Association of common polarographic supporting electrolytes in acetonitrile

Wca

Acta. 1969.Vol 14.pp 135to 142. Peqamon PIWS F?mtedin NorthernItcland

ASSOCIATION OF COMMON POLAROGRAPHIC SUPPORTING ELECTROLYTES IN ACETONITRILE* G. Deprtmcnt

A. FoRcmnt

and J. w.

OLVER

of Chermstry, Umverstty of Massachusetts, Amherst, Mass. 01002, U.S.A.

Abstract-Expermmtd data are reported m conformny wtth the Fuoss-Onsager theory for the ~soclatlorl of several common8” laroi$ aphtc supportmg electrolytes m acetomtrrle. Tetra-ethylammonmm tetrafluoroborate an pert orate are completely dtssocrated, yieldmg a small negatwe value for Ka Sodium perchlorate (completely dtmoctated) and potassmm perchlorate (& = 12) are more extenswely dtssocrated than prevtously reported. Imphcattons of the assoctatron data for tetramethylammonium chlorrde and tetra-ethylammomum bromtde on metal-hahde eqml~brmm studres in acetomtrde are dwussed R&mm&Des dorm&esexperrmentales sent p&se&es conformement a la theone de Fuoss-Onsager, propre 51la Qssoctat:on de divers electrolytes supports polarographrques communs dans l’ac&omtrde. Le t&raSuoroborate et le perchlorate de t&ra&hylammomum sont completement drssoctes, dhotant une valeur lQ&rement negative pour KA. Le perchlorate de sodmm (compl&ement drssoctt) et le perchlorate de potassium (& = 12) sont plus dtssoct6s qu’ant6rmurement mdrque Dtscusston des rmphcattons des don&es d’assocratton propres au chlorure de t&ram&hylammomum et au bromure de tCtra&hylammomum, quant aux equrhbres m&al-halog6nure dans l’ac&omtnle Z---Man pr&entrert expertmentelle Daten zur Theorie von Fuoss-Onsager filr dte Drssoxtatron emtger, be1 der Polarographre m Acetomtril verwendeten Lertekktrolyte Tetraiithylammomum-tetrafiuoroborat turd perchlorat smd vollstilndig drssoxuert und ergeben emen klemen negattven wert fiir KA Natrmmperchlorat (volhg dtssoxiiert) und Kahumperchlorat (ZA = 12) sind st%rkerdissoxuert ah frWer angegeben wurde. Man dtskuttert dte Redeutung der Assovattonsdaten ftlr Tetramethylammomumchlond und Tetra-&hylammomumbromtd betm Studmm der MetalHIalogemdX3letchge~chte m Acetomtrtl. INTRODUCTION

many qualitatrve polarographrc mvestigatrons of both inorgamc and orgamc specres m acetomtrrle have been reported,l-s the solution eqmlibna m which these species take part are poorly characterized. Several conductance studres have been carried out m acctonitnle+6 though many of the salts studed are of little interest m polarography. Furthermore, the conductance theory m low dielectric solvents has been extended* only after many of these studies had appeared. For a general understandmg of the electrochermcal reactions, these solution eqmhbria must be known. For our contmuing mvestigatron of the stability and reactions of unusual oxrdatlon states of transltron metal Ions m acetonitrrle, we have examined uiu conductance and polarographic measurements the solution equilibria for several of our polarographic supportmg electrolytes. Results for tetra-alkylammonium tetrafluoroborates indrcated greater dissoaation than we expected from work on other salts’ and led us to re-cxamine the dissociation of tetramethylammonmm chloride, tetra-ethylammomum bromrde and tetra-ethylammomum perchlorate, among the most common supportmg electrolytes used in

ALTHOUGH

* Manuscript recetved 14 September 1967, as amended, 31 January 1968. tPresentaddress* Chas PflxerandCo,Inc,Groton,Conn,USA 135

136

G. A. FORCIJZR and J. W OLVER

acetomtrile. Furthermore, the large differences between reported values for hrmtmg conductance for NaCIO, and KCLO q4*s*sled us to re-mvestigate those salts conductometrically and m the case of NaClO, polarographrcally for then- dissociation equihbna in acetomtrile. EXPERIMENTAL

TECHNIQUE

Acetomtnle was purified by method F of Forcier and Olver.n The solvent was used only if the specific conductance was less than lO-7 mho and it contamed less than lOa molar water by Fischer titration. Tetrabutylammonmm brormde (Bu,NBr, Eastman Orgamc Chermcals) was recrystalhzed from mixtures of acetone and peroxide-free ether. Tetra-ethylammonium perchlorate (Et,NClO,) was prepared and purified by the method previously described.a Tetra-ethylammonium tetrafluoroborate (Et,NBFk) was prepared by neutrahzation of tetra-ethylammomum hydroxide (10% m water, Eastman) with fluoroboric acid (B + A Reagent) and recrystallized from cold water. Tetramethylammomum chloride (Me,NCl, Eastman White Label) was dried at 110°C. Tetra-ethylammomum bromide (Et,NBr, Eastman) was recrystalhzed from ethanol. Potassmm perchlorate (Fisher Reagent) was recrystalhzed from cold water. Sodium perchlorate was prepared and purified by the procedure of Wawzonek and Runner.l The salts were vacuum-dried at 60°C unless otherwise stated. Stock solutions of the electrolytes were prepared by dissolvmg weighed amounts m acetomtnle and diluting to calibrated volumes in flasks. The water content of the salts was tested by Fischer titration m methanol usmg two platmum indicators, AE = 20 mV. In all cases the water contribution from the salts was less than the residual water present m the solvent. Conductances were measured at 1000 c/s usmg a Fisher ‘Low Conductivity’ cell hghtly platimzed and an Industrial Instruments Conductivity Bridge, Model RC-18 (0.1%). The cell constant (0.1181) was determmed by the method of Lmd et allo using an aqueous KC1 solution 4.265 x 1O-4N. The conductance cell was thermostated m a bath at 25.00 f 0 01°C. Five conductance readings were taken on the eqmhbrated solution at 3 to 5 mm intervals. The range of the readmgs was less than 0.05 %. Polarograpmc results were obtained with an Indiana Instrument and Chermcal Corp. Voltammeter, Model ORNL 1988A using a typical three-electrode system and a bridge between the workmg and reference compartments. The reference electrode was an aqueous solution contsumng 1-OM Me,NCl and calomel over mercury, (+0.04 V(sce)). Th e supportmg electrolyte m all cases was Et,NClO, and the solutions were deaerated with prepunfied mtrogen. The salt bridge solution (O-1 M Et,NClO, in acetomtnle) was changed daily to mmmuze the effect of water leakage (< 1 mg/h) to the workmg compartment. A Honeywell Potentiometer (Model 2730) was used to measure accurately initial and final scan potentials. Calculations were performed on the IBM 1620 computer using simple programmes written in the FORGO system. RESULTS

Conductance measurements Measured equivalent conductances at corresponding electrolyte concentrations are given in Table 1. Data are included between lv and 1O-s molar for Bu,NBr, which

Assoaatlon

of common

pdarograpluc

supportmg electrolytes m a&on&de

TABLE 1 MEASUREI)EQUIVALENTCONDUCl-ANCES

Concentratton mol/l

x 10’

Equwalent conductance

1 359 2 718 3 397 6 899 8 492 13 798 16 985 17 248 11410 18 228

1582 1564 1556 152 8 151 8 148 6 147 3 146 8 149 7 146 5

NaClO,

1 734 2 452 3003 3468 1 372 3431 8 578 12 008 1716 2 573

173 4 1704 168 2 166 1 175 6 1726 1682 1662 1749 1740

KCIO,

0 7232 1446 2 893 2 170 3 616

183 6 181 9 1796 1808 178 5

EtNBF,

1052 2104 4 208 6 312 8 416

191 189 187 185 183

Et,NClO,

0 9851 1 970 2790

1863 1845 183 6

ACETONITRILE

Concentration mol/l

lllh0

Bu,NBr

IN

137

x 1W

Equwalent conductance XllhO

3940 4 926 6 975 9 851 13 950 1 395 2 463 3 488 4 185 5 580 7 880

1822 1814 1798 1776 1752 1850 185 5 182.8 181 9 1812 1791

Et*NBr

0 7922 1458 1 980 2 773 3 961 5 833 11670 1 167 1 584 2042 2 916 4 083

1826 181.3 1810 1790 1777 176 0 171 5 182 1 181 1 1803 1792 1778

Me,NCI

1 393 2 358 3 302 4 717 6604 8 358 9751 1 887 2 786 3 482 4 876 6 965 9 434

1895 185 3 1845 1820 1784 176 3 1744 187 5 185 2 1843 1814 178 1 1747

6 8 1 4 8

was studred by Evans et al4 at higher concentratron The last two entnes for concentration and A,, for Bu,NBr are the lowest concentratron values determined by those authors. The data were analysed in accordance wrth the Fuoss-Onsager conductance theory,6 where A=+-SCIJ”+EClogC+(J-FA,)C

0)

is used for unassocrated electrolytes and A = A, - S(Cy)l12 + ECy log (Cy) + (J - FA,,)Cy - KACyAj=

(2)

is used for assocrated electrolytes. The procedure proposed by Fuoss and Shedlovsky” was used, incorporatmg appropriate corrections for the tlurd and fourth terms on the right side of equatrons above because of ease of computatron and applicability

138

0.

A. FORCES and J W. OLVBR

at lower concentration compared with those used in the work of Evans et al.* Details of the computational procedure are avdable elsewhere.18 The direct evaluation of & by Iteration6 was not meaningful in this case because of the O-1% precision of the measurements. Instead the hydrodynamic rati, 6, were estimated based on Walden’s rule mth an added electrostatic correction”‘m14 using tetra-ethylammomum tetraphenylborate as the standard m the equation (2, = (~t.~%B)(%t,N&%,n) -ho, ) The values of A,, and d for tetra-ethylammomum tetraphenylborate have been reported by Bems and Fuoss. 1s Use of this equation leads to some uncertainty in the evaluation ofJ; however, in the concentration range studied an error of 10 per cent in the value of as would lead to an error of less than 0.1% in the value for A+ at the h@est concentratlon used. The term F&C arising from changes in viscosity was neglectable at the low concentrations used. The physical constants for the acetonitrile were: vlscoslty (O-3449 cP), dielectric constant (36.01). The salts studied along urlth the experimentally obtained values of &, and KA are listed in Table 2.

Salt

A0

KA

Bu,NBr Et,NBr Me,NC!l Et,NBF, Et,NClO, NaCIOI KCIO,

162.7f 0 4 1864+02 1949*03 1951&03 188.8f 0 5 1799*02 1872 ,LO2

1.8 f 0 6 4.7 f 1 6 55 f 3 * * + Y24f32

* Smallnegativevaluesfor KA,mdxatmg complete &ssoclatlon. The final corrected Shedlovsky plots from whch K, values are abstractedU are shown in Fig. 1 for several of the salts. Poltuographicmeasurements Polarographc half-wave-potential data were obtamed for sodmm perchlorate as a function of supportmg electrolyte concentration. Schaapfs denved an equation relatmg half-wave potential to Ion-pan formation constants of both the supporting electrolyte and the reducible species, which IS apphcable to the reduction of sodium Ion to amalgam as the supporting electrolyte concentration changes,

El12~E.‘_~~og~-~~og~-E~_o-059~ogK*c~1’2, n n 8

B

n

q/2

(3)

where the symbols have the usual slgmficance.ls Since the dill&on coefficient in the amalgam, D,,, and the activity coefficient m the amalgam,& do not vary with concentration of supportmg electrolyte, C,, the constants m the last term of (3) can be evaluated if changes m diffusion coefficients in

Association of common polarographrcsupportingelectrolytes111 acetomtnle

139

560m-m-

n

n

540 ??

zi

5*o;;

500 0

I 20

t 40

I 60

1 60

I 100 CASf2

I 120 X

f 140

I 160

I I60

I 200

IO'

FIG.1. Shedlovsky plots for severalsalts 0, Sodmm perchlorate, A, tetraethykunm~mu~ perchlorate,0, tetramethylammolllum .

solution, D,, activity coefficient in solutron,f,, and junction potential, EL, with C,can be evaluated. The value off, was determined by the procedure of Kolthoff and Thomas.l’ & was estimated using the equation of Henderson and Planckl’ where the junction is a simple one between two concentrations of an electrolyte in a single solvent. Changes in D, were estimated to correspond to 2 mV for a ten-fold change in C,.r7 The NaClO, concentration used in this study was 7.8 x 1W M which yielded an I& of -1.764 V&e) in 0.100 M Et,NClO,. The results are summarized in Table 3, showing the cuhduted changes (assuming no association of NaClO, or Et4NC10,) in Byrdbetween pairs of electrolyte concentrations as f,, EL and D, change and the observed Ella changes over the same concentratron ranges. TABLB~.POLARWRAPHIC RESULTS 78 x IO-'M NacIO,

colic. of two solutions compared M 0.005 to 0.050 0.080 0.008 to 0910 to @100 0005 to 0.030 0.005 to 0080 DO30to @100 0005 to @loo 0010 to 0.050

due to

g; due to

Ghan

change

in 8" mV -Z -10 -1: -If -7

m EL mV +6 +6 :4 +7 +3 +8 +4

Es-ted @II, due to ch%e

UlD mV +2 +2 +2 +1 +2 +1 +2 +1

Total cLE,IS CdC

mV 1: -2 0 7; r;

L\E,/r

observed mV

- &I, observed mV

-S

8 2 0 0

z-33 -2

: 0

1; -4

G A. FORCIER and J. W. OLVER

140

DISCUSSION

Our data for Bu,NBr yield a value of 162.7 f 0.4 for the hnntmg conductance and 1.8 f 0.6 for the assoclatron constant, compared with the values A, = 162.1 and K’ = 2 reported by Evans et al. 4 Then data were taken m a hrgher concentratron range wrth greater preasron. The agreement, however, supports the rehabrhty of our data. The data for NaClO, and KClO, are of considerable interest. Evans et al4 rccalculated the data of Mint and Werblans on alkah perchlorates m acetomtrrle to conform with the Fuoss-Onsager equatron. However, lomc conductances of sodmm, potassium, and perchlorate ion lffer considerably, as reported by Mmc and Werblan* and by Harkness and Daggett5 based on data of Walden and Bnr,lQ Humphreyss and Marcinkowsky.sl These data are summarized in Table 4 along with the values from the present work. The preponderance of data favours our A, values, and the K* values here reported are probably more rehable than those recalculated from the work of TABLET. INDUCTANCE

PARAMETERS POR NaClO, ANLIKCIOl OBTAINED FROM vAluoussoURcEs

salt

Source (references)

Aa

KA

NaClO, KCIO, NaClO, KClO,

Mmc and Werblan4** Mmc and Werblan4s’L Harkness and Daggett6Jg-*1 Harkness and DaggettbJO-*l

1912&04 205 7 f 0 5 179.9 f 0.6 1864f06

NaCiO,

Present work

1799&-02

KClO,

Present work

1872f02

23 f 9 33 f 9 not calculated completely tisoclated 12 * 3

Mmc and Werblan.8 Possrble sources for the discrepancy mcludmg (a) different solvent punficatron leadmg to solvent of lower punty with higher specr8c conductrvrty and higher water content than ours or that used by Evans et al, (b) a high dielectrrc constant (37.5) do not appear to account for the large difference between the values of Mmc and Werblan and those of other workers, mcludmg ourselves. Our data and those of Mmc and Werblan do agree on the greater association of KClO, than of NaClO,. As m water solution the limltmg equivalent conductivrtres of sodium and potassmm ions are not m the order expected according to the ionic radn. The values of 1, = 76.9 for sodmm ran and 1, = 84.2 for potassmm ion indmate that the smaller sodmm IS solvated to an extent greater than potassmm. This m turn explains the greater observed assoclatron of the potassmm perchlorate m acetonitrile, whereas for equal solvation of those two Ions one would expect the assocratron of sodium perchlorate to be the greater. The polarographrc data summarized m Table 3 for NaCIO, show no slgn&cant change m half-wave potentral as the concentration of the supportmg electrolyte IS varied, and support independently the high dtssocratron of NaClO,. The calculated data are based on the assumption of no assocratron of NaClO, or of the supportmg electrolyte. Each of the potential values recorded m that table IS uncertam to roughly f 1 mV whether calculated (uncertainty m hquld junction potentrals and actrvrty coefficients) or observed (experimental error) A drfference of 4-5 mV between total AE,,, (calculated) and AE,,, (observed) (final column) would be required for that result to be slgmficant. However, the assocratron constant for NaCIO, must be less than 10 or that difference would appear as a srgmficant quantrty l5

Assccratlonof co-on

polarogaphlc supporting electrolytesm acetonitnle

141

From the conductance data on Et,NCIO,, Et,NBF, and NaCIO,, Tables 1 and 2, these salts are completely dissociated m pure acetomtnle. Analysis for KA m accordance with (2) led to shght negative values for EtaNBFa and Et,NClO, and a value indistmguohable from zero for NaClO,. This analysis is shown m Fig. 1 for Et,NClO, and NaClO,, where the hnes are least-square fits to the data and have shght negative slopes. The physical meamng of the negative KA values reported here and also for several picrate and other salts by Evans et al4 is uncertain but suggests there may yet be uncertamties m the conductance theory which must be tested further. The values of A, and & reported by Coetzee’ for Et,NlCO, were not calculated in accordance with the Fuoss-Onsager theory. e The value KA = 13 & 3 calculate from our data on the same basis used by Coetzee in interpreting his data is m reasonably good agreement, however, with the value J& = 18 which he reported. Recently Ahmed and Schmulbach22 reported the complete dissociation of Et,NClO, m acetorutrile from conductance data. They did not compare their conclusions with those of previous workers7sls and their equivalent conductance values at various concentrations differ considerably from those reported m Table 1 of this communication, which we cannot explam. The values of A, and K, for Et,NBr (185.13 and 14.5) reported by Harkness and Daggett5 also were not calculated m accordance with the Fuoss-Onsager theory and therefore show some deviation from those listed m Table 2. The values of A, and K’ (192.9 and 56 f 6) for Me,NCl recalculated by Evans et al4 from the data of Popov and Skellyl* are m good agreement with those reported m Table 2 Ahmed and Schmulback22 also reported complete dissociation for tetra-ethylammonmm chloride (Et,NCl), which should be compared with results for Me,NCl and Et,NBr. For series of symmetrical tetra-n-alkylammonmm cations in association with a smgle halide, the tetramethyl halides are always more associated than higher alkyl members of the series.4*6 For a single tetra-alkylammomum cation m association with the various halide anions, previous data4*6*18generally show greatest association for the chloride and least for the iodide. The difference between complete dissociation for Et4NClee and KA = 4.7 & 1.6 for Et,NBr (present work) is counter to the trend cited, and more work would be desirable at lugher precision than either of the studies cited afforded. The imphcatlons of the foregoing data for typical polarographic supportmg electrolytes and complexing agents are important in any quantitative study of electrode and solution eqmhbrla m acetomtnle. Perhaps the most used supportmg electrolyte (Et,NClO,) JS fortunately completely dissociated and has a high solubihty in CH,CN, therefore yielding low solution resistances. Quantitative studies of halide complexation and its effect on the electrochemistry of transition metal ions are considerably comphcated by the association of the source of halide, which has been usually Me,NCl for chloride or Et,NBr for bromide. In quantitative studies on the halide complexation of zmc by Kolthoff er al2 and of zucomum and hafnium by Olver and Ross,% the authors used Me,NCl and Et,NBr as halide sources. To obtain meaningful quantrtative data m such cases, highly dissociated hahde sources must be used or their association equilibria must be understood with greater certainty than is at present the case Acknodedgemenr-Thm work was supported by the Directorate of Chemcal Sciences, AFOSR Grant 117-65 2

142

1. 2 3. 4. 5. 6. 7. 8 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20 21. 22. 23.

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