Formal oxidation-reduction potentials of copper(II)-copper(I) systems in acetonitrile

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Formal oxidation-reduction potentials of copper(ll)-copper(I)systems in acetonitrile The oxidation capacities of cupric salts in acetonitrile have been found to depend on the nature of the anion z-8 and in the order: perchlorate > nitrate > chloride. The measurement of the formal redox potentials of Cu(II)-Cu(I) systems confirms this observation. Standard solutions of copper(II) chloride, nitrate and perchlorate were prepared by dissolving the appropriate quantities of anhydrous (or dehydrated) salts in acetonitrile; the exact strength was checked by EDTA titrations z-8. Solutions of copper(I) nitrate were prepared by the interaction of metallic copper and silver nitrate in acetonitrile4. Copper(I) chloride, prepared by the reduction of the cupric salt with metallic copper, was used to prepare solutions of copper(I) chloride in acetonitrile5. Various Cu(II)-Cu(I) systems were prepared by either mixing cupric and cuprous solutions, or by partially reducing the cupric solution. Potentials were measured at 25+0.2 ° using Pleskov's electrode and an aqueous saturated calomel electrode, as reference electrodes. Mixtures of copper(II) and copper(I) salts in various proportions were put in the potentiometric cell fitted with a bright platinum electrode. The potential measurements were made of the following two cells: Cu(II) Cu(I) Pt Ag AgN03 (o.oi M in Cl + c2 acetonitrile) in acetonitrile

(i)

and Hg,Hg~Cl~(s) KCl-agar agar bridge satd KC1

Cu(II)

Cu(I) Pt

cl + c2 in acetonitrile

(ii)

The formal redox potentials of Cu(II)-Cu(I) systems were determined using the Nernst equation, from the straight lines obtained (Fig. I) by plotting the measured potentials against the logarithmic ratio Cu(II):Cu(I). The potentials thus obtained ar~ presented in Table I. It is clear from these results that the formal redox potential of the Cu(II)Cu(I) system in acetonitrile is higher for perchlorate than for nitrate and chloride. The polarographic E½-values for the copper(II) species in acetonitrile reported by KOLTHOI~I~and others 6-7 are in good agreement with the measured formal potential of the Cu(II)-Cu(I) perchlorate system. Perchlorate ions have the least tendency to form complexes with metal ions and therefore the perchlorate system has a higher potential; the other ligands, nitrate and chloride, have a pronounced tendency to complex with metal ions and these systems therefore have considerably lower potentials. Conductometric measurements and molecular weight determinations of copper(II) perchlorate and nitrate in acetonitrile show that perchlorate is extensively dissociated compared to the nitrate s. Solutions of copper(II) chloride in acetonitrile are only partially dissociated9. These j . Electr~anal. Chem., 18 (1968) 2oo-2o4

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20I

facts support gradation in oxidation capacity and formal potentials exhibited by different cupric salts. Similar differences in redox potentials have been observed with the ceric-cerous couple associated with different anions 1°. 11oo1000.-

900-

b

800

E 700 o 600

E 50o L

,,o Y

/

400

300. 200 -0.8

-o14

do

[co (m]

- -

0:4

0.8

log [Cu (I)]

Fig. I. Variation of formal o x i d a t i o n - r e d u c t i o n potentials with log([Cu2+]/[Cu+]) of copper salts in acetonitrile vs. (i) Ag/AgNOs (o.oi M) and (ii) SCE reference electrodes. (a), N i t r a t e salt vs. (i); (b), n i t r a t e salt vs. (ii); (c), perchlorate salt vs. (i); (d), perchlorate salt vs. (ii), (e), chloride salt vs. (ii). TABLE FORMAL

I REI)OX POTENTIALS

C u ( I I ) - C u ( r) system

Redox potential vs. SCE (mV)

O F Cu(II)-Cu(1) COUPLES IN AC~TONITRILE

Redox potential vs. Pleskov's electrode

Redox potential on the standard hydrogen scale (V)

(my) Perchlorate Nitrate Chloride

95 ° 695 56o

65o 395 --

I. 194 o.939 o.8o 4

Effect of perchlorate on the redox potentials of the copper(II)-copper(I) nitrate and chloride systems The formal redox potentials of the Cu(fI)-Cu(I) nitrate and chloride systems increase considerably in the presence of sodium perchlorate. The increase in potential was found to depend on the concentration of perchlorate present in the solution. Maximum values of 860 mV (vs. SCE) and 840 mV (vs. SCE) were observed with j . Eleetroanal. Chem., I5 (I967) 2oo-2o4

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solutions of copper nitrate and copper chloride, respectively, when saturated with sodium perchlorate (Table 2). Perchlorate ions can replace the nitrate group in copper(II) nitrate to form mixed molecular species of the type: [Cu(ClO4)2-x(NO3)xl, the value of x being determined by the experimental conditions. For example, nitrosyl perchlorate m a y be used to replace one equivalent of nitrate in Cu(NO,)~ to produce the compound, ECu(CIO4)(N03)] 11. On the other hand, if anhydrous perchloric acid is used, 89% of nitrate can easily be replaced b y perchlorate to give a compound of formula: [Cu(CIO4)I.Ts (N03)o.2@ 2. TABLE 2 EFFECT NITRATE

OF

P/~IICHLORATE

AND CHLORIDE

IONS

SYSTEMS

ON" T H I ~ F O R M A L

REDOX

POTENTIALS

OF

COPPER(II)--COPPER(I)

IN ACETONITRILE

Redox system

Concn. of anhydrous sodium perchlorate (M)

Formal redox potential vs. Pleskov's electrode (mV)

Formal redox potential vs. SCE ( m V)

Copper(II)-copper(I) nitrate

o O.OLO 0.025 0.05 ° o.ioo Satd. (~2 M)

395 445 525 545 545 555

693 75° 820 845 845 860

Copper(II)-copper(I) chloride

o Satd.

---

560 840

~i

(~2M)

The addition of a soluble perchlorate to a solution of copper(II) nitrate could give rise to similar molecular species. Such mixed molecular species are likely to be more ionic in character and might therefore dissociate further into the component ionic species than pure copper(II) nitrate and perhaps enhance the cupric ion activities in the system. The replacement of the nitrate group in the cuprous nitrate molecule b y perchlorate ions, is also possible. The anion does not appear, however, to have any effect on the cuprous species since these ions are so strongly coordinated with acetonitrile molecules that they are shielded b y the solvent molecules against the effect of anions. Even cuprous nitrate, which is extremely unstable in water, is quite stable in acetonitrile. The addition of a soluble perchlorate is, therefore, unlikely to change the cuprous ion activity in the system. These two factors could be responsible for the observed increase in the formal potentials of the system in the presence of added perchlorate ions. A similar explanation can be extended to the behaviour of the chloride system. The increase in potential of the Cu(II)-Cu(I) nitrate and chloride systems attains a limiting value when the solutions are saturated with perchlorate. This enhanced limiting potential is, however, lower b y nearly ioo mV than the potential observed with the perchlorate system. As already mentioned, the addition of perchlorate to the copper(II) nitrate and chloride systems brings about the formation of J. Electroanal. Chem., 18 (1968) 20o-204

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mixed molecular species which m a y attain an equilibrium condition. I t m a y not, therefore, be possible for the system to reach the potential of pure perchlorate. This difference in the formal potential of the systems is reflected in their capabilities to oxidize iodide ion, When a solution of iodide is titrated with copper(II) perchlorate in acetonitrile, two well-defined potential jumps are observed, corresponding to the formation of triiodide ion, I3-, and its dissociation into iodine. I n the case of copper(II) nitrate, only one j u m p is registered at the addition of two-thirds of an equivalent of oxidant, corresponding to the formation of triiodide ion. If the I000-

9 O0 -

S ×f--xa

J

800-

700-

)

~---- c

//~.f.~/:/

oo.

-

,o

./J

1

ml

2 o f OXIDANT

3

4

5

Fig. ~. Potentiometric titration curves for the oxidation of iodide by Cu(II) salts in acetonitrile using glass electrode--platinum electrode couple. (a), Cu(II) perchlorate; (b), Cu(II) nitrate; (c), Cu(II) nitrate and NaC104 (o.I ,V/); (d), Cu(II) chloride; (e), Cu(II) chloride and NaC104 (o. I m).

same reaction is carried out in the presence of sodium perchlorate (o.I M) two j u m p s are observed (Fig. 2) as in the case of cupric perchlorate, corresponding to the complete liberation of iodine. The addition of perchlorate is found to have a similar effect on the oxidation capacity of copper(II) chloride. These results show t h a t the addition of perchlorate ion enhances the oxidation capabilities of the copper(II) nitrate and chloride systems.

Department of Inorganic and Physical Chemistry, Indian Institute of Science, Bangalore-z2 (India)

H. C. M~RUTHYUNJAYA A. R. VASUDEVA M~URTHY

1 H, C. MRUTHYUN'JAYA AND A. 1:{. VASUDEVA MURTHY, in press.

2 H, C. MRUTI~YUNJAYAAND A. R. VASUDEVAMURTHY,in press. 3 B. KRATOCHVIL, D. A. ZATKO AND R. MARKUSZEWSKI,

Anal. Chem., 3[{ (1966) 770.

4 H. H. MORGAN,J. Chem, Sot., 133 (1923) 29Ol. 5 CttARLXS]3. DEWI~T, Chemist-Analyst, 24 (4) (1935) 15.

J. Electroanal. Chem., 15 (1967) 200-204

204 6 7 8 9 IO II 12

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I. l~. KOLTnOFF A ~ J. F. COETZEE, J. Am. Chem. Soc., 79 (I957) 1852. F. FARHA AI~D a . T. IWAMOTO, J. Electroanal. Chem., 8 (1964) 55. B. J. HATHAWAY AI~D A. E. U1WDERHILL,J. Chem. Soe., (196o) 37o5 . M. BAnz, V. GUTMAI~W, G. HAMPEL AN'D J. R . MASAGUER, Monatsh., 93 (1962) 1416. G. F. SMITIt AND C. A. GETZ, Ind. Eng. Chem. Anal. Ed., IO (1938) x91. B. J. HATH&WAY, Proc. Chem. Soc., (1958) 344. B. J. HATH&WAY AND A. E. UNDERHILL, ]. Chem. Soc., (196o) 648.

Received April 24th, 1967; in revised form, January I8th, 1968 j . Electroanal. Chem., 18 (I968) 2 o o - 2 o 4

A new carbon paste electrode holder and a simple method for preparing reproducible electrode surfaces. It is difficult to find a method for the pretreatment of solid electrodes that will give reproducible results. Practically every investigator using platinum electrodes has developed an individual method of pretreatment. Nevertheless, it is still very difficult to obtain reproducible results. Wax-impregnated graphite rods can give fairly reproducible results if the tips are cut off or sandpapered before each analysis. MORRIS AnD SCHEMPYz reported peak currents and half-peak potentials at a single graphite rod with mean deviations of only + 1.4% and + 2 mV, respectively. However, this order of reproducibility is uncommon. ADAMS and coworkers ~-5 have developed a very useful type of electrode, the carbon paste electrode. A new electrode, e.g., a fresh paste, can be prepared in a few minutes before each run, simply b y using a spatula. This method gives a variation coefficient of about 6% for the measurements of peak currents 3. With some experience and very careful preparation an operator can obtain a variation coefficient of 2-3% 6. In this paper, a new carbon paste electrode holder is described, together with a method for the preparation of reproducible electrode surfaces with the aid of o.2-mm piano wire.

Experimental To 8 g of graphite (Merck 4207) in a glass dish is added a solution of 6 ml of hromonaphthalene (Fluka AG P 51456 ) in 6 ml of n-pentane. The mixture is stirred with a glass rod for 5-1o mill until most of the pentane is evaporated. The construction details of the paste holder are shown in Fig. I. Note that the inner diameter of the Teflon tip should be about 0.2 mm smaller than that of the metal tube. The paste is tamped into the tube with a glass rod. When the threaded piston rod is turned, the paste is pressed out (O.l-O. 3 mm seems to suffice). Instead of removing the excess paste by smoothing with a spatula 8, the electrode is put into a chuck of a lathe or similar device and rotated (about 600-700 rev./min) while the excess paste is cut off with a stretched length of o.2-mm piano wire. The cutting procedure is repeated with clean portions of the wire until all excess paste is removed. The preparation of a new surface takes about I min. Polaxograms were run b y the conventional electroanalytical method with a J. Electroanal, Chem., 18 (1968) 2 o 4 - 2 o 5