Indian Journal of ChemistryVol35A, May 1996, pp. 416-420

Mechanism of vanadium(V) oxidation of thallium(I) in aqueous aceticacid medium: A kinetic study

P L Timmanagoudar, G A Hiremath & S T Nandibewoor*P G Department of Studies in Chemistry, Karnatak University, Dharwad 5ROi)()3

Received 23 August 1995; revised 16 November 1995

The vanadium(V) oxidation of thaJlium(I) in aqueous acetic acid containing hydrochloric acid isconsiderably accelerated both by the hydrogen and chloride ions as well as by increasing acetic acidcontent of the medium. The experimeI1lal results obey the rate law:

d[V(V)]dt

where PI' P2 and P3 are cumulative stability constants of the species TICI, TICli, TIClj - respect-ively, k is the rate constant of the slow step and K 1 is the equilibrium constant for, '

H4VO: +2H+ .,. V(OH)3; +2 H20

The main active species of vanadiumfV] and thallium(I) are V(OH)i+ and TIClj - respectively. Thereaction constants involved in the mechanism are derived.

Although considerable work has been done onvanadium(V) oxidations in acid medium', the na-ture of the active species is not well understoodand different workers have suggested differentforms of active species. in solutions/. The reduc-tion potential" of the couple V(V)!V(IV) being 1.0V, is fairly sensitive to acid concentrations. In-crease in hydrochloric acid and acetic acid con-centration increases the reduction potential" ofV{V)!V(IV) system. While this is the case of V{V)/V{IV) couple, the reduction potential of TI(III)/TI{I) couple in acid medium (H2S04/HCI04: 1-,25V)3 undergoes a decrease with in-creased in acetic acid and hydrochloric acid! con-tents, the value reaching 0.6 V in 50% acetic acidcontaining 3.0 mol dm-3 of hydrochloric acid.Thus there is a possibility of oxidation of thalli-um{I) by vanadium{V) in the latter medium. In-deed, solutions of 6.0 mol dm "? hydrochloric ac-id have been used in the titrimetric analysis ofthallium(I) by chromiumfVlj-". The total reversalof redox potential of two reactants in such mediais reason enough to expect drastic changes in theidentity of the reacting species. While the detailsof the thallium(III)-iron(ll) reaction in dilute acidsolutions are reasonably well understood? as, forexample, involving thallium(II) species, the rever-sal of their roles as to electron transfer needs tobe understood. The reaction is of considerable in-

terest because of its non-complementary natureand thus enceunterirrg different intermediates andspecies. Moreover, as all the oxidimetric titrationmethods for thallium(I) require the presence ofchloride ions", the role of hydrochloric acid, with-out which the reaction does not occur, needs tobe understood. Since the formation of thallium{I)chloride complex is known to be greatly facilitat-ed in aqueous acetic acid solution, it is worth-while to study the title reaction in aqueous aceticacid (50% v/v) containing 3.0 mol dm-3 HC\.Herein we report the detailed study of the titlereaction.

Materials and MethodsReagent grade chemicals and doubly distilled

water were used throughout this work. The stocksolutions of vanadium{V) were prepared by dis-solving a known weight of NH4V03 (Reidel) in0.5 mol dm - 3 of HCI and standardised" with(NH4)2S04.FeS04.6H20(AR) solution using bar-ium diphenylamine sulphonate indicator. The thalli-um(I) . solutions were made by dissolving aweighed quantity of TICI(BDH) in distilled waterand standardising against potassium bromate us-ing methyl orange indicator!", The vanadium{IV)solutions were prepared using VOS04.H20 (Sis-co) in H20 and thallium{III) was prepared by dis-solving required amount of TIZOJ (BDH) in 0.25

TIMMANAGOUDAR et al.: VANADIUM(V) OXIDATION OF THALLIUM(I) 417

mol dm-3 of hydrochloric acid. NaCI04 and HCIwere used to maintain the ionic strength and re-quired acidity respectively. HCI04 and NaCL werethe sources of H + and CI- ions, utilised to studythe acid and chloride ion concentrations respec-tively in the reaction medium. Acetic acid used asthe medium was of Glaxo Excl R, and was alsopurified by refiuxing with chromium(VI) oxide for5-6 hours and then distilling; a further distillationyielded the fraction at 1180 to be used.

Kinetic runThe thermally equilibrated solutions of vanadi-

um(V) and thalliumil) also containing the requiredquantities of HCI, NaCI04 and acetic acid weremixed and transferred to one of the thermostatedcell compartment of Hitachi 150-20 spectropho-tometer instrument and absorbance of vanadi-um(V) at 310 nm was followed. The applicationof Beer's law under the reaction condition hadearlier been verified (E = 575 ± 3).

The kinetic runs were carried out under pseu-do-first order conditions with [thallium(I)] main-tained nearly ten-fold excess over [vanadiumtVj]at 25± 0.1°C unless otherwise stated. The reac-tion was followed for more than three half livesand the initial rates were determined by plottingconcentration of vanadium(V) versus time usingplane mirror method 11. The choice of initial rateswas made owing to the nonlinearity of the plotsof 10g[V(V)]'versus time after 60 to 70% of thereaction and also due to the fractional order in[vanadium(V)]. Initial rates were reproduciblewithin ±5%.

ResuhsStoichiometry

Six different sets of concentrations of reactantswere mixed at constant [acid] and [chloride] (3.0mol dm=' in each case) in the presence of 50%acetic acid. The oxidant, vanadium(V) was ana-lysed by measuring its absorbance at 310 nm and[thallium(I)] was found by titration with bromatein 6.0 mol dm"" of HCl using methyl orange indi-cator. The product [vanadium(IV)] was measuredby recording its absorbance at 745 nm(E = 20 ± 1%). The stoichiometry was found to bein ratio of two moles of oxidant to one mole ofreductant.

2 V(V)+TI(I)-+2 V(IV)+Tl(m) ... (1)Reaction order

The reaction orders were determined from theslopes of log/initial rate) versus log(concentration)plots. Varying [vanadium(V)] in the range of

2.0 x 10- 4 to 2.0 x 10- 3 mol dm - 3 at constant[thallium(I)], [HCI] and ionic strength of3.0 x 10-3, 3.0 and 4.0 mol dm-3 respectively,the order obtained was less than unity with res-pect to [vanadium(V)] (Table 1). With similarreaction conditions and with a constant [vanadi-um(V)J of 7.0 x 10-4 mol dm -3, the [thallium(I)]was varied between 5.0 x 10-4 to 5.0 X 10-3 moldm - 3. The order found was - unity with respectto [thallium(I)](Table 1).

The reaction products vanadium(IV) and thalli-umrlll), added initially in the concentration range2.0x1O-4 to 2.0XlO-' and 5.0x1O-4 to5.0 X 10-3 mol dm-3 respectively (other condi-tions and [reactant] being constant), did not affectthe title reaction significantly.

Effect of varying [CZ-]The effect of [CI-] on the rate of reaction was

studied by varying [Cl-] in the range of 2.0 to 4.0mol dm - 3, other conditions, reactants concentr-ations and [H"l being constant (Table 1). NaClwas used as the source of chloride ions. As thechloride concentration increases, rate of the reac-

Table I-Effect of [V(V)J, [Tl(I)J, [H+J and [Cl-J on the vana-dium(V) oxidation of thallium(I) in aqueous acetic acid (50%

v/v) at 25°C, /=4.0 (mol dm-3)

[V(V)Jx W [Tl(I)Jx W [H+J [Cl-J 104 x In!.mol dm-3 mol dm" ' mol dm "? moldl)1-3 rate

mol dm "?

S-1

2.0 3.0 3.0 3.0 1.124.0 3.0 3.0 3.0 1.997.0 3.0 3.0 3.0 3.3110.0 3.0 3.0 3.0 4.1720.0 3.0 3.0 3.0 7.41

7.0 0.5 3.0 3.0 0.587.0 1.0 3.0 3.0 1.077.0 2.0 3.0 3.0 2.247.0 3.0 3.0 3.0 3.477.0 5.0 3.0 3.0 5.75

7.0 3.0 2.0 3.0 1.747.0 3.0 2.5 3.0 2.517.0 3.0 3.0 3.0 3.387.0 , 3.0 3.5 3.0 4.377.0 - 3.0 4.0 3.0 5.60

7.0 3.0 3.0 2.0 1.257.0 3.0 3.0 2.5 2.197.0 3.0 3.0 3.0 3.557.0 ~.O 3.0 3.5 4.787.0 3.0 3.0 4.0 6.62

Error ±5%

418 INDIAN J CHEM, SEe. A, MAY 1996

tion increases significantly and the order in [CI- Jfound from log-log plots of initial rates versusconcentration was 2.5 (Table 1). The significantincrease in rate with respect to [Cl "] led us tothink of chloride complexes of thallium(I), sincethere is no evidence for the vanadium(V}-chloridecomplexes and this possibility was ruled out else-where".

Thallium(I) is known to form different chloridecomplexes':' of the general formula (TlCIn)l-n,where n is the number of chlorides complexedand the respective equilibria are shown in equ-ations (2) to (4 ),

TI+ +CI- :1TICI ... (2)

T1CI + CI- ~2 TICI;,- ... (3)

TICl- + CI- KJ nC(2-2 ~ .1 ... (4)

TICl4 3 - has also been suggested to form, hut, ifit does, its concentration may be expected to besmall. The complexes (2) to (4) have the equilibri-um constants, K1, K2 and K3 as 16.0, 0.63 and1.0 respectively". The concentrations of thalli-um(I}-chloride complexes were calculated usingEq. (5),

[TI+h=[TI+]f{1 + Pl[Cl-]+ P2[CI-j2+ P~Cl-PI

... (5 J

where [TI +h and [TI+]f are the total and uncom-plexed [TI(I)] respectively arid PI' P2 and P3 arecumulative stability constants of the complexes ofequations (2), (3) and (4) respectively". The con-centration of TICl~ - shows' an increasing trend

Table 2-Effect of [CI-] on Tl(I) species" and initial rate ofvanadium(V) oxidation of thallium(I) in aqueous acetic acid

medium at 25·CI[V(V)]=7.0x 10-4, [TI(I))~3.0X 10-3, [W]=3.0

and /=4.0(moldm ')IUt, x 10.1 al ~ a3 lnl, rate x 104

2.02.53.03.54.0

moldm-.1 S-I

6.50 0.208 0.262 0.524 1.25

3.80 0.153 0.241 0.602 2.192.42 0.116 0.220 0.661 3.551.63 0.091 0.201 0.705 4.781.15 0.073 0.185 0.740 6.62

·a", aI' a-; and a3 are the fractions of total TI(I) of the speciesTI,' ,. TICI. TICI; and TICq - respectively. Error ± 5%.

with increasing. [Cl- j, while the concentration ofother complexes show a decreasing trend (Table2). The variation of TICI~- with log]CI-] is stri-kingly similar to that of log (rate) versus log{Cl-](Fig. 1). Unfortunately lower acid concentrationthan used in the study was not possible due to thelimited solubility of thallium(I) in such acid con-centrations.

Effectof[H+)Increasing [acid], at constant [chloride], acceler-

ates the reaction significantly and the order in[H +], found from log-log plots of initial rates ver-sus concentrations was found to be 1:6 in therange of 2.0 to 4.0 mol dm -.I (Table 1). NaCI wasused to maintain the [Cl"] constant and HCI04was used to vary the [H+]. The order in [H+] in-dicates that probably the reaction involves twoprotons. In fact, vanadium(V) is known to formprotonated species? in acid medium. In higher ac-id concentrations, VO; ion is completely convert-ed into tetrahydroxovanadium(V) ion, H4YO: ,which on protonation in a rapid equilibrium inaybe expected? to form mono and diprotonated oxi-dant species as given in Eqs (6) and (7)

H4VO; + H+ ~ V{OH)2; + H20V(OHW + H i'~ V(OH)~+ + H20

... (6)

... (7)

Effect of ionic strength and solvent polarityThe ionic strength was varied between 2.5 to

4.0 mol dm - 3, and it is found that the rate ofreaction increases substantially with ionicstrength. A plot ·of log(rate) versus (1)1/2 was line-ar with a positive slope.

The relative permittivity (ET) effect was studiedby varying the acetic acid content from 20 to

0.' 0.•

- 0-&•s ...•••s•..

().3 ().6

L- __ -='":__ --::'=- __ ---;;'';;--'0.50.4 0.5 0-6

IOljl[CI-]

Fig. I-Plots of log(rate) versus log{CI-] and Q.1 'tersuslog{CI -]. conditions as in Table 2.

TIMMANAGOUDAR etal.: VANADIUM(V) OXIDATION OF THALLIUM(I)

60% in the reaction mixture with all other condi-tions remaining constant. Attempts to measure therelative permittivity of the media failed due to thecomplexity of the reaction medium. However,they were computed from values of the pure li-quids as in earlier work". Therefore, the solventeffect might be related to mole fraction of aceticacid content in the reaction medium. There wasno reacton of solvent with the oxidant under theexperimental conditions used. Initial rates in-creased with increase in mole fraction of aceticacid. The plot of log(rate) versus lifT was linearwith positive slope.

Effect a/temperatureThe rate constants (k) of' the slow step were

obtained at four different temperatures and wereused for calculation of activation parameters. Thevalues of k (dm! mol-1 s -1) are 2.8 x 102 ± 15,4.2 x 102 ± 20, 5.75 x 102 ± 30 and 8.55 x 102 ± 45at 25, 31, 36 and 42°C respectively. sn : andt!.S '" were calculated to be 45.0 ± 2.0 kJ mol- 1and- 39.44 ± 5.0 J K- 1 mol"! respectively fromthese data.

DiscussionThe stoichiometry of vanadium(V)-thallium(I)

redox reaction is found to be 2:1, and order inoxidant and reductant are fractional and unity re-spectively, No product effect has been observed.Thus the order and stoichiometry of the reactionare computable with the mechanism of Scheme 1,where intervention of thallium(II) species, formedin the first slow step is involved.V(V) +TI(I) -+ V(IV) +TI(II) slow

V(V) +Tl(II) -+ V(IV) +Thfll) fast

Scheme 1The oxidation of thallium(I) by vanadium(V) is

found to be slow in HCI04 and H2S04 media andis also limited to low -concentration due to lesssolubility of thallium(I) in these solutions. Further,as stated earlier, thallium(I) requires the presenceof chloride ions" for its reactions to occur.

At constant [H+], chloride ions have a distinc-tive effect on the reaction, accelerating it to a fas-ter rate. As seen from the result section and Table2, it is apparent that, of the different chloridecomplexes of thallium(I), the onset of TICI~- for-mation has the dominant effect in increasing therate. Figure 1 shows that the variation of the frac-tions of the total thallium(I) present as TlCl~-, a.,withIogl CI-] is similar to that of log(rate) versuslog[Cl-]. The other fractions of the total thalli-

419

um(I) vary in opposite direction to the rate. Theformation of TICI~- has also been expected inconcentrated chloride solution, but its concentra-tion is too small to affect the reaction appreciably.

The significant effect of hydrogen ions to accel-erate the reaction at constant chloride concentra-tion suggests the involvement of one or more pro-tonated species=-" in the reaction. As mentionedin the result section, the order of more than unityin [H"] indicates, probably, the involvement of di-protonated species of oxidant in the reaction. Therole of proton is to weaken the strong bond be-tween the negative oxygen atom and the centralmetal atom. The hydroxide group is converted in-to water molecule which can easily be lost2.17.1Hand the increased positive charge on the oxidantfacilitates the attack by TICI~- ion.

It is possible, however, that the reaction inaqueous media containing lower percentages ofacetic acid is mainly between TICI and the oxi-dant. As the acetic acid content in the medium in-creases, the formation of TlCI~- is facilitatedmore and more as is apparent from its formationconstant of 1.37 in an aqueous 3.0 mol dm-3 hy-drochloric acid!" solution rising to 10.1 undercomparable conditions in a 50% acetic acid solu-tion'" containing 3.0 mol dm " hydrochloric acidwhile the reaction between TICI and oxidant isslow, the formation of TICI~- greatly acceleratingthe reaction. Therefore, the effect of increasingacetic acid content on the reaction is to increasethe rate of reaction by a gradual change in theidentity of the substrate itself.

The mechanism, therefore, may involve the spe-cies V(OH)~ + of oxidant and TICl~ - of reductantmainly. Such a mechanism (Scheme 2) involvesthe formation of above species in prior equilibria,which are followed by steps similar to those ofScheme 1. The step forming thallium(II) species isthe rate determining step. In agreement with theobserved rate law and other experimental results,a mechanism like that in Scheme 2 may be envi-saged.

H4VO: + 2H+ ~1 V(OH)32++ 2HzO (I)

TI+ + 3 cr !3 TICe3- (II) ~

TICl\- +V(OH)3rt V(OH)~+ +TICli slow

... (III)

V(OH)~+ +TICl:l-+V(OH)~+ +TICI1 fast... (IV)

Scheme 2

420 INDIAN J CHEM, SEe. A, MAY 1996

Although experimental evidences for the forma-tion of thallium(II) was not observed, such a spe-cies was observed in earlier studies'v". Scheme 2leads to the rate law (8)

d [V(V)] kKI/33[V(V)]~TI(I)h{H+]~[CI-]idt \1 + /31[Cl-]+ ~2[Cl-]2 + /33[CI-n

1x------------~------~~--~------~~!1 + 2Kl[V(V)][H+] + K1[H+l' + 2K ~[V(V)][H+n

... (8)

which explains all the observed orders. The term,lK flV(V)]2[H+j2},supposed to be in the denomin-ator of the right hand side of the Eq. (8) is, neg-lected due to low concentrations of vanadium(V).By rearranging Eq. (8) into form (9),

/3,[Tl(I)h [Cl-H 1- x---{1+/31[CI ]+/32[CI ]2+/33[CI ]3} rate

= 1 _! 1 +11+~!~+Kl[H+ll ... (9)k[V(V)] K1[H+] k [H ]

using Eq. (9) a plot of L.H.S. versus 1/[V(V)] isexpected to be linear, which is verified in Fig. 2.The slope and intercept of such plots led to thevalues of K, and k at 25°C as 46.55±2.0 dm!mol-1 and 2.8 x 102± 15 dm ' mol-1 s -1 respect-ively. These values are utilised to calculate initialrates under several' experimental conditions andthere is a reasonable agreement between them.

It, thus, becomes apparent- that the reaction be-tween vanadium(V) and. thallium(I) in 50% aceticacid containing hydrochloric acid is rendered.pos--sible by the altered redox potentials broughtabout due to the facile formation of chloride com-plexes of thallium(I) and protonated oxidant spe-cies such as TICI~- and V(OHH+ respectively.

•

2 345

[V(V)r'• 10-3 dm3 mol-1

Fig. 2- Verification of rate law( H).conditions as in Table I

The role of chloride in aqueous acetic acid mediais crucial to the reaction. The effect of ionicstrength and dielectric constant on the reactionmay be understood in terms of their opposing ef-fects on the first three steps of Scheme 1, whilestep (I) is favoured by increasing ionic strength,the steps (II) and (Ill) may occur to lesser extent.The effect of solvent polarity'? favours the step(II) and (III), not step (1). The less negative en-tropy of activation is indicative of the less proba-bility of involvement of intermediate complex inthe reaction (Scheme 1).

References1 Babel D K, Banu R, Shankar R & Bekore G V, J sci ind

Res; 43 (1984) 250; Waters W A & Littler J S, Oxidationsin organic chemistry, Part I, (Academic Press, New York)(1965), p 196.

2 Timmanagouder P L, Hiremath G A & Nandibewoor S T,J chem Soc Dalton Trans (1995) (In press); Secco S &Grati C, J chem Soc Dalton Trans, 1 (1972) 1675; SubbaRao P V, Venkateshwara Rao N, Murthy R V S & MurtyK S, Indian J Chern, 15A (1977) 16; Subba Rao P V S,Murty P S N & Murty R V S, J inorg nucl Chern, 40(1978) 295; Dyrssen D & Sekine T, Acta Chern Scand, 15(1961) 1399.

3 Day M C & Se1bin J, Theoretical inorganic chemistry(Reinhold, New York) (1964), pp. 226.

4 Corgell CD & Yost D M, J Am chem Soc, 55(1933)1909; Carpenter J E, JAm chem Soc. 56 (1934) 1847.

5 Sagi S R, Prakasa Raju G S & Ramana K V, Talanta, 22(1975) 93; Gokavi G S & Raju J R, Polyhedron, 6 (1987)1721.

6 Buzas I & Erdey L, Talanta, 10 (1963) 467.

7 Johnson Jr C E, J Am chem Soc, 74 (1952) 959; AshurstKG & Higginson W C E, J chem Soc, (1953) 3044; Chi-matadar S A & Raju J R, J inorg nucl Chern, 43 (1981)1947.

8 Kolthoff I M & Belcher R, Volumetric analysis, . Vol. 3. (Interscience, New York) (1957), pp, 102, 103, 104, 453,

461,519,641.9 Pal B B, Mukherjee D C & Sengupta K K, J inorg nucl

Chern, 34 (1972) 3433.10 Gokavi G S & RajuJ R, Polyhedron, 6 (1987) 1721.

11 Hiremath S C, Chimatadar S A & Raju J R, Trans MetChern, 19(1994)636.

12 Me1win W S & Gardon G, Inorg Chern, 11 (1972) 1912.13 Lee A G, The chemistry oj thallium (Elsevier, New York)

1971,29.14 Gokavi G S & Raju J R, Int J Chern Kinet; 20 (1988j

365.15 Blackburn T R, Equilibrium (Rinehart and Winston, New

York)(1969),p.75.16 Radhakrishnamurthy PS& PatiSC,IlldiaIlJChem, 7A( 1969)

fiR7.17 GuptaK S&Gupta Y K,J chemEduc, 61 (1984) 972.18 Edwards J 0, Inorganic reaction mechanism (Benjamin W

A, New York) (1955), p. 137.19 Frost A A & Pearson R G, Kinetics and mechanism (Wi-

ley Eastern. New Delhi) (1970) p. 147.