t h e mono-oxalato complexes of iron(iii)

8/7/2019 T H E MONO-OXALATO COMPLEXES OF IRON(III)

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THE MONO-OXALATO COMPLEXES OF IRON(II1)P A R T 11. K I N E T I C S O F F O R M A T IO N

R. F. BAUERAND W. RIIAcF. SMITHDepartnzent of C hem istry, Queen's U nivers ity, Kin gst on, OntarioReceived October 1, 1964

ABSTRACTThe kinetics of t he formation of th e mono-oxalato complexes of i ron( II1) have been examinedspectropho tometrical ly over the range of t emper atur es 5 to 25 "C in an aqueous medium ofionic strength 0.50 and the range of hydrogen ion concentrat ions 0.03 to 0.45 M . The kinetic-ally significant paths under the conditions studied involve reactions first order in iron(II1)and in bioxalate b ut there appears t o be some decrease in the second order rate con stant withincrease in hydrogen ion concentration a t the highest acidities and a t the highest temperatur es.Although there is no significant contribution to the rate by an acid-independent path firstorder in free oxalate under the experimental conditions, the possibility of t he r at e constant forsuch a path being greater th an t ha t first order in bioxalate is not precluded.

INTRODUCTIONTh e kinetics of the anation reaction of iron (111) ion with fluoride (I ), chloride ( 2 ) ,

bromide (3) , thiocyanate (4), and sulfate (5) have been examined by techniques rangingfrom relaxation methods to stopwatch timing and show a fairly inarlted dependence ofthe second order rate constan t on ligand type. Such behavior is in contrast to tha t exhibitedby certain bivalent metal ions such as nicltel (11) whose rate constants for correspondinganation reactions show little dependence on ligand type (6) and are presumably controlledby the rat e of water loss from the metal ion.

This paper presents more information bearing on the influence of ligand type on theanation reactions of iron(II1). T he reaction studied , tha t of i ron(II1) with oxalate ionand it s protonated forms, is convenient to examine because of its relative slowness in acidsolution. I t differs from the anat ion reaction of iron( II1) with singly charged ligands whichhave been previously examined in that the principal complex formed is probably bidentaterather than monodentate. EXPERIMENTALMaterialsSolutions were prepared from triply distilled water , th e second distillation being from alkaline perman-ganate. Ferric perchlorate (non-yellow), perchloric acid (double vacuum distilled 70 %), sodium perchlorate(anhydrous ), and oxalic acid were all reagent grade. T he sod ium perchlorate was dried a t 110' for 2 h beforeweighing. Stock solutions ferric perchlorate an d perchloric acid were analyzed for acid a nd iron content bythe method previously discussed (1).Method and ApparatusThe reaction was followed spectrophotometrically with a Beckman DIG-1 spectrophotometer a t X 310 mpusing thermostated standard 1 cm silica cells as reaction vessels. Two solutions, one containing iron(II1)perchlorate, the other oxalic acid, and bo th adjusted for ionic strengt h and acidity, were simultaneouslyinjected into the reaction cell from thermostated 2 ml syringes fitted with polythene tips to ensure mixing.The solution in the cell in the reference beam had a composition identical with the reacting solution exceptthat it contained no oxalic acid. At X 310 mp the uncomplexed oxalic acid and its ions do not con tributesignificantly to optical densi ty and t he difference in optical densities ( A D )of reference and sample solutions is

where x is the tot al concentra tion of t he complexes, r is th e fraction of the coinplexes which is FeHC204++,an d E is th e extinct ion coefficient in mole-' 1 cm-I of th e species indicated by the subscript. If i t is assumedtha t a t a given acidity the ratio [FeC204+]/[FeHC204++]remains at the equilibrium value, t he change ofCana dian Journal of Chem istry. Volume 43 (1965)


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2764 CANADIAN JOURNAL OF CHEMISTRY. VOL. 43 , 1965optical density is proportional t o x, the co ncentration of complexed iron. Since as indicated in the precedingpaper the ex tinction coefficients of the two complexed species differ by less tha n 2075, AD would remainnearly proportioned t o x even if there were moder ate deviations from th e equilibrium value.

I N T E R P R ET A T I O N O F T H E K I N E T I C D A TAIn interpreting the experimental rate data we have used the integrated form of therelation

where a is the total concentration of iron(II1) in all species, b is the total concentrationof oxalate, and x is the total concentration of iron(II1) complexed as FeC204+ andFeHC204++.

kt" is a composite rate constant for the formation of the complexes and will depend onthe rate constants of the various paths. I t will involve factors for the fractions of uncom-plexed oxalate in the protonated and unprotonated forms and will consequently dependon the concentration of hydrogen ions. kb is a corresponding composite constant for thereactions involving breakup or aquation of the conlplexes and like ktr' will be aciddependent.

The relation [2] will be valid for a variety of possible mechanisms, e.g. three parallelforward reactions opposed by back reactions i f the ratio of the concentrations of the pro-tonated and unprotonated complexes is maintained close to the equilibrium value.

kzFe++++ C?OaP=F e C ? 0 4 +k-z

In this case eq. [2] takes the form

which may be written

* TBe t r an s it i on st at e f or tlu is r ea ct io n w ~ ~ ~ l dof course be differen t fr om that for reaction [I].


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BAUER A N D SMITH: MONO-OSALATO COMPLEXES OF IRON(II1).PART I1where Qn2is the acid ionization quotien t for HC204-,

Equation [ 2 ]would of course also app ly if the protonated complex were formed solely b yprotonation of FeCZO4+if th e rat es of protonation a nd deprotona tion of th e complexeswere rapid compared to th e rat e of complex formation. I n this event k twould equal

instead of k1+ k3 [H'] + k2.Qn 2

There is the marked possibility that the protonated and unprotonated complexes areformed from a common intermediate. Th e rate expression [ 2 ]would then apply only if th erate-controlling steps involved solely the transformation of the initial reactants to th eintermediate or if the intermedia te occurred in the relatively small concentrations con-sistent with applicability of the stat iona ry stat e approximations.

If x , represents the tota l concentration of complexes a t equilibrium[51 & -- ( a - xc)s(b - xe)kt x e

Substitu tion of [ 5 ]into [ 3 ]with rearrangement yields

which on integration and introduction of t he condition tha t x = 0 when t = 0 yieldsk islogt*) + log( ab ) =-( xe2- ab)t .a6 - x , x 2 . 3 0 3 ~ ~

The second term on the left of [G I is negligible relative to the first. a is much greaterthan b under our experimental conditions and even when x approaches x , the product x,xis rarely as great a s 1% of ab.

I t follows from [ I ]that

where x , and AD , are corresponding values for equilibrium.Substituting [ 7 ]into [G I and dropping t he second term on the left of [G I yields

k tslog 1 -- - ---( t i ~ )- 2 .303~ . ( x e 2- ab) t .


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2766 CANADIAS JOURXAL OF CI-IEMISTRY. VOL. 43, 1965Th e slope S1of the plot of log (1 - (ADJAD,)) against t should then be related to k

by t he relationk f = 2.303x,S1s (xo" ab) '

x, and s can be calculated from the values of the association quot ien ts of th e mono-oxalate conlplexes and t he ionization quotients fo r oxalic acid given in the previous paper.

R E S U LT S O F T H E K I N E T I C M E A S U R E M E N T S A N D D IS C US S IO NTh e da ta from kinetic ru ns f or all conditions investigated yielded linear plots for

log (1 - (ADJAD,)) against time. Ra te cons tant s k pwere calculated from the slopes ofthese plots using eq. [9] and showed no significant dependence on total concentration ofiron(II1) or oxalate a t a given acidity over the range of concentrations investigated, i.e.2 X to 8 X l o p4 for iron(III), 2 X to 12 X for oxalate. Mean values fork p for each acidity and tempe rature along with the standard deviations are given inTabl e I and plotted in Fig. 1against the concentration of hydrogen ions.

G i E250 QIO 0.20 03 0 [H4 (mole0 4 0 I'FIG.1. Plots of observed forward rate constant kr against hydrogen ion concentration.As indicated above

if there are three parallel paths, one first order in oxalate and the other two first order inbioxalate.

Figure 1indicates a linear relationship between kp and [I-I+] (i.e. proportionality betweenra te and bioxalate ion concentration) over the range of acidities investigated for the twolowest temperatures and a significant falloff from first order dependence on hydrogen ionconcentration for the highest acidities a t th e two highest temperatures. Lines of regressionthrough the points a t each temperature, ignoring those for concentrations of hydrogenion in excess of 0.2 iM a t 20 and 25 "C, have slopes implying, through eq. [4], values fork 1+ k3 listed in Table 11. Cursory exainination of Fig. 1 suggests t ha t k fapproaches zeroa t zero hydrogen ion concentration and consequently t ha t k: equals zero. However,


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BAUER AXD SMITH: MONO-OSALATO COMPLEXES OF IRON(II1).PART I1T A B L E I

Average values for the forward rate constant(kr) for react ion between iron(II1 ) and oxalateat various temperatures and acidities. Ionics trength 0.50T e m p . H+ k t X("c) (mole 1 - I ) (mole-' 1 s-1)

statis tical trea tme nt of the dat a, implies values for kz of (19f20*) X lo3a t 15 O C and(30 f 32) X l o 3mole-l 1 s-I a t 25 "C. I t is appa rent th at there is the possibility tha t k2is larger than kl + k3. However, under all condit ions inves tigated k2[C204=] mus t benegligible compared with (kl + k3) [HCz04-] because of t he very large value of t he ra tio[HC204-]/[C204=]and consequently there will be negligible contributions to the overallrat e by the path with first order dependance of r ate on free oxalate ion concentration.

The plot of log (k l + k3) against 1 / T implies effective values for AH * and A F of18 A 4 kcal molep1 and 16f 10 cal deg-l mole-l respectively. In th e subsequent argu -ment we tentatively identify the experinlental (kl + k3) with kl and consequently implyth at t he values for AH * an d AS * apply to the path with ra te constant kl. I t is of interestthat AH * for other anation reactions of iron(II1 ) range from 1 3 ltcal (thiocyanate) t o22.8 kcal (fluoride) and A F from -5 cal deg-I (thiocyanate) to +35 cal deg-l (fluoride).However, the points in t he plot of log k against 1 / T for oxalate show considerable dis-persion and do not rule o ut curvature. We d o not therefore consider our values for AH *and AS * suitable for critical comparison with those for other anat ion reactions. I t is possibleth at a n extensive group of measurements obtained with a technique such as "stop flow"which would reinove coinplications due to back reaction and perinit a n extension of t he*95%con$dence limit.


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CANADIAN JOURNAL OF CHEMISTRY. VOL. 43. 1965TABLE I1

Values for effective rate constant(kl + ks) for reactions formingmono-complexes with first orderdependence on bioxalate ion con-centrationT em p. ( k ~+ kt) X

("c) (mole-' 1 s-l)

temperature range might help elucidate through the shape of the log k agains t 1/T plotwhether (kl + k3) is indeed composite.

I t is apparent t ha t an acceptable mechanism for the reaction between iron(II1) andoxalate must be in accord with a second order rate law, first order in iron(II1) ion andfirst order in bioxalate for all conditions except the highest acidities a t the two highesttemperatures where there is some decrease in the magnitude of the second order rateconstant with increase in acidity.

k1 (and also k3) could conceivably be composite and involve ra te constants for steps in aseries reaction leading t o the formation of the final complex. For example, if t he protonatedcomplex could be ignored the reaction path might be

where FeC204+ .represents a monodentate intermediate. Th at the formation of a mono-dentate intermediate is rate controlling in the formation of the multiden tate coinplexes ofnickel(I1) is suggested by the findings of Wilkins and co-workers (7) and of Margerum andco-workers (8).

If the stationary st ate approach is applicable, the following relation should apply for theforward rat e of formation of the bidentate complex.

At a ny one hydrogen ion concentration the r ate should be first order in iron(II1) and inbioxalate with an effective rate constant kl equal to k,jka/(kb[H+] + kd. If kb[H+] ismuch less than kd then the r ate constant kl equals k, (i.e. the first step is rate controlling)and the ra te constant is independent of acidity. If [H+] is such th at kb[H+] becomes sig-nificant relative to k, there will be a decrease in kl. A corresponding mechanism involvinga protonated monodentate intermediate would imply that kl should be independent ofhydrogen ion concentration. In fact, in the range of higher acidities where there appears tobe a falloff in the observed rat e consta nt , the protonated complex is a significant species a tequilibrium and should not be ignored. Identification of a protonated intermediate withthe final protonated complex would imply a series reaction which in general would show adependence of ra te on concentration more involved tha n the second order dependencewhich is actual ly observed. Th e possibility of there being a reaction pa th parallel to th at


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BAUER AND SMITH: MONO-OXALATO COMPLEXES OF IRON(II1).PART I1 2769summarized by eq. [l o] and involving a protonated intermedia te different from th e finalproduct cannot be definitely ruled out but we shall ignore this possibility and assume t ha tthe essential reaction path is th at summarized b y eq. [l o] and th at the final protonatedand unprotonated species are in labile equilibrium. Such an assumption is in accord withour dat a and our method of interpre ting them and implies values quoted for k1$ k3aboveare those for kl.

Th e work of Eigen and co-workers (6, 9, 10 ) and of Wend t and Strehlow (11) has led togeneral acceptance of a mechanism of ion association which emphasizes th a t there is arate-determining exchange between ligand and t he wate r of the fi rst coordination shell ofthe metal ion within ion pairs where the partners are separated by one water molecule.Th e rat e of water loss may be independent of or dependent on the na tur e of the potentialligand of the ion pair. In the lat ter case the reaction could approach in character t ha t oftendesignated b y SN2. Ion pair formation should be a labile process and an equilibrium con-centration of ion pairs relative to reactan t species may be expected. On th is basis theconcentration of ion pairs between iron( II1) and bioxalate would be Klo[Fe+++J[HC204-]where Klo is an equilibrium quotient whose magnitude may be roughly estimated b y th emethod of Hammes and Steinfeld (12, 13). For ions of oppos ite charge and valences threeand one the value of the association quotient is abou t 14 mole-' 1if th e separation of centersis assumed to be 5 A. Th e values for ion pairs between iron(II1) and various singly chargedanions may be expected t o vary somewhat but not markedly and the value 14 may beconsidered to represent th e order of magnitude of association quotients .

Th e rat e constant k1 for complex formation should on t his picture be Klokl* where kl*is a first order rat e constant. Also, if our value for kl($ k3) is indeed th a t for formation ofthe monodentate intermediate, kl($ k3)/K10 will represent the effective first order r atecons tant for transformation of the ion pairs into th e intermediate and a t 25 OC will be ofthe order of 1440 mole-' 1 s-'/I4 mole-I 1 = 100 s-'. The rate const ant s a t 25 OC for theanat ion reactions of iron (111) in aqueous solution are listed in Table 111 along with va luesfor the reciprocal of the ionization constants (K,) for the acids conjugate to the associatingligand. Th e ionic strengths to which t he ra te constant s apply var y fro m 0.1 to 1.0 bu t theeffect of such variation on the constants should be small. On the basis of the theory out -lined above, the first order rate constants kl* for transformation of ion pair t o complexshould be roughly parallel to the rat e constant s listed for all singly charged ligands. Th evalues listed for the ionization co nstants apply t o zero ionic strengths b ut their reciprocalsshould give fair measure of th e abil ity of the ligand in the ion pair to accept protons fromwater in the inner hydration sphere of the metal ion. I t is apparent t ha t there is a fairlymarked increase in rate constant with increase in ligand basicity and th at the value whichwe have determined for bioxalate is roughly in accord with there being a correlationbetween rate constant and ligand basicity for singly charged anions and monodentatecomplex formation.

Wendt and Strehlow (11) have suggested t ha t the dependence of r ate on ligand type isrelated t o the abili ty of the potential ligand in an ion pair to accep t a proton from one ofthe waters coordinated with t he me tal ion. In thi s picture such an inner hydrolysis leadsto a n increase in the rate of loss of one of t he oth er water molecules coordinated with t hemetal ion. If i t is assumed t ha t the partially desolvated metal ion is the precursor to thecomplex, such an increase in the r at e of water loss could lead to a n increase in the firstorder rate constan t for complex formation from the ion pair and in the second order rateconstant with regard to the primary reactants.

Superficial assessment of the d at a of Table I11 suggests t ha t they a re in accord with such


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2770 CANADIAN JOURNAL OF CHEMISTRY. VOL.43, 1 9 ~ 5TABLE I11

Rate constants of ana tion reactions yielding mono-complexes of iron(I I1) a t 25 "C an destimates of the basicities of the anions involved in the reactions

Anion involved Rat e constant Ref. l/k, Ref.- -B r- 20 3 14C1- 9 .4 2 lo-& 15SCN- 127 4 7 16HCa0.i- 1440 This paper 17.5 17,18HF 12* 1F- 4 000* 1 1.5X103 19Sod- 6 370 5 1O2 20

"Obtained by extrapolation.a view. However, we doubt t he direct applicability of W end t and Strehlow's hypothesisto t he reactions of iron(II1) mainly because the rate constant for water exchange fromiron(II1) is markedly in excess of the value of the ra te const ant for complex formation.Th e rat e constant for the exchange of a water nlolecule from the first coordination sphereof hexahydrated iron(I11) ion has been estimated by Connick and co-workers on the basisof nuclear magnetic resonance (n.1n.r.) measurements to be in the range 2.4 X lo4 to1.1 X lo6s-I (21, 22). The presence of a ligand in the ion pair could conceivably raise thevalue for water exchange of the hydra ted metal ion through inner hydrolysis in the ionpair above this. We have implied th at th e first order rat e constant for formation of t heoxalate complex from the ion pair is approximately 100 s-l, a value substanti ally less thanth at for water exchange. If it is assumed t ha t th e rat e of water exchange measures therate of formation of the pentahydr ated iron(II 1) and th at thi s is the precursor to complexformation it is appar ent t ha t very few of t he partially desolvated ion pairs completereaction to form th e complex and tha t th e reaction

is essentially in equilibrium. If the equilibrium constant for this reaction is representedby IP and the rate constant for transformation of t he partially desolvated ion pair intomono-complex (presumably with proton transfer from the bioxalate ion to t he water ofthe solvent) by k1 then

The associating ligand mill then be influential via its effect on K 1 an d k'. Presumablythere will be partial proton trans fer from water coordinated with th e metal t o the associ-ating ligand in bo th t ypes of ion pairs and although such inner hydrolysis may well affectthe lability of loss and gain of water i t is difficult t o see how there could be much influenceon the equilibrium constant K1. On t he other hand the rat e consta nt k1 for the relativelyslow nucleophilic attack by the ligand would doubtlessly be dependent on the nucleo-philicity of the ligand. We suggest tha t the influence of ligand is through its nuc'eophilicitytowards the metal ion rather tha n its basicity with regards to proton acceptance, and tha tthe appar ent parallelism of ra te cons tan t on the basicity is a reflection of the parallelismbetween basicity and nucleophilicity.

Th e behavior of i ron(II1) and nickel(I1) in anation reactions seems to be in markedcontrast . For nickel the ra te of reac tion is close to the ra te of wa ter loss and in terms of theabove picture this implies th at th e probability of nucleophilic at ta ck by ligand within the


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BAUER AN D SMITI-I: MONO-OXALATO COMPLEXES OF IROX(1II). PART I1 2771p ar t i a l ly des o lv a ted io n p a i r i s mu c h g r ea te r th an t h a t for attack b y water, wh ereasfo r i r o n (I I1 ) th e converse applies . As th e re la t ive nucleophi l ic i t ies of wa ter and a givenligand (e.g. oxalate) should no t be great ly d i f ferent for the reactions involving nicltel(I1)and i ron( I I1) perha ps th e di f ferences in behavior a re to be ex p la in ed i n t e r ms of avai l -ab i l i t y of w a t e r for nucleophilic attack w i t h i n the ion pairs .

REFERENCES1. D. POULIand \\I. NIAcF. S~ II TH .Can. J . Chem. 38, 567 (1960).2. R. E. CON NICK^^^ C. P. COPPELL.J. Am. Chem. Soc. 81,6389 (1959).3. P. MATHIESand H. WENDT. Z. Physik. Chem. Frankfurt , 30, 137 (1961).4. J . F. BELOW,JR. , R. E. CONNICIC,and C. P. COPPEL. J. Am. Chem. Soc. 80,2961 (1958).5. G G. DAVISand W. MAcF. SMITH. Can. J . Chem. 40, 1836 (1962).6. M . EIGENand I

t h e mono-oxalato complexes of iron(iii)

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