH~)~MOO~+ from CO(NH~)~H~O~+ and MoOd2- in slightly basic solution (pH 7.1-8.0). The observed second-order rate dependence upon [MOODY-]may be explained by HMoO,- acting as a proton-donating catalyst, in contrast to the previously postulated mechanism involving the formation of dimeric molybdenum(W) species. An increase in reaction rate was also caused by the addition of other protonated anions (HC03-, H,PO,-) to the reaction solution, thus supporting the contention of a general proton-assisted mechanism. No evidence of dimeric molybdenum(V1) was found.
Taylor has reported that the formation of Co (NH3)5M004+ from MoOd2- and CO(NH~)~OH*+, in the pH range 7.1-8.0, could be conveniently followed using the stopped-flow technique.’ It was also found that the complexation reaction, given in eq. (l), showed a greater than first-order dependence upon [MOODY-].
Co(NH&Mo04+
+H20
(1)
The forward rate constant, kf, has four terms, according to eq. (2) : kf = k,+k,[H+]+k,[Mo0,2-]+kJH+][Mo0,2-]; (2)
where: k, = 96+7 M-’ s-l, kb = (1.1 f0.2) x 10’ Me2 s-l, k, = (2.2+0.2)x lo3 M-2 s-l and kd = (1.02f0.05) x lo” Me3 SK’ at 25°C and p = 1.0 M (NaClOJ. Taylor also noted that the rate of reaction was consistent with substitution at the molybdenum(V1) rather than cleavage of the inert cobalt(III)-oxygen bond. It was suggested that the pathways k, and kb correspond to the reaction of HMoO,with respecCO(NH~)~OH~~+ and Co(NH3),0H2+, * Author to whom correspondence should be addressed. t Present address : Chemistry Department, University of Calgary, Calgary, Alberta, Canada T2N lN4.
tively. The second-order [MoO,~-] dependence for paths k, and kd was ascribed to a reaction involving the dimolybdate species Mo20g4-, HMo~O~~- or H2M020g2-, depending upon the active cobalt(II1) entity, in order to conform to the observed acid dependence. The dimolybdate anion, M0207~-, has been isolated from acetonitrile solution as the (n-C4Hg)4N+ salt, containing corner-shared MOODY- tetrahedra.* However, Mo20T2- has never been detected in aqueous solution, although a recent study by Cruywagen and Heyns3 reported evidence of the protonated species HMo207- under mildly acidic conditions (2.3 < pH < 6.1). The discrete dimeric species, Mo~O~(H~O)~~+, has been identified as the major component of molybdate solutions in 2-6 M HC104 and HN03 by g5Mo NMR and Raman spectroscopy.4 The structure involves two molybdenum(W) centres bridged by a single oxygen. Two terminal oxygens and three aquo ligands constitute an octahedral environment around each molybdenum ion. However, the existence of an aqueous dimeric molybdenum(V1) species, under normal conditions, must remain tentative since another recent potentiometric study’ found no molybdate oligomers between monomers and the heptameric series Mo~O~~H~(~-“- (n = 0, 1, 2 and 3). The major objective of the present study was to reinvestigate the complexation between MOODY-
2317
2318
M. R. GRACE and P. A. TREGLOAN
and CO(NH~)~OH~~+ in basic solution, in order to investigate Taylor’s postulated reaction between cobalt(II1) ions and dimeric molybdenum(V1) species. The presence (or absence) of a term involving dimolybdate in the rate expression could then be used as independent evidence for its existence and assist in the continuing development of the models used to describe molybdenum(V1) speciation in acid solution. We thought that the second order term in [MOODY-]might also arise from HMoO,- acting as a general acid catalyst. Consequently, the effect of a range of anions upon the rate of the complexation reaction was also examined. EXPERIMENTAL Materials
[Co(NH 3)SOH 2](ClO4)3 was prepared by established methods.677 W-vis spectra were in good agreement with those reported previously.7*8 Na,MoO, - 2H20 (May & Baker, LR) was twice recrystallized from base solution (pH x 11). Triethylanolamine (Ajax, LR) was distilled in vacua. Na2HP04 (Ajax, AR), NaHCO, (BDH, Analar), Na,SO, (BDH, Analar) and NaOH (Ajax, AR) were used without further purification. Twice distilled, once deionized water was used throughout Both equilibrium and kinetic measurements were performed in the presence of 0.1 M triethanolamine in order to buffer the solutions at the required pH and prevent the formation of polymeric molybdenum(V1) species. ’ Solutions were adjusted to the desired pH by addition of either 2 M HClO, or 2 M NaOH. These solutions showed negligible drift in pH (< 0.05 pH) units) over a period of several months. The synthesis of [Co(NH3),Mo04](C104) involved mixing slightly alkaline 0.05 M solutions (8 < pH < 9) of MoOd2- and CO(NH~)~OH~+ at 5&6O”C in the presence of ca 0.5 M NaClO,. The mixed solution was maintained at this temperature for a further 15 min. The pink solid formed upon cooling was filtered and then washed with ice-cold water, CH3CH20H and then (CH3CH2),0 and dried at 60°C for several hours. The crude complex was recrystallized from a minimum of hot (x 80°C) water at pH 8, filtered, washed and dried as above. Yields were typically 3040%. The analysis of NH, content has been described *The term “pH”, used throughout this paper, refers to the quantity -log,o[H+]. t All concentrations are at the “instant of .mixing” = stock concentration/2.
previously.’ For [Co(NH,),MoO,](ClO~) : Found : 20.6. Calc. : 21.1%. The residue after product decomposition was filtered and the molybdenum concentration in the filtrate was determined gravimetrically as Mo02(CvH60N), after complexation with 8-hydroxyquinoline.” Found : 23.5. Calc. : 23.8%. Physical methods
All pH* measurements were performed using a Radiometer GK240 1C combined glass electrode, in which the saturated KC1 reference solution was replaced with 1 M NaCl to prevent the precipitation of KC104 at the liquid junction. The electrode was calibrated daily with a series of acid (HC104, pH 2-3.5, p = 1.0 M) and base (NaOH, pH 11-12.5, p = 1.0 M) standards. The slope of the pH vs electrode potential plot was typically 0.059+0.002 V per pH unit. Absorbance measurements were performed on a Varian Techtron 635D UV-vis spectrophotometer, using a cell block thermostatted to 25.O+O.l”C. Spectra were recorded in a quartz mixing cell (Hellma, 238-QS) which is divided into two equal path-length compartments by a thin wall ; a hole at the top of this wall allows mixing of the two reactant solutions. Consequently any absorbance change upon mixing can be directly attributed to chemical reaction. IR spectra were recorded as KBr discs on a Jasco 302 IR spectrometer. Raman spectra were obtained using a Spex Ramalog 5 Raman spectrometer interfaced with an M6800 microcomputer. During Raman studies of molybdate/phosphate solutions, NaC104 was added to each solution to maintain an ionic strength of 1.OM and also provide an intensity calibration for each spectrum [at v,(C104-) = 935 cm- ‘I. HCl was used to obtain the desired solution PH. Diffuse reflectance spectra were recorded over the range 34&750 nm on a Beckman DK-2A spectrophotometer (BaS04 reference). Variable temperature magnetochemical measurements were performed over the range 21&350 K on a thermostatted Gouy balance, calibrated with Hg[Co(NCS),]. Diamagnetic corrections were applied according to Kiinig. ’ I Rate measurements The formation of CO(NH~)~MOO~+ from MoOd2- and CO(NH~)~OH~~+ was followed spectrophotometrically at 550 nm, using the stoppedflow technique. The cobalt(II1) concentration was maintained at 0.0020 M,t whilst [MOODY-] was
Formation
of cobalt(II1) molybdate cation
varied between 0.010 and 0.100 M. All rate measurements were made at an ionic strength of 1.O M (NaC103 and at 25.0 + 0.1 “C over the pH range 7.1-8.0. A stopped-flow apparatus, designed and constructed within the Chemistry Department at Melbourne University, was used to perform kinetic experiments with half-lives ranging from 10 ms to several minutes. The molybdate and cobalt(II1) reactant solutions were drawn into 2 cm3 glass, gastight drive syringes (Hamilton, 1002-TEFLL) from plastic reserve syringes via three-way taps (Hamilton, No. 86727). A kinetic experiment (“run”) was initiated by triggering a solenoidcontrolled pneumatic drive, which forced fresh reactant solutions through a Kel-F mixer” and into the observation chamber. Up to 20 runs can be recorded before refilling the drive syringes. The observation cell consists of either: (a) a “straight-through” 2 mm internal diameter quartz tube where the light-beam passes through the tube at a 90” angle, or (b) a “Z-shape” 18 mm pathlength chamber in a Kel-F block, enclosed into the brass stopped-flow block by quartz windows. Choice of cell is governed by the magnitude of the absorbance change and the intensity of the transmitted light. Optically dense solutions necessitated using the short path-length tube to ensure adequate signal-to-noise response of the photomultiplier detection system. The stopped-flow unit is encased in a brass block through which thermostatted water is circulated, allowing temperatures between 7 and 45°C to be maintained to within 0.05”C. A Hewlett Packard 34740A platinum resistance thermometer is inserted into a cavity in the brass casing to enable temperature resolution of 0.01 “C. All reactant solutions were thermostatted within the drive syringes for 20 min to ensure thermal equilibration. A 100 W quartz-halogen lamp, powered by a Hewlett Packard Model 6286A supply, provided the light source for the spectrophotometric detection of reaction progress. A Czerny-Turner design Spex type 1670 Minimate monochromator was used to select the required wavelength. The light source is focussed onto the entrance slit of the monochromator with a quartz lens. An adjustable diaphragm shutter, mounted on the lens, enables the light beam to be blocked between measurements to minimize any photolytic side reactions. The transmitted light from the observation cell is detected by a photomultiplier (RCA type lP28). All optical detection components are mounted on an optical rail (Oriel), with the light source, quartz lens, entrance and exit slits of the monochromator and the photomultiplier at the same
2319
height as the observation cell of the stopped-flow apparatus. The difference between the output voltage from the photomultiplier and an adjustable constant voltage source is amplified (Analog Devices AD521 J) to maximize the reaction signal change in the range -2.5 to 2.5 V (the bipolar limits of the analogue to digital converter). A passive RC type circuit is used to remove high-frequency noise. The filtered signal is simultaneously displayed by oscilloscope and sampled by the AR1 1 analogue subsystem of a PDP-1 1/10-S minicomputer (or an M6802 microcomputer) and stored for subsequent analysis. Each stopped-flow run consisted of collecting 200 voltage/time data pairs at a specified sampling rate, waiting a predetermined interval and then collecting a further 20 data pairs, representing reaction equilibrium. The sampling rate was usually chosen so that 200 data pairs covered 95% of the total reaction amplitude. The collection and preliminary analysis of stopped-flow data were performed under the control of the BASIC program RDATA, developed within this laboratory. This program allowed the collection of up to 20 runs for each set of reactant solutions with the capability to maximize the signalto-noise ratio. Pseudo first-order rate constants for each run and the signal-average could be calculated by either a simple integrated rate expression or by the Guggenheim method. ’ 3 RESULTS Characterization
of [CO(NH~)~MOO~](C~O~)
IR and Raman spectra were used to ascertain the coordination mode of the molybdate species in [Co(NH,),MoO,](ClO,). Table 1 lists the principal observed IR and Raman frequencies for this compound and for [Co(NH,),MoO,]Cl, [Co(NH,), Mo0,]N03, [CO(NH~)~CI]MOO~and free Mo04’-, reported by Coomber and Griffith. I4 Fundamental Raman frequencies for Mo04’and C104-, measured from solid samples of Na*MoO, and NaClO,, are also included. The v1 symmetric stretch of the MoO,~- moiety can be used to discriminate between alternate modes of coordination. The observed v I frequencies for [CO(NH~)~MOO~](C~O~), 912 (IR) and 909 cm- ’ (R), are in excellent agreement with the values cited for [Co(NH3),Mo04]CI, thus indicating monodentate, inner-sphere coordination to the cobalt(II1) ion. Not surprisingly, the frequencies of the perchlorate anion are similar in the cobalt(II1) complex and in NaClO,, indicating no significant
M. R. GRACE and P. A. TREGLOAN
9
f
II%1
I
I3
I
r
I,&
I
I
z
I
interaction between the cobalt(II1) and Clodin the solid. The W-vis spectrum of [CO(NH~)~MOO~] (Clod) was recorded by diffuse reflectance from the solid to avoid complications arising from concurrent rapid aquation to Mo04*- and CO(NH~)~ 0HZ3+ when the complex is dissolved in solution. The spectrum has absorbance maxima at 360 and 525 nm, corresponding to the ‘T% + ‘A ,Jv2) and ‘T lg t ‘A ,&v ,) electronic transitions, respectively. The effective magnetic moment, peff, of [Co (NH,),MoO,](ClOJ at 298 K was determined as 0.54 B.M. This value is consistent with a wide variety of low-spin cobalt(II1) compounds with moments ranging from 0 to 1 B.M., “*I6 including the analogous monodentate chromate complex, [Co(NH 3)5 CrO,](ClO,), which has a peffof 0.62 B.M.’ As was observed with the chromate species, the non-zero magnetic moment for the d6 cobalt(II1) complex can be explained by temperature-independent paramagnetism, where there is a coupling of the ground and excited states under the influence of the applied magnetic field. This contention is supported by the invariance of the molar susceptibility (xhl = 1.5550.05 x lo-’ m3 mol-‘) over the temperature range 210-350 K. The molybdate moiety would also most probably be contributing to the observed A~, since applying diamagnetic corrections to the uncorrected susceptibilities for Na,MoO, and K2Mo04 yield room temperature, magnetic moments of 0.33 and 0.34 B.M., respectively (again arising from temperature-independent paramagnetism). ’ ’ Equilibrium measurements
As noted by Taylor in his original determination of K,, the equilibrium constant describing the complexation given in eq. (l), mixing solutions of Co (NH&OHz3+ and molybdenum(V1) under slightly basic conditions results in a rapid intensification of the purple colour. Figure 1 shows spectra from a mixing cell experiment illustrating the total absorbance change resulting from the complexation reaction. Taylor attributed this absorbance enhancement to the formation of CO(NH~)~MOO~+. This proposal is supported by the close similarity in absorbance maxima between the aqueous spectrum (Lm,, = 365 and 525 nm) and the diffuse reflectance measurements on [CO(NH~)~MOO.J(C~O,J in this study (&,,, = 360 and 525 nm). The evaluation of K, involved analysis of the absorbance changes at 540, 550 and 560 nm (corresponding to the wavelengths of maximum change) using the Benesi-Hildebrand method ’ ’
Formation
2321
of cobalt(II1) molybdate cation
Table 2. Experimental parameters and observed firstorder rate constants for the CO(NH~)~OH*+/MOO,~complexation reaction at 25°C and 1.O M ionic strength
Absorbance
PH
wavelength
0.05 0.10 0.075 0.025 0.01 0.0375 0.0175
17.1 53.8 35.1 6.39 1.94 11.5 3.77
0.7 4.6 1.8 0.31 0.08 0.3 0.19
7.73
0.05 0.10 0.075 0.025 0.01 0.0375 0.0175
3.61 9.84 6.15 1.44 0.49 2.41 0.93
0.12 0.28 0.19 0.07 0.05 0.07 0.03
7.43
0.05 0.10 0.075 0.025 0.01 0.0375 0.0175
7.03 21.8 13.94 2.93 1.02 4.78 1.93
0.33 1.2 0.50 0.17 0.05 0.21 0.07
7.92
0.05 0.10 0.075 0.025 0.01 0.0375 0.0175
1.96 5.17 3.33 0.87 0.3 1.37 0.57
0.09 0.20 0.15 0.07 0.05 0.10 0.05
Fig. 1. The effect of complexation upon the UV-vis spectrum of Co(NHJSOH2+ and Mo04’- at 25°C and 1.0 M ionic strength.
with the molybdate
Kinetic measurements The acid dependence of the reaction was measured over the pH range 7.1-8-O. Lower pH values and higher molybdate concentrations were avoided to minimize formation of isopolymolybdate species. l8 All reaction transients were first-order for at least five half-lives. The observed rate constants, kobs,and the relevant reaction conditions are listed in Table 2. Data analysis involved the determination of the forward rate constant, kf, and its resolution into the four terms given in eq. (2). Plots of kf vs [Mo04’-] were linear at all pH values indicating a second-order dependence upon [Mood*-1. The slope and intercept terms, shown in Table 3, correspond to the molybdate-dependent and -independent pathways in the forward rate law. Plotting these slope and intercept terms against p+] yields the four individual rate constants in Taylor’s rate expression. ’ The resultant values for k,, kb, k, and kd are included in Table 3. The agreement between the two studies is very reasonable if the constants are recalculated using molybdenum(V1) stock concentrations, rather than those at the “instant of mixing”.
Error
7.17
(nm)
concentration in large excess. Calculations were performed according to Taylor and resulted in an average value for K, of 960 + 80 M- ’ at 25°C and 1.O M ionic strength (cf. 475 + 30 M-l).’
[Mo04’-] (M)
All measurements were made at 25.O+O.l”C, 1.0 M ionic strength (NaClO,) and with a total cobalt(II1) concentration of 0.002 M. Solutions were buffered with 0.1 M triethanolamine. Quoted errors are one standard deviation and are derived from the signal-averaged reaction transient. Concentrations are at the “instant of mixing”.
Catalysis of C~(NH~)~MoO,+jiirmation As mentioned in the introductory remarks, it is possible that the observed second-order dependence upon [Mood*-] may be due to general acid catalysis of complex formation, presumably by HMoO,-. In order to test this hypothesis, various potential catalysts (Na2HP04, NaHCO, and Na2S04) were added to the molybdate solution prior to the stopped-flow run. The potential catalyst was mixed with the molybdate solution since Sod*-, CO,*+ and POd3- all form complexes with CO(NH~)~OH’+ in alkaline
2322
M. R. GRACE and P. A. TREGLOAN Table 3. Evaluation of second- and third-order rate constants for the formation of Co(NH&MoO,+ (a) Results of the kr versus [MOO,‘-] plots
PH
Slope
Intercept
Correction coefficient
7.17 7.43 7.73 7.92
(3.42kO.13) x lo4 (2.02 f0.08) x lo4 (1.65kO.05) x lo4 (1.20+0.06)x lo4
467 f 37 379f25 233+ 19 161+26
0.9960 0.9954 0.9941 0.9949
(b) Results of plots of molybdate-dependent
Intercept
Slope
Correction coefficient “Calculated centrations.
and -independent paths versus [H+]
Molybdateindependent path
Molybdatedependent path
Reference
k, (M-’ SC’) 123k23 166+ 16 96+7
kc (M-‘s-‘) (8.52f0.60) x lo3 (2.26k0.19) x 10’ (2.2&0.2)x 103
This work This work” 1
kb (M-2 s-‘) (5.74& 0.73) x IO9 (3.02kO.47) x 10’ (1.1+0.2)x 109
(3.64kO.23) x 10” (8.12f0.68) x 10” (1.02_+0.05) x 10”
This work This work 1
0.9567 0.946 1
0.9730 0.9656
This work This work”
using stock [MoO,~-]
solution.19 Before assessing the effect of these anions upon the rate of CO(NH~)~MOO~+ formation, some preliminary experiments, designed to determine any direct reaction between these anions
and MoOd2-, were performed. This precaution was deemed essential, since the ability of molybdenum(V1) to form heteropolymolybdate species, including heterophosphates, in acidic solution is well known. 2oMurata and Ikeda have used Raman spectroscopy to follow the formation of the 12molybdophosphate species, PMo , 20403-. 21 These brief investigations involved quantitatively monitoring the intensity of the Raman vi symmetric stretch of MOODY-at 897 cm-‘, upon the addition of increasing concentrations of the proposed catalytic species. The formation of any molybdenum(VI)-anion complex would result in a loss of Mo04’- tetrahedral symmetry and a consequent diminution in the v, peak intensity. In all cases, the concentration of free molybdate was greater than 99% of total molybdenum(W), even with “catalyst” concentrations in lo-fold excess. A realistic estimate for the accuracy of this technique is + 1 %, thus trace amounts of a molybdo-anion complex would not be detected. Pettersson” determined the
kd (M-3
s- ‘)
rather than “instant
of mixing” con-
speciation of molybdophosphates at 25°C over the pH range 3-9. At low pH, the total molybdenum(W) concentration was comprised almost solely of complex molybdenum-phosphate species, but above pH 7, only MoOa2- was detected. The combined results of these two studies suggest that the concentration of possible molybdophosphate species at pHs higher than 7.2 would be negligibly small. (a) The ej2ct of SOd2- addifion. The effect of S042- upon the rate of CO(NH~)~MOO,+ formation was assessed by adding Na2S04 to a 0.10 M MOODY- stock solution prior to mixing with 0.0040 M cobalt(II1) at pH 7.54. The results of these experiments are summarized in Table 4(a) and show that SOJ2- has only a marginal effect upon kobs. (b) The e&ct of HP04 2- addition. In view of the small rate enhancement caused by the addition of sodium sulphate to the reactant system, a preliminary stopped-flow experiment was performed, under identical solution conditions, using 0.025 M Na2HP04. The resultant reaction was too fast for accurate detection using the stopped-flow technique. The concentration of HP042- was then reduced by a factor of 50 and the reaction transient
2323
Formation of cobalt(II1) molybdate cation Table 4. General catalysis of CO(NH~)~MOO.,+ formation at 25°C and /L= 1.0M (a) Catalysis by S04*-
[Co”‘]
[Mo04’-]
PH
(M)
(M)
7.54
0.0020
0.050
[SO‘$-]
($?)
(M) 0 0.0025 0.025 0.100
5.05+0.13 5.57f0.22 6.22 f0.22 6.59f0.45
(b) Catalysis by HPO,*-
[Co”‘]
[Mood*-]
PH
(M)
(M)
7.35
0.0020
0.050
[HPO‘,*-] x lo3 (M) 0 0.25 0.50 0.75 1.00
11.5kO.6 18.1f1.2 21.3kO.9 27.7 f 2.5 31.3+ 1.5
7.54
0 0.50 1.00 1.75 2.50 25.0
5.01 kO.18 11.6kO.3 18.7* 1.0 25.3 + 1.2 37.3k3.1 Too fast
7.84
0 0.50 1.oo 1.75 2.50
4.33 f0.23 7.28 kO.26 11.5kO.8 15.9f0.7 20.7+ 1.8
(c) The effect of [NaHCO,] upon the rate of complex formation
[Co”‘]
[MoOa*-]
PH
(M)
(M)
7.25
0.0020
0.025
[HC03-] x lo3 (M) 0 1.25 2.50 3.75 5.00
5.62f0.06 7.96kO.12 10.3 +0.4 12.6kO.5 14.6f0.5
7.50
0 1.25 2.50 3.75 5.00
2.53 f 0.06 4.67kO.14 6.8 f0.2 8.6kO.2 10.9 f0.4
7.81
0
1.38kO.05 5.4f0.2 10.1 f0.9 14.4f0.5 19.3f1.5
3.13 6.25 9.38 12.5
LIRefers to the added concentrations of Na2S04, Na2HP04 and NaHCO, at the “instant of mixing”.
2324
M. R. GRACE
and P. A. TREGLOAN
2cQlm.o -
1.Ym.o
-
12Lxa.o
I 0.001
0.0015
0.032
O.WZ3
O.W3
6.0
Added[Na,HPO,l,i, 0.4)
Fig. 2. The effect of pH and total phosphate concentration on the rate of CO(NH,)~OH*+/MOO~*~ complexation at 25°C and 1.0 M ionic strength. [Co”‘] = 0.0020 M. [Mo04’-] = 0.050 M.
Fig. 3. The catalytic activity of H,PO,- on the rate of CO(NH,),OH~+/MOO~~ complexation at 25°C and 1.0 M ionic strength. [Co”*]= 0.0020 M. [Mo04*-] = 0.050 M.
even at 5 x lop4 M, this species doubled the rate of Co(NH&Mo04+ formation. Further stopped-flow runs were then conducted to determine the dependence of this rate enhancement upon the concentration of added Na2HP04 and the solution pH. Solution conditions and the observed first-order rate constants for these experiments are listed in Table 4(b). Figure 2 shows that rate enhancement is linearly dependent upon Na2HP04 at a given pH, but also that the degree of enhancement is acid dependent. The slope of the plot at each pH is a measure of “catalytic activity”. The results of the linear analysis of these plots are presented in Table 5. In the pH range 7-8, phosphate solutions contain both HP04’- and H2P04-. The relative proportions of these species are determined by the deprotonation equilibria : showed
that
H,PO,-
L ---HP04’-
+H+.
A value of 6.4OkO.06 was determined for p& at 250°C and p = 1.OM (NaClO,), by potentiometric
8.0 10.0 IZ.0 ‘k of total Phosphate as H2P0,-
titration with NaOH in CO,-free water, using the automatic titration equipment described elsewhere.23 This value is in good agreement with the 6.46 + 0.02 quoted by Smith and Martell under the same conditions.24 Figure 3 shows a plot of the catalytic activity vs the percentage of total phosphate as H2PO4-. It is apparent from this plot that rate enhancement is linearly dependent upon [HzPO4-] and also, that this species is the only detectable catalyst in solution. Complications arising from anation reactions between the cobalt(II1) and phosphate species in the stopped-flow chamber were not expected since these processes occur by slow cobalt-oxygen bond fission.” Hence, Co(NH3),Mo04+ formation would be essentially complete before any cobalt(II1) phosphate complexation occurs. (c) The eJ?ixt of HC03- addition. The addition reaction of CO2 to CO(NH~)~OH’+ occurs over the same time frame as the molybdate complexation processz6 and hence poses a potential interference to the desired reaction sequence (due to competition
Table 5. Determination of the pH dependence of Na,HPO, catalytic efficiency on the rate of CO(NH,),OH~~+/MOO,*- complexation at 25°C and 1.O M ionic strength
IX+1
% of Na,HPO,
PH
x 108
as H,PO,-
7.35 7.54 7.84
4.47 2.88 1.45
10.30 6.90 3.55
Slope (catalytic activity)
Intercept
Correction coefficient
19900+1400 12600+420 6580 f 340
11.7f0.6 5.1 fO.l 4.2kO.2
0.9956 0.9978 0.9976
Formation
of cobalt(II1) molybdate cation
2325
Table 6. Determination of the pH dependence of NaHC03 catalytic efficiency on the rate of CO(NHJ~OH~~+/MOO~~- complexation at 25°C and 1.O M ionic strength
P+l
% of carbonate
PH
x 108
as H&O3
7.25 7.50 7.81
5.62 3.16 1.55
1.15 x 10-2 6.62 x lo- 3 3.26 x lo-’
% of carbonate as HCO,- ’
Slope
Intercept
Correction coefficient
94.47 96.27 96.88
1830+62 1655f45 1363+37
5.63 +0.06 2.54kO.06 1.37f0.09
0.9998 0.9996 0.9994
between MOO,*- and CO2 for the cobalt(II1) ions).
However, the nature of the Co”‘/Mo042- complexation reaction and a careful choice of experimental parameters enables some very informative, semi-quantitative results to be obtained. Taylor found that with [MO”‘] in large excess, the rate of Co(NH 3)SMoO_,+ formation was independent of [Co”‘] and observations during preliminary stoppedflow runs in this work indicated that the presence of NaHC03 had no obvious effect on the reaction transient for this process. The transients were firstorder for at least five half-lives and also showed no diminution in absorbance change during reaction upon increased bicarbonate concentration. There was also no indication of a second, first-order transient corresponding to Co(NH 3)&O 3+ formation. By limiting the lowest pH solutions to 7.25, the amount of dissolved CO2 was kept below 5% of the total carbonate concentration and the ratio of Mo04’- to CO2 was never less than 2000. It was found that the presence of NaHCO, in solution during stopped-flow runs substantially increased the rate of CO(NH~)~MOO~+ formation. Solutions were prepared from CO,-free water in a nitrogen-filled glove-box. The results of the catalysis experiments are summarized in Table 4(c). The catalytic activity of NaHCO, was determined in an identical manner to the phosphate system and the linear fit parameters to the kobsvs [NaHCO,] plots are given in Table 6. The determination of the active catalytic species is complicated by extensive carbonate speciation and potential reaction with CO(NH~)~OH*+. Over the pH range studied, CO3*-, HC03-, H&O3 and dissolved CO2 are all present in solution. The relationship between these species and the cobalt(II1) reaction products is shown in the following equilibrium scheme : H20+CO&H2C03&
The deprotonation equilibrium involving HCO,and CO3*- has been well-characterized with a pK3 value of 9.57 at 25°C and p = 1.0 M.24 Dissolved CO2 and H2C03 are in an equilibrium heavily favouring CO3, to the extent that the equilibrium constants K, and K2 are normally combined into one apparent pK,,, for the “deprotonation of CO2”26 (pKapp= 6.02 at 25°C and p = 1.0 M24). A limited amount of work has been done on resolving Kapp into its K, and K2 components. van Eldik and Palmer report values for these constants of 2.28 x 10m3and 4.6 x lop4 M, respectively in 0.5 M NaCl.*’ In the absence of values at 1.0 M NaClO,, these results were used, in combination with K3, to calculate preliminary carbonate speciation in the molybdate solutions, at the pH values used in this study. The close agreement between the resultant Kapp value at p = 0.5 M NaC126 and the 1.0 M NaClO, value indicates that the error introduced by this procedure would be negligible in comparison to the inherent errors in the measurement of the observed rate constants and the effects of cobalt(III)-carbonato complexation. The calculated percentages of total carbonate as H2CO3 and HC03- at the “instant of mixing” are included in Table 6. The plot of catalytic activity against apparent percentage HzCO3, shown in Fig. 4, demonstrates that although the doubly protonated species appears to be the major catalyst for the formation of CO(NH~)~MOO~+, ‘the large positive intercept term indicates that anot.her species, most probably HC03-, is also responsible for some rate-enhancing effect. As there is considerable doubt as to the actual concentrations of these species, no further attempt was made to accurately quantify these effects.
H+ +HCO,-
K1’H+
KE
CO(NH~)~OH*+ 11
CO(NH~)~HCO~ ‘+ &Co(NH3)$03+
+H+
+CO,*-
M. R. GRACE and P. A. TREGLOAN
2326 2om.c
1ml.~
1600.(
MOOS
1X0.(
L
1cao.o 0.0
0.002
O.oW
0.026
The catalytic species, HA, may be any moiety in solution capable of proton donation to the molybdate to assist in water elimination. In this scheme, the reaction between cobalt@) and MoOd2-, observed by Taylor, is catalysed by HMo04- in solution. Although the concentration of HMoO,constitutes less than 0.1% of the total molybdate, it is still present at levels 100-1000 times higher than /H+]. The added catalysts H2P04-, H2C03 and possibly HC03- contribute to the total solution “pool” of proton-donating species. If the mechanism involves rate-determining proton-transfer and water departure, then the degree of rate enhancement should reflect the acidity of the proton which migrates to the molybdenum(W) hydroxy group. The slope of the catalytic activity plot is a direct measure of the effectiveness of that species to increase the rate of CO(NH~)~MOO~+ formation. The slope of the plot for H2C03, (6.0 kO.9) x 104, is much greater than that for H2P04-, (1.9 kO.1) x 103, which can be attributed to the greater acidity of the carbonate species. The slope estimate for H&JO3 is a lower limit since no account was taken of the effect of CO(NH~)~CO~+ formation upon carbonate speciation. Clearly, the existence of any correlation between the pK, of the proton-donating anion and the catalytic efficiency of the species, based only upon two anions, must remain speculative. Further investigation using a wide variety of substitution-inert, proton-donating species may help clarify this situation. It might also be expected that a second-order dependence upon cobalt(II1) concentration would be observed with this species in excess over [MOODY-]. In addition, CO(NH&OH~~+ would be the likely catalyst due to the greater acidity of the aquo ligand’s proton relative to the CO(NH~)~OH~+ proton. Another advantage of using excess .cobalt(III) is that the complications arising from HMoO,catalysis could be minimized. Path (b) in Scheme 1 involves the initial formation of the protonated complex, CO(NH~)~ HMoO, 2+, followed by rapid deprotonation to CO(NH~)~MOO~+. It has been demonstrated that the incorporation of a CO(NH~)~ group into the coordination sphere of various 0x0 anions
J 1.011
o.txm 0.01 , ck of total Carbonate as H2CC ‘3
Fig. 4. The catalytic activity of H&O3 on the rate of CO(NH,),OH~+/MOO,~- complexation at 25°C and 1.0 M ionic strength. [Co”‘] = 0.0020 M. [Mo04*-] = 0.050 M.
DISCUSSION
The reinvestigation of the pH and molybdate dependence for the formation of Co(NH3)5Mo0,+, resulted in an expression for the forward rate term which corresponded exactly with that proposed by Taylor2 [eq. (2)]. After allowing for a plausible difference, by a factor of two, in evaluating the MoOd2- concentration (stock or “instant of mixing”), the individual second- and third-order rate constants comprising this term were also in good agreement. However, after consideration of the catalytic effect of several protonated anions, an alternative reaction scheme can be proposed to account for the apparent second-order dependence upon [MoO,~-], observed in both studies. Scheme 1, shown below, involves direct reaction between the cobalt(II1) and MoOd2- species and also requires proton donation from an external “acid catalyst”. The resultant rate expression from this scheme also conforms to the observed pH and concentration dependences expressed in eq. (2).
(a) CO@JH,)~OH,~++MOO,~- +CO(NH~)~MOO~++H~O, (h) CO(NHJ@H,~+ +HMoO,- --‘Co(NH&Mo04+ (c) Co(NH,),0HZ++Mo0,2-
+H@+H+,
+HA=Y+-CO(NH~),MOO~++H~O+A-,
(d) Co(NH,),0HZ3+ +MoO,‘- +HA --‘Co(NHJ5Mo04++H20+HA. Scheme 1. Various possible reaction paths for the formation cf c~(NI-I~)~M~G,+.
Formation
t 1 2321
of cobalt(II1) molybdate cation
(AsOd3-, POd3-, SeOs3-, Cr04’-) dramatically enhances the acidity of any ionizable protons on that anion. 28-31It is therefore assumed that the pK, of CO(NH~)~HMOO‘,~+ will be significantly lower than HMoO,- itself (3.92 at 25°C and fi = 1.O M’) and consequently, at pH values above 7, the concentration of the protonated product would be negligible. The almost complete indifference of the reaction rate to added SOd2- suggests that the catalysis involves proton transfer. It may be argued that the very small dependence of kobsupon increasing sulphate concentration Fable 4(a)] is due to trace amounts of HS04-, but it is equally plausible that this dependence is simply a medium effect resulting from the relatively high concentrations of added Na2S04 (up to 0.1 M). Taylor noted that changing the electrolyte from NaClO, to LiC104 caused a marked decrease in kobs. It is interesting to note that the addition of Na,SO, did not retard the complexation reaction as was observed with the formation of Co(tetren)HS032+.22 This decrease in rate upon addition of SOd2- was attributed to the formation of unreactive {Co(tetren)OH23+ . . . SOd2-) ion pairs. Unfortunately, definitive assignment of reaction pathways to the terms in the rate expression is impossible due to the inherent proton ambiguities. Paths (a) and (d) in Scheme 1 may be rewritten as reactions involving CO(NH~)~OH’+ and HMo04-, as proposed by Taylor. It is believed that the reaction between Co-OH and HMoO,- is more likely than Co-OH, and MoOd2-, since the latter option requires transfer of both protons from the aquo ligand to a molybdenum(V1) oxygen, whereas only one transfer to the molybdenum(V1) hydroxo ligand is necessary from Co(NH3),0H2+. Conversely, significant ion-pairing and internal hydrogen-bonding would favour the + 3/- 2 reaction. Paths (a) and (b) may also be expressed in terms of external proton donation or catalysis as shown below : (e) CO(NH~)~OH~+ +MoOd2- +H+ + CO(NH~)~MOO~+ +H20 (f) Co(NH3),0HZ3+ +MoOA2- +H+ e CO(NH,)~MOO‘,+ +H20+H+. The attractive feature of this last proposal is that paths (e) and (f) are identical to the general catalysis paths (c) and (d) (except that HA is now specifically H+) and therefore, only two reactions are required to describe the entire complexation process. Figure 5 depicts possible transition states for the various pathways proposed for the complexation
0 CbW&
cd+ + H?aoo;
ii) CcWL&+
+
IMoO;
Fig. 5. Proposed
transition states for molybdate complexation.
cobalt(II1)
process. The HMoO,- anion has been assigned as the reactive molybdenum(W) species in the uncatalysed reactions on the basis of the proton-transfer argument discussed above, and also to allow direct comparison with other molybdate substitution processes which are all expressed in terms of this protonated species. The specific proton-catalysed reaction pathways are incorporated into the general acid terms (iii) and (iv), where HA = H+. The rate constants for the formation of transition states (iF (iv) are given by : (i) kl = k,xK,/K,, (ii) k2 = kblKHM (iii) k3 = k, x KH (iv) k4 = kd
=6.6x 104M-Is-’ = 1.3 x 106M-‘s-’ = 2.1 x 10’“M-3s-’ = 3.6x lO”M-3s-‘*
9
where &, and KHM are the acidity constants for CO(NH~)~OH~~+ (4.79 x lo-’ at 25°C and p= 1.0 M”) and HMoO,-, respectively. The rate constants, k, and k2, for the uncatalysed reaction of HMoO,- with CO(NH~)~OH~+ and CO(NH~)~OH~~+, respectively, are much lower than for a series of bidentate ligands, shown in Table 7, where the second-order rate constants falls in the range (0.42.0) x lo* M-’ SK’. These faster reactions have all been attributed to ligand addition processes where the molybdenum(W) centre expands its coordination to six and Mo-0 bond cleavage is unnecessary. Although the reaction of HMoO,- with H2EDTA2- forms an octahedral molybdenum(V1) complex, the slow rate of reaction results from the necessity to break one Mo-0 bond.
M. R. GRACE
2328
and P. A. TREGLOAN
Table 7. Second-order rate constants for complexation with HMoO,Ligand (L) Co(NH&OH’+ Co(NH&OHz3+ H2EDTA2Oxine(Oxine)SO,‘Hcatechol-
Rate constant (M- ’ s- ‘) 6.6 x 1.3 x 2.3 x 1.5 x 4.0 x 1.9 x
104 106 10’ 108 10’ 108
at 25°C”
Product
Denticity number of L
Reference
Mo03L+ MoOSL+ Mo03L3’Mo02(0H),L,MoO~(OH),L~~~ MoO~(OH)~L~*-
1 1 3 2 2 2
This work This work 33 34 35 36
a Adapted from ref. 1.
A comparison of the rate constants for the two uncatalysed processes (k, and k2), reveals that complexation occurs more quickly with Co(NH,), OHz3+ than the corr e sp onding hydroxo species. The importance of ion-pairing in this reaction scheme was dismissed in the previous study, ’ by analogy with the weak outer-sphere interaction between CO(NH~)~~+ and MOO,‘-. However, the formation of hydrogen-bonding contacts between the anions HSeO,or HCr04and CO(NH~)~ 0Hz3+ resulted in values for the ion-pairing constant, K,,, of between 6 and 8 M-‘.3’*37 Even allowing for an order of magnitude difference between K,, values for the + 3/ - 1 and + 2/ - 1 ionpairs, the intrinsic rate constant for the reaction with the aquo species is still twice that for the cobalt(II1) hydroxo ligand. Thus, the effect of increased hydrogen bonding in the transition state and greater proton acidity for reaction with Co slightly outweighs the higher nucleo(NH3)&Hz3+ philicity of the hydroxo ligand in CO(NH~)~OH~+, in determining the more reactive cobalt(II1) species towards HMoO,-. Even with the absence of supporting activation parameters for this reaction, it still seems reasonable that, in view of the associative nature of the Mo042--oxygen exchange reaction3* the extent of bond formation in the complexation transition state outweighs bond breaking. Consequently, changing the nature of the inert cation substituting into the molybdenum coordination sphere should affect the rate of complex formation. CONCLUSION Whilst the form of the rate law for the complexation reaction between Mo04*- and aquo- and hydroxo-pentamminecobalt(II1) cations in basic solution is consistent with a previous proposal for the existence of dimeric molybdenum(V1) species, general anion catalysis experiments show that an
alternative mechanism is likely. The observed rate of reaction is enhanced by anionic species capable of proton donation to the molybdenum(V1) hydroxo ligand, thereby facilitating water elimination from the transition state. The apparent second-order dependence upon [Mood*-] may be explained by the involvement of HMoO,in proton donation. Consequently, this study found no definitive evidence for the formation of molybdenum dimers in slightly basic solution. Proton ambiguities preclude definitive identification of reactant species in aqueous solution. An investigation into the direct reaction of MoOd2with CO(NH~)~OH~‘+ or CO(NH~)~OH*+ in an aprotic solvent may prove beneficial in this regard. Acknowledgement-M.R.G. Government for the assistance graduate Research Award.
is grateful to the Australian of a Commonwealth Post-
REFERENCES R. S. Taylor, Znorg. Chem. 1977, 16, 116. V. W. Day, M. F. Fredrich, W. G. Klemperer and W. J. Shum, J. Am. Chem. Sot. 1977,99,6146. J. J. Cruywagen and J. B. B. Heyns, Znorg. Chem. 1987,26, 2569. M. J. Taylor and J. M. Coddington, Conference Abstract M25, XXVII International Conference on Coordination Chemistry, Queensland, Australia (1989). 5. P. L. Brown, M. E. Shying and R. N. Sylva, J. Chem. Sot., Dalton Trans. 1987, 2149. 6. F. Basolo and R. K. Murmann, Znorg. Synth. 1953, 4, 171.
7. R. van Eldik and G. M. Harris, Znorg. Chem. 1980, 19, 880.
8. W. E. Jones and T. W. Swaddle, Can. J. Chem. 1967, 45, 2647.
9. B. F. Abrahams,
M. R. Grace, B. F. Hoskins and P. A. Tregloan, Zrzorg. Chem. Acta 1991,182, 135. 10. A. I. Vogel, Textbook of Quantitative Inorganic Analysis, 4th edn. Longmans, London (1978).
Formation
of cobalt(II1) molybdate cation
Il. E. K&rig, Magnetic Properties of Transition Metal Compounds, Vol. 2. Springer-Verlag, Berlin (1966). 12. Q. H. Gibson and L. Milnes, Biochem. J. 1964, 91, 161. 13. A. A. Frost and R. G. Pearson, Kinetics and Mechanism, 2nd edn. John Wiley and Sons, New York (1961). 14. R. Coomber and W. P. Griffith, .Z. Chem. Sot. A 1968,1128. 15. K. Nakamoto, Infrared and Raman Spectra of Znorganic and Coordination Compounds, 3rd edn. WileyInterscience, New York (1978). 16. E. K&rig and G. Konig, Magnetic Properties of Transition Metal Compoundc, Supplement 2. SpringerVerlag, Berlin (1979). 17. H. A. Benesi and J. H. Hildebrand, J. Am. Chem. Sot. 1949,71,2703. 18. C. F. Baes Jr and R. E. Mesmer, The HydroZysis of Cations. Wiley-Interscience, New York (1976). 19. W. G. Jackson, C. N. Hookey, M. L. Randall, P. Comba and A. M. Sargeson, Znorg. Chem. 1984,23, 2473. 20. For example: A. Yagasaki, I. Andersson and L. Pettersson, Znorg. Chem. 1987, 26, 3926. 21. K. Murata and S. Ikeda, Polyhedron 1987,6, 1681. 22. L. Pettersson, Acta Chem. Stand. 1971,25, 1958. 23. A. Campisi and P. A. Tregloan, Znorg. Chim. Acta 1985,100,251. 24. R. M. Smith and A. E. Martell, Critical Stability Constants, Vol. 4. Plenum Press, New York (1976).
2329
25. S. F. Lincoln and D. R. Stranks, Aust. J. Chem. 1968, 21, 37. 26. E. Chaffee, T. S. Dasgupta and G. M. Harris, J. Am. Chem. Sot. 1973,95,4169. 27. R. van Eldik and D. A. Palmer, J. Soln Chem. 1982, 11, 339. 28. Stability Constants of Metal-Zon Complexes, IUPAC
29. 30. 3 1. 32. 33. 34. 35. 36. 37. 38.
Chemical Data Series, No. 21. Pergamon Press, Oxford (1982). T. A. Beech, N. C. Lawrence and S. F. Lincoln, Aust. J. Chem. 1973,26, 1877. A. D. Fowless and D. R. Stranks, Znorg. Chem. 1977, 16, 1271. M. R. Grace and P. A. Tregloan, manuscript in preparation. A. C. Dash, A. A. El-Awady and G. M. Harris, Znorg. Chem. 1981, 20, 3160. D. S. Honig and K. Kustin, J. Am. Chem. Sot. 1973, 95, 5525. P. F. Knowles and H. Diebler, Trans. Faraday Sot. 1968,&l, 977. H. Diebler and R. E. Timms, J. Chem. Sot. A. 1971, 273. K. Kustin and S.-T. Liu, J. Am. Chem. Sot. 1976, 98,2487, A. D. Fowless and D. R. Stranks, Znorg. Chem. 1977, 16, 1276. H. Gamsjiiger and R. K. Murmann, Advances in Inorganic and Bioinorganic Mechanisms (Edited by A. G. Sykes), Vol. 2. Academic Press, London (1983).
0
0277-5387/91 $3.00+.00 1991 Pergamon Press plc
AN ALTERNATIVE MECHANISM FOR THE FORMATION OF THE COBALT(III) MOLYBDATE CATION, Co(NH,),MoO,+ MICHAEL
R. GRACE*? and PETER A. TREGLOAN
Inorganic Chemistry Section, School of Chemistry, University of Melbourne, Parkville, Victoria 3052, Australia (Received 7 March 199 1; accepted 5 June 1991)
Abstract-An alternative mechanism is proposed for the formation of CO(NH~)~MOO~+ from CO(NH~)~H~O~+ and MoOd2- in slightly basic solution (pH 7.1-8.0). The observed second-order rate dependence upon [MOODY-]may be explained by HMoO,- acting as a proton-donating catalyst, in contrast to the previously postulated mechanism involving the formation of dimeric molybdenum(W) species. An increase in reaction rate was also caused by the addition of other protonated anions (HC03-, H,PO,-) to the reaction solution, thus supporting the contention of a general proton-assisted mechanism. No evidence of dimeric molybdenum(V1) was found.
Taylor has reported that the formation of Co (NH3)5M004+ from MoOd2- and CO(NH~)~OH*+, in the pH range 7.1-8.0, could be conveniently followed using the stopped-flow technique.’ It was also found that the complexation reaction, given in eq. (l), showed a greater than first-order dependence upon [MOODY-].
Co(NH&Mo04+
+H20
(1)
The forward rate constant, kf, has four terms, according to eq. (2) : kf = k,+k,[H+]+k,[Mo0,2-]+kJH+][Mo0,2-]; (2)
where: k, = 96+7 M-’ s-l, kb = (1.1 f0.2) x 10’ Me2 s-l, k, = (2.2+0.2)x lo3 M-2 s-l and kd = (1.02f0.05) x lo” Me3 SK’ at 25°C and p = 1.0 M (NaClOJ. Taylor also noted that the rate of reaction was consistent with substitution at the molybdenum(V1) rather than cleavage of the inert cobalt(III)-oxygen bond. It was suggested that the pathways k, and kb correspond to the reaction of HMoO,with respecCO(NH~)~OH~~+ and Co(NH3),0H2+, * Author to whom correspondence should be addressed. t Present address : Chemistry Department, University of Calgary, Calgary, Alberta, Canada T2N lN4.
tively. The second-order [MoO,~-] dependence for paths k, and kd was ascribed to a reaction involving the dimolybdate species Mo20g4-, HMo~O~~- or H2M020g2-, depending upon the active cobalt(II1) entity, in order to conform to the observed acid dependence. The dimolybdate anion, M0207~-, has been isolated from acetonitrile solution as the (n-C4Hg)4N+ salt, containing corner-shared MOODY- tetrahedra.* However, Mo20T2- has never been detected in aqueous solution, although a recent study by Cruywagen and Heyns3 reported evidence of the protonated species HMo207- under mildly acidic conditions (2.3 < pH < 6.1). The discrete dimeric species, Mo~O~(H~O)~~+, has been identified as the major component of molybdate solutions in 2-6 M HC104 and HN03 by g5Mo NMR and Raman spectroscopy.4 The structure involves two molybdenum(W) centres bridged by a single oxygen. Two terminal oxygens and three aquo ligands constitute an octahedral environment around each molybdenum ion. However, the existence of an aqueous dimeric molybdenum(V1) species, under normal conditions, must remain tentative since another recent potentiometric study’ found no molybdate oligomers between monomers and the heptameric series Mo~O~~H~(~-“- (n = 0, 1, 2 and 3). The major objective of the present study was to reinvestigate the complexation between MOODY-
2317
2318
M. R. GRACE and P. A. TREGLOAN
and CO(NH~)~OH~~+ in basic solution, in order to investigate Taylor’s postulated reaction between cobalt(II1) ions and dimeric molybdenum(V1) species. The presence (or absence) of a term involving dimolybdate in the rate expression could then be used as independent evidence for its existence and assist in the continuing development of the models used to describe molybdenum(V1) speciation in acid solution. We thought that the second order term in [MOODY-]might also arise from HMoO,- acting as a general acid catalyst. Consequently, the effect of a range of anions upon the rate of the complexation reaction was also examined. EXPERIMENTAL Materials
[Co(NH 3)SOH 2](ClO4)3 was prepared by established methods.677 W-vis spectra were in good agreement with those reported previously.7*8 Na,MoO, - 2H20 (May & Baker, LR) was twice recrystallized from base solution (pH x 11). Triethylanolamine (Ajax, LR) was distilled in vacua. Na2HP04 (Ajax, AR), NaHCO, (BDH, Analar), Na,SO, (BDH, Analar) and NaOH (Ajax, AR) were used without further purification. Twice distilled, once deionized water was used throughout Both equilibrium and kinetic measurements were performed in the presence of 0.1 M triethanolamine in order to buffer the solutions at the required pH and prevent the formation of polymeric molybdenum(V1) species. ’ Solutions were adjusted to the desired pH by addition of either 2 M HClO, or 2 M NaOH. These solutions showed negligible drift in pH (< 0.05 pH) units) over a period of several months. The synthesis of [Co(NH3),Mo04](C104) involved mixing slightly alkaline 0.05 M solutions (8 < pH < 9) of MoOd2- and CO(NH~)~OH~+ at 5&6O”C in the presence of ca 0.5 M NaClO,. The mixed solution was maintained at this temperature for a further 15 min. The pink solid formed upon cooling was filtered and then washed with ice-cold water, CH3CH20H and then (CH3CH2),0 and dried at 60°C for several hours. The crude complex was recrystallized from a minimum of hot (x 80°C) water at pH 8, filtered, washed and dried as above. Yields were typically 3040%. The analysis of NH, content has been described *The term “pH”, used throughout this paper, refers to the quantity -log,o[H+]. t All concentrations are at the “instant of .mixing” = stock concentration/2.
previously.’ For [Co(NH,),MoO,](ClO~) : Found : 20.6. Calc. : 21.1%. The residue after product decomposition was filtered and the molybdenum concentration in the filtrate was determined gravimetrically as Mo02(CvH60N), after complexation with 8-hydroxyquinoline.” Found : 23.5. Calc. : 23.8%. Physical methods
All pH* measurements were performed using a Radiometer GK240 1C combined glass electrode, in which the saturated KC1 reference solution was replaced with 1 M NaCl to prevent the precipitation of KC104 at the liquid junction. The electrode was calibrated daily with a series of acid (HC104, pH 2-3.5, p = 1.0 M) and base (NaOH, pH 11-12.5, p = 1.0 M) standards. The slope of the pH vs electrode potential plot was typically 0.059+0.002 V per pH unit. Absorbance measurements were performed on a Varian Techtron 635D UV-vis spectrophotometer, using a cell block thermostatted to 25.O+O.l”C. Spectra were recorded in a quartz mixing cell (Hellma, 238-QS) which is divided into two equal path-length compartments by a thin wall ; a hole at the top of this wall allows mixing of the two reactant solutions. Consequently any absorbance change upon mixing can be directly attributed to chemical reaction. IR spectra were recorded as KBr discs on a Jasco 302 IR spectrometer. Raman spectra were obtained using a Spex Ramalog 5 Raman spectrometer interfaced with an M6800 microcomputer. During Raman studies of molybdate/phosphate solutions, NaC104 was added to each solution to maintain an ionic strength of 1.OM and also provide an intensity calibration for each spectrum [at v,(C104-) = 935 cm- ‘I. HCl was used to obtain the desired solution PH. Diffuse reflectance spectra were recorded over the range 34&750 nm on a Beckman DK-2A spectrophotometer (BaS04 reference). Variable temperature magnetochemical measurements were performed over the range 21&350 K on a thermostatted Gouy balance, calibrated with Hg[Co(NCS),]. Diamagnetic corrections were applied according to Kiinig. ’ I Rate measurements The formation of CO(NH~)~MOO~+ from MoOd2- and CO(NH~)~OH~~+ was followed spectrophotometrically at 550 nm, using the stoppedflow technique. The cobalt(II1) concentration was maintained at 0.0020 M,t whilst [MOODY-] was
Formation
of cobalt(II1) molybdate cation
varied between 0.010 and 0.100 M. All rate measurements were made at an ionic strength of 1.O M (NaC103 and at 25.0 + 0.1 “C over the pH range 7.1-8.0. A stopped-flow apparatus, designed and constructed within the Chemistry Department at Melbourne University, was used to perform kinetic experiments with half-lives ranging from 10 ms to several minutes. The molybdate and cobalt(II1) reactant solutions were drawn into 2 cm3 glass, gastight drive syringes (Hamilton, 1002-TEFLL) from plastic reserve syringes via three-way taps (Hamilton, No. 86727). A kinetic experiment (“run”) was initiated by triggering a solenoidcontrolled pneumatic drive, which forced fresh reactant solutions through a Kel-F mixer” and into the observation chamber. Up to 20 runs can be recorded before refilling the drive syringes. The observation cell consists of either: (a) a “straight-through” 2 mm internal diameter quartz tube where the light-beam passes through the tube at a 90” angle, or (b) a “Z-shape” 18 mm pathlength chamber in a Kel-F block, enclosed into the brass stopped-flow block by quartz windows. Choice of cell is governed by the magnitude of the absorbance change and the intensity of the transmitted light. Optically dense solutions necessitated using the short path-length tube to ensure adequate signal-to-noise response of the photomultiplier detection system. The stopped-flow unit is encased in a brass block through which thermostatted water is circulated, allowing temperatures between 7 and 45°C to be maintained to within 0.05”C. A Hewlett Packard 34740A platinum resistance thermometer is inserted into a cavity in the brass casing to enable temperature resolution of 0.01 “C. All reactant solutions were thermostatted within the drive syringes for 20 min to ensure thermal equilibration. A 100 W quartz-halogen lamp, powered by a Hewlett Packard Model 6286A supply, provided the light source for the spectrophotometric detection of reaction progress. A Czerny-Turner design Spex type 1670 Minimate monochromator was used to select the required wavelength. The light source is focussed onto the entrance slit of the monochromator with a quartz lens. An adjustable diaphragm shutter, mounted on the lens, enables the light beam to be blocked between measurements to minimize any photolytic side reactions. The transmitted light from the observation cell is detected by a photomultiplier (RCA type lP28). All optical detection components are mounted on an optical rail (Oriel), with the light source, quartz lens, entrance and exit slits of the monochromator and the photomultiplier at the same
2319
height as the observation cell of the stopped-flow apparatus. The difference between the output voltage from the photomultiplier and an adjustable constant voltage source is amplified (Analog Devices AD521 J) to maximize the reaction signal change in the range -2.5 to 2.5 V (the bipolar limits of the analogue to digital converter). A passive RC type circuit is used to remove high-frequency noise. The filtered signal is simultaneously displayed by oscilloscope and sampled by the AR1 1 analogue subsystem of a PDP-1 1/10-S minicomputer (or an M6802 microcomputer) and stored for subsequent analysis. Each stopped-flow run consisted of collecting 200 voltage/time data pairs at a specified sampling rate, waiting a predetermined interval and then collecting a further 20 data pairs, representing reaction equilibrium. The sampling rate was usually chosen so that 200 data pairs covered 95% of the total reaction amplitude. The collection and preliminary analysis of stopped-flow data were performed under the control of the BASIC program RDATA, developed within this laboratory. This program allowed the collection of up to 20 runs for each set of reactant solutions with the capability to maximize the signalto-noise ratio. Pseudo first-order rate constants for each run and the signal-average could be calculated by either a simple integrated rate expression or by the Guggenheim method. ’ 3 RESULTS Characterization
of [CO(NH~)~MOO~](C~O~)
IR and Raman spectra were used to ascertain the coordination mode of the molybdate species in [Co(NH,),MoO,](ClO,). Table 1 lists the principal observed IR and Raman frequencies for this compound and for [Co(NH,),MoO,]Cl, [Co(NH,), Mo0,]N03, [CO(NH~)~CI]MOO~and free Mo04’-, reported by Coomber and Griffith. I4 Fundamental Raman frequencies for Mo04’and C104-, measured from solid samples of Na*MoO, and NaClO,, are also included. The v1 symmetric stretch of the MoO,~- moiety can be used to discriminate between alternate modes of coordination. The observed v I frequencies for [CO(NH~)~MOO~](C~O~), 912 (IR) and 909 cm- ’ (R), are in excellent agreement with the values cited for [Co(NH3),Mo04]CI, thus indicating monodentate, inner-sphere coordination to the cobalt(II1) ion. Not surprisingly, the frequencies of the perchlorate anion are similar in the cobalt(II1) complex and in NaClO,, indicating no significant
M. R. GRACE and P. A. TREGLOAN
9
f
II%1
I
I3
I
r
I,&
I
I
z
I
interaction between the cobalt(II1) and Clodin the solid. The W-vis spectrum of [CO(NH~)~MOO~] (Clod) was recorded by diffuse reflectance from the solid to avoid complications arising from concurrent rapid aquation to Mo04*- and CO(NH~)~ 0HZ3+ when the complex is dissolved in solution. The spectrum has absorbance maxima at 360 and 525 nm, corresponding to the ‘T% + ‘A ,Jv2) and ‘T lg t ‘A ,&v ,) electronic transitions, respectively. The effective magnetic moment, peff, of [Co (NH,),MoO,](ClOJ at 298 K was determined as 0.54 B.M. This value is consistent with a wide variety of low-spin cobalt(II1) compounds with moments ranging from 0 to 1 B.M., “*I6 including the analogous monodentate chromate complex, [Co(NH 3)5 CrO,](ClO,), which has a peffof 0.62 B.M.’ As was observed with the chromate species, the non-zero magnetic moment for the d6 cobalt(II1) complex can be explained by temperature-independent paramagnetism, where there is a coupling of the ground and excited states under the influence of the applied magnetic field. This contention is supported by the invariance of the molar susceptibility (xhl = 1.5550.05 x lo-’ m3 mol-‘) over the temperature range 210-350 K. The molybdate moiety would also most probably be contributing to the observed A~, since applying diamagnetic corrections to the uncorrected susceptibilities for Na,MoO, and K2Mo04 yield room temperature, magnetic moments of 0.33 and 0.34 B.M., respectively (again arising from temperature-independent paramagnetism). ’ ’ Equilibrium measurements
As noted by Taylor in his original determination of K,, the equilibrium constant describing the complexation given in eq. (l), mixing solutions of Co (NH&OHz3+ and molybdenum(V1) under slightly basic conditions results in a rapid intensification of the purple colour. Figure 1 shows spectra from a mixing cell experiment illustrating the total absorbance change resulting from the complexation reaction. Taylor attributed this absorbance enhancement to the formation of CO(NH~)~MOO~+. This proposal is supported by the close similarity in absorbance maxima between the aqueous spectrum (Lm,, = 365 and 525 nm) and the diffuse reflectance measurements on [CO(NH~)~MOO.J(C~O,J in this study (&,,, = 360 and 525 nm). The evaluation of K, involved analysis of the absorbance changes at 540, 550 and 560 nm (corresponding to the wavelengths of maximum change) using the Benesi-Hildebrand method ’ ’
Formation
2321
of cobalt(II1) molybdate cation
Table 2. Experimental parameters and observed firstorder rate constants for the CO(NH~)~OH*+/MOO,~complexation reaction at 25°C and 1.O M ionic strength
Absorbance
PH
wavelength
0.05 0.10 0.075 0.025 0.01 0.0375 0.0175
17.1 53.8 35.1 6.39 1.94 11.5 3.77
0.7 4.6 1.8 0.31 0.08 0.3 0.19
7.73
0.05 0.10 0.075 0.025 0.01 0.0375 0.0175
3.61 9.84 6.15 1.44 0.49 2.41 0.93
0.12 0.28 0.19 0.07 0.05 0.07 0.03
7.43
0.05 0.10 0.075 0.025 0.01 0.0375 0.0175
7.03 21.8 13.94 2.93 1.02 4.78 1.93
0.33 1.2 0.50 0.17 0.05 0.21 0.07
7.92
0.05 0.10 0.075 0.025 0.01 0.0375 0.0175
1.96 5.17 3.33 0.87 0.3 1.37 0.57
0.09 0.20 0.15 0.07 0.05 0.10 0.05
Fig. 1. The effect of complexation upon the UV-vis spectrum of Co(NHJSOH2+ and Mo04’- at 25°C and 1.0 M ionic strength.
with the molybdate
Kinetic measurements The acid dependence of the reaction was measured over the pH range 7.1-8-O. Lower pH values and higher molybdate concentrations were avoided to minimize formation of isopolymolybdate species. l8 All reaction transients were first-order for at least five half-lives. The observed rate constants, kobs,and the relevant reaction conditions are listed in Table 2. Data analysis involved the determination of the forward rate constant, kf, and its resolution into the four terms given in eq. (2). Plots of kf vs [Mo04’-] were linear at all pH values indicating a second-order dependence upon [Mood*-1. The slope and intercept terms, shown in Table 3, correspond to the molybdate-dependent and -independent pathways in the forward rate law. Plotting these slope and intercept terms against p+] yields the four individual rate constants in Taylor’s rate expression. ’ The resultant values for k,, kb, k, and kd are included in Table 3. The agreement between the two studies is very reasonable if the constants are recalculated using molybdenum(V1) stock concentrations, rather than those at the “instant of mixing”.
Error
7.17
(nm)
concentration in large excess. Calculations were performed according to Taylor and resulted in an average value for K, of 960 + 80 M- ’ at 25°C and 1.O M ionic strength (cf. 475 + 30 M-l).’
[Mo04’-] (M)
All measurements were made at 25.O+O.l”C, 1.0 M ionic strength (NaClO,) and with a total cobalt(II1) concentration of 0.002 M. Solutions were buffered with 0.1 M triethanolamine. Quoted errors are one standard deviation and are derived from the signal-averaged reaction transient. Concentrations are at the “instant of mixing”.
Catalysis of C~(NH~)~MoO,+jiirmation As mentioned in the introductory remarks, it is possible that the observed second-order dependence upon [Mood*-] may be due to general acid catalysis of complex formation, presumably by HMoO,-. In order to test this hypothesis, various potential catalysts (Na2HP04, NaHCO, and Na2S04) were added to the molybdate solution prior to the stopped-flow run. The potential catalyst was mixed with the molybdate solution since Sod*-, CO,*+ and POd3- all form complexes with CO(NH~)~OH’+ in alkaline
2322
M. R. GRACE and P. A. TREGLOAN Table 3. Evaluation of second- and third-order rate constants for the formation of Co(NH&MoO,+ (a) Results of the kr versus [MOO,‘-] plots
PH
Slope
Intercept
Correction coefficient
7.17 7.43 7.73 7.92
(3.42kO.13) x lo4 (2.02 f0.08) x lo4 (1.65kO.05) x lo4 (1.20+0.06)x lo4
467 f 37 379f25 233+ 19 161+26
0.9960 0.9954 0.9941 0.9949
(b) Results of plots of molybdate-dependent
Intercept
Slope
Correction coefficient “Calculated centrations.
and -independent paths versus [H+]
Molybdateindependent path
Molybdatedependent path
Reference
k, (M-’ SC’) 123k23 166+ 16 96+7
kc (M-‘s-‘) (8.52f0.60) x lo3 (2.26k0.19) x 10’ (2.2&0.2)x 103
This work This work” 1
kb (M-2 s-‘) (5.74& 0.73) x IO9 (3.02kO.47) x 10’ (1.1+0.2)x 109
(3.64kO.23) x 10” (8.12f0.68) x 10” (1.02_+0.05) x 10”
This work This work 1
0.9567 0.946 1
0.9730 0.9656
This work This work”
using stock [MoO,~-]
solution.19 Before assessing the effect of these anions upon the rate of CO(NH~)~MOO~+ formation, some preliminary experiments, designed to determine any direct reaction between these anions
and MoOd2-, were performed. This precaution was deemed essential, since the ability of molybdenum(V1) to form heteropolymolybdate species, including heterophosphates, in acidic solution is well known. 2oMurata and Ikeda have used Raman spectroscopy to follow the formation of the 12molybdophosphate species, PMo , 20403-. 21 These brief investigations involved quantitatively monitoring the intensity of the Raman vi symmetric stretch of MOODY-at 897 cm-‘, upon the addition of increasing concentrations of the proposed catalytic species. The formation of any molybdenum(VI)-anion complex would result in a loss of Mo04’- tetrahedral symmetry and a consequent diminution in the v, peak intensity. In all cases, the concentration of free molybdate was greater than 99% of total molybdenum(W), even with “catalyst” concentrations in lo-fold excess. A realistic estimate for the accuracy of this technique is + 1 %, thus trace amounts of a molybdo-anion complex would not be detected. Pettersson” determined the
kd (M-3
s- ‘)
rather than “instant
of mixing” con-
speciation of molybdophosphates at 25°C over the pH range 3-9. At low pH, the total molybdenum(W) concentration was comprised almost solely of complex molybdenum-phosphate species, but above pH 7, only MoOa2- was detected. The combined results of these two studies suggest that the concentration of possible molybdophosphate species at pHs higher than 7.2 would be negligibly small. (a) The ej2ct of SOd2- addifion. The effect of S042- upon the rate of CO(NH~)~MOO,+ formation was assessed by adding Na2S04 to a 0.10 M MOODY- stock solution prior to mixing with 0.0040 M cobalt(II1) at pH 7.54. The results of these experiments are summarized in Table 4(a) and show that SOJ2- has only a marginal effect upon kobs. (b) The e&ct of HP04 2- addition. In view of the small rate enhancement caused by the addition of sodium sulphate to the reactant system, a preliminary stopped-flow experiment was performed, under identical solution conditions, using 0.025 M Na2HP04. The resultant reaction was too fast for accurate detection using the stopped-flow technique. The concentration of HP042- was then reduced by a factor of 50 and the reaction transient
2323
Formation of cobalt(II1) molybdate cation Table 4. General catalysis of CO(NH~)~MOO.,+ formation at 25°C and /L= 1.0M (a) Catalysis by S04*-
[Co”‘]
[Mo04’-]
PH
(M)
(M)
7.54
0.0020
0.050
[SO‘$-]
($?)
(M) 0 0.0025 0.025 0.100
5.05+0.13 5.57f0.22 6.22 f0.22 6.59f0.45
(b) Catalysis by HPO,*-
[Co”‘]
[Mood*-]
PH
(M)
(M)
7.35
0.0020
0.050
[HPO‘,*-] x lo3 (M) 0 0.25 0.50 0.75 1.00
11.5kO.6 18.1f1.2 21.3kO.9 27.7 f 2.5 31.3+ 1.5
7.54
0 0.50 1.00 1.75 2.50 25.0
5.01 kO.18 11.6kO.3 18.7* 1.0 25.3 + 1.2 37.3k3.1 Too fast
7.84
0 0.50 1.oo 1.75 2.50
4.33 f0.23 7.28 kO.26 11.5kO.8 15.9f0.7 20.7+ 1.8
(c) The effect of [NaHCO,] upon the rate of complex formation
[Co”‘]
[MoOa*-]
PH
(M)
(M)
7.25
0.0020
0.025
[HC03-] x lo3 (M) 0 1.25 2.50 3.75 5.00
5.62f0.06 7.96kO.12 10.3 +0.4 12.6kO.5 14.6f0.5
7.50
0 1.25 2.50 3.75 5.00
2.53 f 0.06 4.67kO.14 6.8 f0.2 8.6kO.2 10.9 f0.4
7.81
0
1.38kO.05 5.4f0.2 10.1 f0.9 14.4f0.5 19.3f1.5
3.13 6.25 9.38 12.5
LIRefers to the added concentrations of Na2S04, Na2HP04 and NaHCO, at the “instant of mixing”.
2324
M. R. GRACE
and P. A. TREGLOAN
2cQlm.o -
1.Ym.o
-
12Lxa.o
I 0.001
0.0015
0.032
O.WZ3
O.W3
6.0
Added[Na,HPO,l,i, 0.4)
Fig. 2. The effect of pH and total phosphate concentration on the rate of CO(NH,)~OH*+/MOO~*~ complexation at 25°C and 1.0 M ionic strength. [Co”‘] = 0.0020 M. [Mo04’-] = 0.050 M.
Fig. 3. The catalytic activity of H,PO,- on the rate of CO(NH,),OH~+/MOO~~ complexation at 25°C and 1.0 M ionic strength. [Co”*]= 0.0020 M. [Mo04*-] = 0.050 M.
even at 5 x lop4 M, this species doubled the rate of Co(NH&Mo04+ formation. Further stopped-flow runs were then conducted to determine the dependence of this rate enhancement upon the concentration of added Na2HP04 and the solution pH. Solution conditions and the observed first-order rate constants for these experiments are listed in Table 4(b). Figure 2 shows that rate enhancement is linearly dependent upon Na2HP04 at a given pH, but also that the degree of enhancement is acid dependent. The slope of the plot at each pH is a measure of “catalytic activity”. The results of the linear analysis of these plots are presented in Table 5. In the pH range 7-8, phosphate solutions contain both HP04’- and H2P04-. The relative proportions of these species are determined by the deprotonation equilibria : showed
that
H,PO,-
L ---HP04’-
+H+.
A value of 6.4OkO.06 was determined for p& at 250°C and p = 1.OM (NaClO,), by potentiometric
8.0 10.0 IZ.0 ‘k of total Phosphate as H2P0,-
titration with NaOH in CO,-free water, using the automatic titration equipment described elsewhere.23 This value is in good agreement with the 6.46 + 0.02 quoted by Smith and Martell under the same conditions.24 Figure 3 shows a plot of the catalytic activity vs the percentage of total phosphate as H2PO4-. It is apparent from this plot that rate enhancement is linearly dependent upon [HzPO4-] and also, that this species is the only detectable catalyst in solution. Complications arising from anation reactions between the cobalt(II1) and phosphate species in the stopped-flow chamber were not expected since these processes occur by slow cobalt-oxygen bond fission.” Hence, Co(NH3),Mo04+ formation would be essentially complete before any cobalt(II1) phosphate complexation occurs. (c) The eJ?ixt of HC03- addition. The addition reaction of CO2 to CO(NH~)~OH’+ occurs over the same time frame as the molybdate complexation processz6 and hence poses a potential interference to the desired reaction sequence (due to competition
Table 5. Determination of the pH dependence of Na,HPO, catalytic efficiency on the rate of CO(NH,),OH~~+/MOO,*- complexation at 25°C and 1.O M ionic strength
IX+1
% of Na,HPO,
PH
x 108
as H,PO,-
7.35 7.54 7.84
4.47 2.88 1.45
10.30 6.90 3.55
Slope (catalytic activity)
Intercept
Correction coefficient
19900+1400 12600+420 6580 f 340
11.7f0.6 5.1 fO.l 4.2kO.2
0.9956 0.9978 0.9976
Formation
of cobalt(II1) molybdate cation
2325
Table 6. Determination of the pH dependence of NaHC03 catalytic efficiency on the rate of CO(NHJ~OH~~+/MOO~~- complexation at 25°C and 1.O M ionic strength
P+l
% of carbonate
PH
x 108
as H&O3
7.25 7.50 7.81
5.62 3.16 1.55
1.15 x 10-2 6.62 x lo- 3 3.26 x lo-’
% of carbonate as HCO,- ’
Slope
Intercept
Correction coefficient
94.47 96.27 96.88
1830+62 1655f45 1363+37
5.63 +0.06 2.54kO.06 1.37f0.09
0.9998 0.9996 0.9994
between MOO,*- and CO2 for the cobalt(II1) ions).
However, the nature of the Co”‘/Mo042- complexation reaction and a careful choice of experimental parameters enables some very informative, semi-quantitative results to be obtained. Taylor found that with [MO”‘] in large excess, the rate of Co(NH 3)SMoO_,+ formation was independent of [Co”‘] and observations during preliminary stoppedflow runs in this work indicated that the presence of NaHC03 had no obvious effect on the reaction transient for this process. The transients were firstorder for at least five half-lives and also showed no diminution in absorbance change during reaction upon increased bicarbonate concentration. There was also no indication of a second, first-order transient corresponding to Co(NH 3)&O 3+ formation. By limiting the lowest pH solutions to 7.25, the amount of dissolved CO2 was kept below 5% of the total carbonate concentration and the ratio of Mo04’- to CO2 was never less than 2000. It was found that the presence of NaHCO, in solution during stopped-flow runs substantially increased the rate of CO(NH~)~MOO~+ formation. Solutions were prepared from CO,-free water in a nitrogen-filled glove-box. The results of the catalysis experiments are summarized in Table 4(c). The catalytic activity of NaHCO, was determined in an identical manner to the phosphate system and the linear fit parameters to the kobsvs [NaHCO,] plots are given in Table 6. The determination of the active catalytic species is complicated by extensive carbonate speciation and potential reaction with CO(NH~)~OH*+. Over the pH range studied, CO3*-, HC03-, H&O3 and dissolved CO2 are all present in solution. The relationship between these species and the cobalt(II1) reaction products is shown in the following equilibrium scheme : H20+CO&H2C03&
The deprotonation equilibrium involving HCO,and CO3*- has been well-characterized with a pK3 value of 9.57 at 25°C and p = 1.0 M.24 Dissolved CO2 and H2C03 are in an equilibrium heavily favouring CO3, to the extent that the equilibrium constants K, and K2 are normally combined into one apparent pK,,, for the “deprotonation of CO2”26 (pKapp= 6.02 at 25°C and p = 1.0 M24). A limited amount of work has been done on resolving Kapp into its K, and K2 components. van Eldik and Palmer report values for these constants of 2.28 x 10m3and 4.6 x lop4 M, respectively in 0.5 M NaCl.*’ In the absence of values at 1.0 M NaClO,, these results were used, in combination with K3, to calculate preliminary carbonate speciation in the molybdate solutions, at the pH values used in this study. The close agreement between the resultant Kapp value at p = 0.5 M NaC126 and the 1.0 M NaClO, value indicates that the error introduced by this procedure would be negligible in comparison to the inherent errors in the measurement of the observed rate constants and the effects of cobalt(III)-carbonato complexation. The calculated percentages of total carbonate as H2CO3 and HC03- at the “instant of mixing” are included in Table 6. The plot of catalytic activity against apparent percentage HzCO3, shown in Fig. 4, demonstrates that although the doubly protonated species appears to be the major catalyst for the formation of CO(NH~)~MOO~+, ‘the large positive intercept term indicates that anot.her species, most probably HC03-, is also responsible for some rate-enhancing effect. As there is considerable doubt as to the actual concentrations of these species, no further attempt was made to accurately quantify these effects.
H+ +HCO,-
K1’H+
KE
CO(NH~)~OH*+ 11
CO(NH~)~HCO~ ‘+ &Co(NH3)$03+
+H+
+CO,*-
M. R. GRACE and P. A. TREGLOAN
2326 2om.c
1ml.~
1600.(
MOOS
1X0.(
L
1cao.o 0.0
0.002
O.oW
0.026
The catalytic species, HA, may be any moiety in solution capable of proton donation to the molybdate to assist in water elimination. In this scheme, the reaction between cobalt@) and MoOd2-, observed by Taylor, is catalysed by HMo04- in solution. Although the concentration of HMoO,constitutes less than 0.1% of the total molybdate, it is still present at levels 100-1000 times higher than /H+]. The added catalysts H2P04-, H2C03 and possibly HC03- contribute to the total solution “pool” of proton-donating species. If the mechanism involves rate-determining proton-transfer and water departure, then the degree of rate enhancement should reflect the acidity of the proton which migrates to the molybdenum(W) hydroxy group. The slope of the catalytic activity plot is a direct measure of the effectiveness of that species to increase the rate of CO(NH~)~MOO~+ formation. The slope of the plot for H2C03, (6.0 kO.9) x 104, is much greater than that for H2P04-, (1.9 kO.1) x 103, which can be attributed to the greater acidity of the carbonate species. The slope estimate for H&JO3 is a lower limit since no account was taken of the effect of CO(NH~)~CO~+ formation upon carbonate speciation. Clearly, the existence of any correlation between the pK, of the proton-donating anion and the catalytic efficiency of the species, based only upon two anions, must remain speculative. Further investigation using a wide variety of substitution-inert, proton-donating species may help clarify this situation. It might also be expected that a second-order dependence upon cobalt(II1) concentration would be observed with this species in excess over [MOODY-]. In addition, CO(NH&OH~~+ would be the likely catalyst due to the greater acidity of the aquo ligand’s proton relative to the CO(NH~)~OH~+ proton. Another advantage of using excess .cobalt(III) is that the complications arising from HMoO,catalysis could be minimized. Path (b) in Scheme 1 involves the initial formation of the protonated complex, CO(NH~)~ HMoO, 2+, followed by rapid deprotonation to CO(NH~)~MOO~+. It has been demonstrated that the incorporation of a CO(NH~)~ group into the coordination sphere of various 0x0 anions
J 1.011
o.txm 0.01 , ck of total Carbonate as H2CC ‘3
Fig. 4. The catalytic activity of H&O3 on the rate of CO(NH,),OH~+/MOO,~- complexation at 25°C and 1.0 M ionic strength. [Co”‘] = 0.0020 M. [Mo04*-] = 0.050 M.
DISCUSSION
The reinvestigation of the pH and molybdate dependence for the formation of Co(NH3)5Mo0,+, resulted in an expression for the forward rate term which corresponded exactly with that proposed by Taylor2 [eq. (2)]. After allowing for a plausible difference, by a factor of two, in evaluating the MoOd2- concentration (stock or “instant of mixing”), the individual second- and third-order rate constants comprising this term were also in good agreement. However, after consideration of the catalytic effect of several protonated anions, an alternative reaction scheme can be proposed to account for the apparent second-order dependence upon [MoO,~-], observed in both studies. Scheme 1, shown below, involves direct reaction between the cobalt(II1) and MoOd2- species and also requires proton donation from an external “acid catalyst”. The resultant rate expression from this scheme also conforms to the observed pH and concentration dependences expressed in eq. (2).
(a) CO@JH,)~OH,~++MOO,~- +CO(NH~)~MOO~++H~O, (h) CO(NHJ@H,~+ +HMoO,- --‘Co(NH&Mo04+ (c) Co(NH,),0HZ++Mo0,2-
+H@+H+,
+HA=Y+-CO(NH~),MOO~++H~O+A-,
(d) Co(NH,),0HZ3+ +MoO,‘- +HA --‘Co(NHJ5Mo04++H20+HA. Scheme 1. Various possible reaction paths for the formation cf c~(NI-I~)~M~G,+.
Formation
t 1 2321
of cobalt(II1) molybdate cation
(AsOd3-, POd3-, SeOs3-, Cr04’-) dramatically enhances the acidity of any ionizable protons on that anion. 28-31It is therefore assumed that the pK, of CO(NH~)~HMOO‘,~+ will be significantly lower than HMoO,- itself (3.92 at 25°C and fi = 1.O M’) and consequently, at pH values above 7, the concentration of the protonated product would be negligible. The almost complete indifference of the reaction rate to added SOd2- suggests that the catalysis involves proton transfer. It may be argued that the very small dependence of kobsupon increasing sulphate concentration Fable 4(a)] is due to trace amounts of HS04-, but it is equally plausible that this dependence is simply a medium effect resulting from the relatively high concentrations of added Na2S04 (up to 0.1 M). Taylor noted that changing the electrolyte from NaClO, to LiC104 caused a marked decrease in kobs. It is interesting to note that the addition of Na,SO, did not retard the complexation reaction as was observed with the formation of Co(tetren)HS032+.22 This decrease in rate upon addition of SOd2- was attributed to the formation of unreactive {Co(tetren)OH23+ . . . SOd2-) ion pairs. Unfortunately, definitive assignment of reaction pathways to the terms in the rate expression is impossible due to the inherent proton ambiguities. Paths (a) and (d) in Scheme 1 may be rewritten as reactions involving CO(NH~)~OH’+ and HMo04-, as proposed by Taylor. It is believed that the reaction between Co-OH and HMoO,- is more likely than Co-OH, and MoOd2-, since the latter option requires transfer of both protons from the aquo ligand to a molybdenum(V1) oxygen, whereas only one transfer to the molybdenum(V1) hydroxo ligand is necessary from Co(NH3),0H2+. Conversely, significant ion-pairing and internal hydrogen-bonding would favour the + 3/- 2 reaction. Paths (a) and (b) may also be expressed in terms of external proton donation or catalysis as shown below : (e) CO(NH~)~OH~+ +MoOd2- +H+ + CO(NH~)~MOO~+ +H20 (f) Co(NH3),0HZ3+ +MoOA2- +H+ e CO(NH,)~MOO‘,+ +H20+H+. The attractive feature of this last proposal is that paths (e) and (f) are identical to the general catalysis paths (c) and (d) (except that HA is now specifically H+) and therefore, only two reactions are required to describe the entire complexation process. Figure 5 depicts possible transition states for the various pathways proposed for the complexation
0 CbW&
cd+ + H?aoo;
ii) CcWL&+
+
IMoO;
Fig. 5. Proposed
transition states for molybdate complexation.
cobalt(II1)
process. The HMoO,- anion has been assigned as the reactive molybdenum(W) species in the uncatalysed reactions on the basis of the proton-transfer argument discussed above, and also to allow direct comparison with other molybdate substitution processes which are all expressed in terms of this protonated species. The specific proton-catalysed reaction pathways are incorporated into the general acid terms (iii) and (iv), where HA = H+. The rate constants for the formation of transition states (iF (iv) are given by : (i) kl = k,xK,/K,, (ii) k2 = kblKHM (iii) k3 = k, x KH (iv) k4 = kd
=6.6x 104M-Is-’ = 1.3 x 106M-‘s-’ = 2.1 x 10’“M-3s-’ = 3.6x lO”M-3s-‘*
9
where &, and KHM are the acidity constants for CO(NH~)~OH~~+ (4.79 x lo-’ at 25°C and p= 1.0 M”) and HMoO,-, respectively. The rate constants, k, and k2, for the uncatalysed reaction of HMoO,- with CO(NH~)~OH~+ and CO(NH~)~OH~~+, respectively, are much lower than for a series of bidentate ligands, shown in Table 7, where the second-order rate constants falls in the range (0.42.0) x lo* M-’ SK’. These faster reactions have all been attributed to ligand addition processes where the molybdenum(W) centre expands its coordination to six and Mo-0 bond cleavage is unnecessary. Although the reaction of HMoO,- with H2EDTA2- forms an octahedral molybdenum(V1) complex, the slow rate of reaction results from the necessity to break one Mo-0 bond.
M. R. GRACE
2328
and P. A. TREGLOAN
Table 7. Second-order rate constants for complexation with HMoO,Ligand (L) Co(NH&OH’+ Co(NH&OHz3+ H2EDTA2Oxine(Oxine)SO,‘Hcatechol-
Rate constant (M- ’ s- ‘) 6.6 x 1.3 x 2.3 x 1.5 x 4.0 x 1.9 x
104 106 10’ 108 10’ 108
at 25°C”
Product
Denticity number of L
Reference
Mo03L+ MoOSL+ Mo03L3’Mo02(0H),L,MoO~(OH),L~~~ MoO~(OH)~L~*-
1 1 3 2 2 2
This work This work 33 34 35 36
a Adapted from ref. 1.
A comparison of the rate constants for the two uncatalysed processes (k, and k2), reveals that complexation occurs more quickly with Co(NH,), OHz3+ than the corr e sp onding hydroxo species. The importance of ion-pairing in this reaction scheme was dismissed in the previous study, ’ by analogy with the weak outer-sphere interaction between CO(NH~)~~+ and MOO,‘-. However, the formation of hydrogen-bonding contacts between the anions HSeO,or HCr04and CO(NH~)~ 0Hz3+ resulted in values for the ion-pairing constant, K,,, of between 6 and 8 M-‘.3’*37 Even allowing for an order of magnitude difference between K,, values for the + 3/ - 1 and + 2/ - 1 ionpairs, the intrinsic rate constant for the reaction with the aquo species is still twice that for the cobalt(II1) hydroxo ligand. Thus, the effect of increased hydrogen bonding in the transition state and greater proton acidity for reaction with Co slightly outweighs the higher nucleo(NH3)&Hz3+ philicity of the hydroxo ligand in CO(NH~)~OH~+, in determining the more reactive cobalt(II1) species towards HMoO,-. Even with the absence of supporting activation parameters for this reaction, it still seems reasonable that, in view of the associative nature of the Mo042--oxygen exchange reaction3* the extent of bond formation in the complexation transition state outweighs bond breaking. Consequently, changing the nature of the inert cation substituting into the molybdenum coordination sphere should affect the rate of complex formation. CONCLUSION Whilst the form of the rate law for the complexation reaction between Mo04*- and aquo- and hydroxo-pentamminecobalt(II1) cations in basic solution is consistent with a previous proposal for the existence of dimeric molybdenum(V1) species, general anion catalysis experiments show that an
alternative mechanism is likely. The observed rate of reaction is enhanced by anionic species capable of proton donation to the molybdenum(V1) hydroxo ligand, thereby facilitating water elimination from the transition state. The apparent second-order dependence upon [Mood*-] may be explained by the involvement of HMoO,in proton donation. Consequently, this study found no definitive evidence for the formation of molybdenum dimers in slightly basic solution. Proton ambiguities preclude definitive identification of reactant species in aqueous solution. An investigation into the direct reaction of MoOd2with CO(NH~)~OH~‘+ or CO(NH~)~OH*+ in an aprotic solvent may prove beneficial in this regard. Acknowledgement-M.R.G. Government for the assistance graduate Research Award.
is grateful to the Australian of a Commonwealth Post-
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Formation
of cobalt(II1) molybdate cation
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