Controlled-potential coulometry and voltammetry of manganese in

Jakob , L. P. Rigdon , and Jackson E. Harrar. Analytical Chemistry 1970 42 ... P. S. Jaglan , R. B. March , and Francis A. Gunther. Analytical Chemist...
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Controlled-Potential Coulometry and Voltammetry of Manganese in Pyrophosphate Medium J. E. Harrar and L. P. Rigdon Chemistry Department, Lawrence Radiation Laboratory, University of California, Livermore, Calif. 94550

A procedure has been developed for the determination of manganese by controlled-potential coulometry. The method is based on the oxidation of Mn(ll) to Mn(lll) at +1.05 to +1.10 V VS. SCE at a platinum electrode, in a supporting electrolyte of 0.25M Na4P20,at pH 2. Samples containing 0.5 to 10 mg Mn may be analyzed with an accuracy and precision of 0.1%. The method is applicable to a variety of materials; among the few interferences are TI(I), As(llI), Ce(llI), and Sb(ll1). Vanadium(1V) interference is eliminated in the sample pretreatment by oxidation to V(V) with (NH,)2SzOs. Excess amounts of Cr(lll) and CI- can be tolerated. The oxidation reaction has been studied in detail by means of voltammetry, coulometry, spectrophotometry, and by ion microprobe mass spectrometry of the electrode surface. The reaction appears to involve an absorbed or electrodeposited manganese species, and the platinum oxide film. At high Mn(ll) concentrations this gives rise to unusual current-time behavior in the electrolysis.

surface, which extends the electrolysis time but has no effect on the accuracy of the determination. This phenomenon has been investigated in more detail by means of voltammetry and other techniques, and the observed effects can be interpreted qualitatively on the basis of a strong adsorption or electrodeposition of a manganese species on the electrode during the course of the reaction. There is also evidence that the platinum oxide film influences the electrode reaction. EXPERIMENTAL

AMONGTHE NUMBER of suitable titrimetric methods for the determination of manganese ( I ) , one of the most versatile is that of Lingane and Karplus (2), in which Mn(I1) is oxidized to Mn(II1) in neutral pyrophosphate medium with potassium permanganate. The method is highly accurate and is relatively free from interferences (2, 3 ) ; however, an unstable titrant is required. Thus it appeared there would be a definite advantage if the oxidation could also be carried out by direct, controlled-potential coulometry. This possibility has been investigated and the conditions necessary for a successful coulometric determination have been delineated. The controlled-potential coulometric procedure developed is simple and compares favorably with titrimetric methods in its tolerance to interferences. Solutions need not be deoxygenated and background corrections are low. The only previously reported controlled-potential coulometric method for manganese (4), involving the reduction of Mn(I1) to Mn(1) at - 1.50 cs. SCE in cyanide solution, requires complicated background corrections, and would be subject to numerous interferences. In the method described below, electrolysis is carried out at a platinum electrode in 0.25M Na4P207. Although the determination is quantitative over a pH range of 1-6 at appropriate applied potentials, a pH of 2 with a potential of +1.05 to $1.10 V os. SCE is optimum for most samples. The method is especially suitable for the determination of milligram and submilligram amounts of manganese. With larger quantities of manganese, the electrolysis current is partially controlled by a reaction involving the electrode

Apparatus. The controlled-potential coulometer used was an operational amplifier-type instrument designed in this laboratory (5), and fabricated by the M-T Electronics Co. (San Leandro, Calif.). The instrument has been improved by the substitution of a Burr-Brown (Tucson, Ariz.) Model 3042 amplifier for both the control and power booster amplifiers in the original potentiostat circuit. The integrator was calibrated electrically and its readout voltages were measured with a Non-Linear Systems (Del Mar, Calif.) Model 484A digital voltmeter. Voltammetric measurements were made with both a controlled-potential, fast-sweep differential polarograph (6), and an ORNL Model Q-1988-ES controlled-potential dc derivative polarograph (7, 8). A special derivative circuit was designed for the faster scan rates employed. Log i us. t curves were obtained directly by the use of a Pacific Measurements (Palo Alto, Calif.) Model 1002 logarithmic converter. Current signals observed with stirred solutions in both coulometry and voltammetry were smoothed with an Analog Devices (Cambridge, Mass.) 1-Hz, active low-pass filter, and were recorded with a Honeywell Test Instruments Div. (San Diego, Calif.) Model 400 X-Y recorder. Two electrolysis cell assemblies for coulometry were used in the course of this work, both based on a design described previously (9). Details of the construction and operation of the improved version of this cell assembly are available (IO). The principal modification is the use of a Coleman Instruments No. 3-702 asbestos-fiber-tipped tube for the reference electrode salt bridge. For the measurement of the log i us. t curves, the reference electrode salt-bridge tube tip was drawn out and placed close to the working electrode to minimize uncompensated resistance. The planar area of 45-mesh platinum gauze in this cell was about 69 cm2. For analytical coulometry, a second cell was built in which the working electrode area was increased by 24 cm2. Both cells had a working solution volume of 25 ml. Both the reference electrode and counter electrode salt-bridge tubes were filled with NaaPz07 supporting electrolyte solution. Beckman No. 39270 and Corning No. 476015 reference electrodes were

(1) M. D. Cooper and P. K. Winter in “Treatise on Analytical Chemistry,” Part 11, Vol. 7, I. M. Kolthoff and P. J. Elving, Eds., Interscience, New York, N. Y.,1961, pp 461-466. (2) J. J. Lingane and R. Karplus, IND.ENG.CHEM.,ANAL.ED., 18, 191 (1946). (3) W. G . Scribner and R. A. Anduze, ANAL.CHEM.,33, 770 (1961). (4) S. A. Moros and L. Meites, J. Electroanal. Chem., 5, 90, 103 (1963).

(5) J. E.Harrar and E. Behrin, ANAL.CHEM., 39, 1230 (1967). (6) F. B. Stephens, E. Behrin, and J. E. Harrar, U. S . At. Energy Comm. Report UCRL-50374, Livermore, Calif., April, 1968. 31, (7) M. T. Kelley, H. C. Jones, and D. J. Fisher, ANAL.CHEM., 1475 (1959). (8) M. T. Kelley, D. J. Fisher, H. C. Jones, ibid., 32,1262 (1960). (9) J. E. Harrar and I. Shain, ibid., 38,1148 (1966). (10) J. E. Harrar, U. S. At. Energy Comm. Rept. UCRL-50417, Livermore, Calif., March, 1968.

758

ANALYTICAL CHEMISTRY

used, and their potentials were checked occasionally against a laboratory-prepared SCE. Voltammetric measurements were carried out by inserting a platinum wire microelectrode into the coulometric cell. This electrode was fabricated from Johnson, Matthey (London, England) spectrographically standardized l-mmdiameter wire and had an apparent geometric area of 0.259 cm2. All measurements were carried out at an ambient temperature of 23 f 1 “C. Other apparatus utilized included a Cary 14 spectrophotometer, a GCA Corp. (Bedford, Mass.) ion microprobe mass spectrometer, and a locally designed electron spin resonance spectrometer. Reagents. For the analytical work, supporting electrolyte solutions were prepared from Na4P207.H20(J. T. Baker) without further purification. Appropriate amounts of salt were dissolved in water, and concd H2S04was added to give solutions 0.25M in Na4P207and the desired pH. Acidic solutions of Na4P20i should be freshly prepared every 3 weeks. Standard solutions of manganese metal were prepared from manganese chips, (Alfa Inorganics, Beverly, Mass.) which were nominally 99.9% Mn with respect to metallic impurities and 99.8y0 Mn, based on total impurities. The chips were etched with 0.1M H3P04,which imparted a bright luster, rinsed with water, and air dried. Weighed portions were then dissolved in 1 M H3P04,and solutions were prepared in certified volumetric flasks to contain 1 or 10 mg Mn per ml in 1 M H3P04. Calibrated micropipets were used to take aliquots of the solutions for coulometric analysis. An auxiliary manganese standard, MnC204 2H20, was prepared by the method of Knoeck and Diehl (11). This salt was dissolved in with ”03, and the solution was fumed to destroy the oxalate before dilution to volume. Solutions of National Bureau of Standards materials also were fumed with H2S04prior to coulometric analysis. The chemicals used in the interference studies were reagent grade or comparable purity. Vanadium solutions were prepared from v o s o 4 , and their V(1V) concentrations were determined by titration with K M n 0 4 standardized against AstOa. For the study of the reduction of Mn(III), cell solutions were deoxygenated with high-purity nitrogen. Procedure. SAMPLEPREPARATION. Large amounts of halogen acids in the sample solution should be removed by evaporation with either H2S04 or HC104. Hydrofluoric acid must be completely removed to avoid attack of the glass cell components ; moderate amounts of hydrochloric acid can be tolerated. If vanadium is present, after removal of the halides, add an excess of (NH4)2S208 to oxidize the vanadium to V(V), then boil the solution for 15 minutes to destroy the unreacted persulfate. Sample solutions in which large amounts of iron are present (>25 mg per ml) should contain between 0.5 and 1.OM H 3 P 0 4to prevent excessive precipitation of ferric hydroxide when an aliquot is added to the supporting electrolyte. ELECTRODE PRETREATMENT. When the working electrode has been used for other determinations, or is being used for the first time, pretreat it as follows. Immerse the electrode in hot concentrated HC1 for 5 minutes, rinse it with water, and place it in the coulometric cell assembly. Polarize the electrode at +1.6 V us. SCE for 5 minutes, then at +0.8 V us. SCE for 2 minutes in pH 2, 0.25M NaaP207supporting electrolyte solution. Replace the cell solution with a fresh portion of the supporting electrolyte and polarize the electrode at the desired analytical control potential for 5 minutes. At pH 2 and +1.10 V us. SCE, the background current should be less (11) J. Knoeck and H. Diehl, Talanfa, 14, 1083 (1967).

7 6

5 4

lp 3 2

ELECTROLYSIS TIMES > 20 min

1 0

i.4

1.3

1.2

1.1

1.0

0.9

0.8 0.7

0.6

0.5

0.4

E vs. SCE (volts)

Figure 1. Practical potential-pH range for controlled-potential coulometric determination of Mn in 0.25M Na4Pz0,

than 30 PA. The electrode is now prepared for the analysis of samples. The treatment with concentrated HC1 followed by anodization and cathodization need not be employed except when high backgrounds or loss of electrolysis efficiency is encountered. Generally, one 5-minute prepolarization at the analytical control potential is all that is necessary for daily preconditioning of the electrode. COULOMETRIC ANALYSIS.Place sufficient 0.25M Na4P207 supporting electrolyte in the cell, so that the final solution volume will be 25 ml, the concentration of Na4P207will be between 0.15 and 0.25M, and the pH of the solution will be 2.0 + 0.5. If HNO, has been used in the dissolution of samples, add five drops of 0.5M sulfamic acid. Pipet an aliquot of manganese solution containing not more than 5 mg Mn into the cell. Electrolyze the solution at +1.10 V us. SCE and measure the integrator readout voltage when the current has decreased to 30 PA. For samples containing chloride or chromium, electrolyze the solution at +l .OS V us. SCE. Determine the background correction by carrying out an electrolysis of the supporting electrolyte alone for the same length of time as required for the manganese sample. At pH 2 and +1.05 to 1-1.10 V us. SCE, this correction should not exceed the equivalent of 0.020 mg Mn for a 20-minute electrolysis. If chloride is present in the manganese solution and the working electrode has not been used previously with chloride-containing solutions, the background correction for the first run will be somewhat larger than that for succeeding runs, because of the establishment of the platinum chloride film on the electrode (12, 1.3). Therefore, reject the result of the first run, and determine the quantity of manganese as well as the blank correction from succeeding runs.

RESULTS AND DISCUSSION Coulometric Analysis. As shown in Figure 1, considerable latitude is allowed in the choice of control potential and pH for the quantitative determination of manganese in pyrophosphate medium. The practical range is limited anodically by the background current of the evolution of oxygen, and cathodically by excessively long electrolysis times. However, because the Mn(I1)-Mn(II1) system is irreversible, quantitative coulometry is still possible at much lower control potentials. (12) D. G. Peters and J. J. Lingane, J. Elecrroanal. Chem., 4, 193 (1962). (13) S. Gilman in “Electroanalytical Chemistry,” Vol. 2, A. J. Bard, Ed., Marcel Dekker, New York, N. Y.,1967, pp 137-139. VOL. 41, NO. 6, MAY 1969

759

Table I. Analyses of Standard Solutions of Manganese by Controlled-Potential Coulometry (0.25M Na4P207)

Control potential, V VS. SCE

Manganese takena, mg

Mean background correction, mg

Mean manganese found, mg

n

Time of electrolysis, min

Re1 std dev, Z

Re1 error,

pH 2.0

0,5354 1.003 4.935 9.814

+1.05 $1. 10 + l . 10

+1.10

0.0050 0.006 0.010 0.015

10 6

0.5351 1.003 4.936 9.877

6

6

15 10

-0.06

30

0.16 0.30 0.16 0.06

10 18

0.23 0.06

0.0 0.0

20

0.0 +0.02

$0.03

pH 6.0

5

1.003 $0.80 0.006 6 4.898 $0.80 0.007 6 Calcd assuming 100.00% purity for the Mn metal after etching.

The effects of pH on the optimum control potentials are in general agreement with the findings of Watters and Kolthoff (24), who measured the formal potential of the Mn(I1)Mn(II1) couple in 0.40M Na4P207. They found that it varies from +0.86 V us. SCE at pH 1 to $0.31 V at pH 6. These investigators (24, 1.5) also established the identity of the complexed manganese species in this medium and gave the following equation for the predominant half reaction in strongly acidic solutions:

+ H4Pz0;F1

Mn(H2P20;)22-

+ 2 H+ + e-

Mn(HzP20;)aa-

~~

(14) J. I. Watters and I. M. Kolthoff, J. Amer. Chem. SOC.,70, 2455 (1948). (15) I. M. Kolthoff and J. I. Watters, IND.ENG. CHEM.,ANAL. ED., 15, 8 (1943). (16) D. M. Yost and H. Russell, “Systematic Inorganic Chemistry,” Prentice-Hall,New York, N. Y., 1946, pp 224-227. 760

ANALYTICAL CHEMISTRY

IWd-

(1)

The solution of pyrophosphate has low buffer capacity in the pH range of 2.5-4.5, thus this region is best avoided for most routine analytical work. Because some solubility problems may be encountered at high pH, and because most manganese samples are dissolved in acid media, a pH of 2 appears to be most generally useful for coulometry. The stability of the supporting electrolyte is of some concern because of the hydrolysis of pyrophosphate to orthophosphate in acid solution (16). Supporting electrolyte solutions of 0.25M Na4P2O7at pH 2 were used in this work for as long as 3 weeks without effect on the analytical results; however, a careful study was not made of this factor. A solution of Na4P20; before pH adjustment is strongly alkaline and is stable for much longer periods of time (16). It is not necessary to control the Na4P207concentration carefully for satisfactory analytical results. Table I indicates the precision and accuracy that were obtained in the analysis of standard manganese metal solutions. To minimize further the background correction, a control potential of f1.05 V us. SCE was employed for the 0.5-mg samples. For an ideal system in which the electrolysis current is proportional to the bulk concentration of electroactive species, plots of log i us. t should be linear. However as shown in Figure 2, such curves for the electrolysis of Mn(I1) assume an unusual shape for amounts greater than 1.4 mg; and with a very sample (18.8 mg) a virtually constant current is observed. Because 1-10 mg of Fe(I1) in 1 M H2S04 can be oxidized in this cell within 7 to 11 minutes without anomalous ~~

1.003 4.898

Figure 2. Log i us. f , controlled-potential electrolysis of Mn(1I)

0.25M NarPzO,,pH 2, E A . Background D. 4.71 mg

= +1.10

V

VS.

SCE

B. 0.47 mg Mn(I1) E. 18.8 mg

C. 1.41 mg

current decay behavior, the effect observed with Mn(I1) cannot be attributed to either improper cell design or potential control, but is apparently due to a kinetically controlled process. In an effort to characterize this phenomenon more completely, additional experiments were conducted as described below. As far as the analytical coulometry is concerned, the kinetic complication imposes a practical upper limit on the amount of manganese that should be taken for analysis. The electrolysis of quantities of manganese greater than 5 mg places a premium on low background and low integrator drift. Although not imperative for this determination, it would be advantageous to use a more efficient electrolysis cell such as that designed for high-speed coulometry (27, 18). However, because of the kinetic factors, a significant improvement in electrolysis rate would be obtained mainly by increasing the electrode area/solution volume ratio, rather than by increasing the effectiveness of stirring. Interferences. Table 11 summarizes the results of tests to determine the effects of various possible interferences. In addition to the species tested, the following are known to be quantitatively oxidized at the control potential and their

(17) A. J. Bard, ANAL.CHEM., 35,1125(1963). (18) G. C. Goode and J. Herrington, Anal. Chim. Acra, 33, 413 (1965).

~~

-30 t l .4

I

I

t1.0

t0.7

E

VI

I t0.4

I 0

I -0.4

SCE (volts)

Figure 3. Voltammetric curves for the oxidation of Mn(1I) and reduction of Mn(II1) in stirred solution 0.25M Na4PZ07, pH 2; scan rate, 100 mV/min A . Residual currents B. 2.93 mg Mn(II) 11.7 X C. 5.86 mg Mn(I1) D. 8.79 mg Mn(I1) E. 2.93 mg Mn(II1)

M]

level of interference can be calculated: Br-, Fe(II), Fe(CN)B4-, Hz02, Ir(III), I-, Pu(III), and Ru(II1). Antimony(II1) was found to poison the electrode in some manner and to extend the electrolysis time; the recommended pretreatment was necessary to restore the electrode to its normal condition. Vanadium is potentially the most serious interference among the elements commonly associated with manganese, because V(1V) is quantitatively and rapidly oxidized to V(V) at $1.10 V 6s. SCE. However, this bias can be eliminated in the sample pretreatment by chemically oxidizing the V(1V) with ammonium persulfate. In the absence of pyrophosphate, phosphate, or silver, no Mn(I1) is concurrently oxidized. Cerium(II1) also is apparently quantitatively oxidized in this medium, and its current-time behavior is similar to that of Mn(I1). Chromium(II1) is slowly oxidized; a 2 : l ratio of Cr:Mn can be tolerated in the analysis at the lower control potential of +1.05 V us. SCE. The influence of foreign ions on the coulometric method is similar in many respects to that reported for the LinganeKarplus potentiometric titration (2, 3). However, there are several differences that arise mainly from the difference in pH of the pyrophosphate media of the two procedures. The acid solution recommended for the coulometric analyses avoids some of the solubility problems (3) as well as the effect of chromium (19) reported for the pH 6-7 medium of the Lingane-Karplus procedure. A number of ions that do not otherwise interfere with the coulometry were tested for their solubility at pH 2. Of these, three were of low solubility: Ba(II), 0.5 mg per 25 ml of supporting electrolyte; Hg(I), 2 mg; and Sr(II), 9 mg. The coulometric technique at pH 2 also circumvents an inaccuracy of the Lingane-Karplus method resulting from the air oxidation of Mn(I1) in the presence of Cu(II), Co(II), and Fe(II1) (20). In the coulometric procedure, samples containing excess acid are partially neutralized, if necessary, before an aliquot is added to the supporting electrolyte, thus no pyrophosphate is present during pH adjustment and air oxidation of Mn(I1) does not occur. (The buffer capacity of 25 ml of the supporting electrolyte is such that 4.6 mmoles (19) W.G.Scribner,ANAL.CHEM., 32,970(1960). (20) W.G.Scribner, ibid., 32,966 (1960).

Table 11. Tolerances of Diverse Substances in the Determination of Manganese (5 mg Mn, 0.25M Na4P207,pH 2.0, E = +1.10 V US. SCE) Amount to cause Substance Added as 0.5% re1 error, mg 0.02 As(II1) AS203 >10 &(I) Bi(II1) Bi(NOd3 >10 0.05 Ce(II1) Ce(N03h Co(I1) CO(N03h >50 Cr(II1) Cr(ClOd3 0.50; 10at +1.05V Cu(I1) cuso4 >50 >100 Fe(II1) Fe(C104h Ni(I1) Ni(N03)~ >50 NH4+ NHaOH > lo00 0.05 Os(II1) oScl3 Pb(W Pb(N03)z > 10 Pd(I1) Na2PdC14 > 10 0.05 R(IU PtClZ Rh(II1) RhC13 > 10 0.03 St$III) SbC13 0.06 TU) TIC1 0.03 V(W VOSOa V(V) NaV08 >25 Anions CINaCl 30; 70 at +1.05 V c104HClO4 > 10 C2042HzCz04 0.05 CHaCOOCH,COONa > 10 NO1 NaNOZ 0,020 NOaNaNO; >60 a Interference eliminated with sulfamic acid.

of H+ are required to lower the pH to 1.5 and 2.0 mmoles of OH- are needed to raise the pH to 2.5.) With solutions of pH 2, the coulometric procedure cannot tolerate as much halide as the Lingane-Karplus method. For samples containing moderate amounts of chloride, a control potential of f1.05 V us. SCE is recommended to minimize the oxidation of C1- to Cln. If larger amounts of chloride are present, the analysis can be carried out at pH 6 and +0.80 V US. SCE. Under these conditions, at least 425 mg of chloride, or a concentration of 0.5M in the cell solution, causes no interference. Similar improvement in the tolerance to bromide, which is electrolytically oxidized at pH 2, but not in the Lingane-Karplus method (3), could probably be realized at pH 6. Analysis of Standard Materials. The results of the determinations of manganese in various standard materials are shown in Table I11 and, in general, are in good agreement with the theoretical or certificate values. The manganese oxalate dihydrate salt has recently been evaluated and recommended as a primary standard by Knoeck and Diehl (11).

STUDIES OF THE MECHANISM OF THE OXIDATION OF Mn(I1) IN PYROPHOSPHATE MEDIUM Voltammetry in Stirred Solutions. Current-voltage curves obtained at a platinum microelectrode with stirred solutions for the oxidation of Mn(I1) and for the reduction of Mn(II1) under identical conditions are shown in Figure 3. Voltammetry was carried out in the coulometric cell, and the Mn(II1) solution was prepared by controlled-potential electrolysis at +1.10 V us. SCE. Kolthoff and Jacobsen (21) also studied (21) I. M. Kolthoff and E. Jacobsen, Microchem. J., 1, 3 (1957). VOL. 41, NO. 6, MAY 1969

761

-10-

-20

-

3

z -a-

a

1

-I

-3

I

I

I

I

I

J

l

l

l

l

l

40

-

-w

-

-bo-

-m +I.$

1 il.4

1

rl.3

,

+l,l

I

+I.I

, +1.0

l

io.? I

I 10.8

, +0.7

, i0.b

I

10.5

,

10.4

+0.1

I,. ICE ( " O l l l l

Figure 5. Regular and first derivative voltammetric curves for the oxidation of Mn(I1) at a stationary electrode 0.25M NaP207, pH 2; scan rate 18.4 mV/sec A. Residual current B. 9.76 mg Mn(I1) 15.7 X 10-3M] C. Derivative of curve B.

the reduction of Mn(II1) at a rotated platinum electrode in pyrophosphate solutions, and the data obtained here are in good agreement with their work. From the present studies, it is evident, first of all, that the Mn(I1)-Mn(1II) system is totally irreversible. Secondly, the apparent limiting current for the oxidation of Mn(I1) is onesixth that of the same concentration of Mn(II1). The limiting current for Mn(I1) is also relatively insensitive to the stirring conditions; thus, the magnitude of the current for the solution stirred at 1800 rpm is only slightly larger than that observed under natural convection conditions. This insensitivity is also shown by the absence of noise on the traces, in contrast to the curves for Mn(II1). Thirdly, the apparent limiting current for Mn(I1) oxidation is not linearly proportional to concentration in the higher portion of the range applicable for coulometric analysis. Curve D of Figure 3 exhibits both a peaked character and an indication of another oxidation process beginning at about $1.1 V GS. SCE. Although the qualitative behavior of the Mn(1I) oxidation is the same over the pH range 1 to 6, these characteristics become more distinct at higher pH. The principal wave shifts cathodically with increasing pH by about 110 mV per pH unit. The peaked nature of the stirredsolution waves results from the fact that a steady-state, mass-

transfer-controlled current is never established. If the applied potential is held at any value more anodic than about +0.9 V US. SCE at pH 2, the current decays steadily. Stationary Electrode Voltammetry. The voltammetric behavior of the oxidation of Mn(I1) in both stirred and unstirred solutions is significantly influenced by the surface condition of the platinum electrode. The curves of Figure 4 show the effect of electrode pretreatment on the characteristics of the oxidation at a stationary electrode in an unstirred solution. In general agreement with what is found for many electrode reactions (22,23),the combined anodic and cathodic polarization places the electrode in its most active state. After such pretreatment, the electrode was found to remain active for several hours in water or in the supporting electrolyte, at open circuit, or while being used in measurements between +0.7 and +1.3 V us. SCE at pH 2. The electrode gradually ages to yield curves such as D of Figure 4 and this condition appeared to be stable indefinitely. As shown by curve C of Figure 4, the treatment with hydrochloric acid renders the electrode very inactive. The coulometric and voltammetric results for stirred solutions described above were obtained with aged electrodes; however, it was noted that there was no qualitative difference in the data for active electrodes. Although cathodization at +0.2 V us. SCE does permit higher electrolysis currents, (22) Reference 13, pp 117-118. 36, 71R (1964). (23) A. J. Bard, ANAL.CHEM.,

Table 111. Controlled-Potential Coulometric Determination of Manganese in Mean Theoretical or certified Mn n Found Material value, Mn Manganese oxalate dihydrate 30.694 5 30.694 NBS 100b, Mn steel 1.89 4 1.903 (Cr 0.06, Mo 0.2)

z

z

NBS 25c, Mn ore

NBS 162, Ni-Cu alloy (Ni 66, Cu 29, Cr 0.2) NBS 162a, Ni-Cu alloy (Ni 64, Cu 31, Fe2)

762

ANALYTICAL CHEMISTRY

Standard Materials Std dev, Z Mn

Z Mn

0.01 0.007

0.00 +0.013

Error

57.85 2.34

4 4

57.83 2.339

0.04 0.005

-0.02

1.60

5

1.631

0.008

+O. 03

0.0

-501 11.5

I

i1.b

I

r1.3

I

11.2

I 61.1

f 11.0

I bo.?

I

~0.8

I

80.7

I 10.6

!

a 5

I

r0.4

1 +O.J

f n. IC8 Ird,,)

Figure 6. Cyclic stationary electrode voltammetric curve for the oxidation of Mn(I1) 9.76 mg Mn(I1) [5.7 X 10-aM], 0.25M N ~ P z O IpH , 2 ; scan rate 10.0 mV/sec

l W l lmlnl

Figure 7. Log i us. t , controlled-potential oxidation of Mn(I1) with interruption of electrolysis 10.0 mg Mn(II), 0.25M NaPzO?, pH 2, E = +1.10 V Electrolysis interrupted for 30 seconds; 30 oxidized

SCE.

background currents were also higher (as shown by curve B of Figure 4); this more than offsets the gain in electrolysis efficiency. Thus, the recommended electrode pretreatment for coulometry is to cathodize at +0.8 V, which results in an electrode condition close to that of the aged electrode. For the oxidation of Mn(I1) in unstirred solution, the peak current is linearly proportional to concentration up to about 6 mg Mn in the cell; but it then increases too slowly with concentration, as in the case of the stirred solutions. Estimates based on the peak currents at low concentrations give a value for the apparent diffusion coefficient of Mn(I1) of 1.4 X 10-6 cm*/sec, and 1.1 X lop6 cm*/sec for Mn(III), assuming n = 1 for both reactions. The variation of peak current, i,, and the half-peak potential, Epi2,was determined as a function of the rate of voltage scan, between values of 0.01 and 1.0 V per second. Faster scan rates were not employed because the wave was increasingly obscured by the residual current. The experimental correlations for this system were typical of an uncomplicated irreversible charge transfer (24), and a value of AEg/2of 60 mV per decade of scan rate was obtained. The stationary electrode current-voltage curves for solutions containing more than 6 mg, or about 4 X 10-aM Mn(II), become distorted and show evidence of a second oxidation process. This is illustrated in Figure 5 , which shows both a regular and first derivative curve for the oxidation of 5.7 X 10-3M Mn(I1). The smaller peak of the derivative curve corresponds to the normal wave observed at lower concentrations of Mn(I1). Experiments in which Mn(I1) was oxidized at pH 2 in the absence of pyrophosphate indicated that the larger peak of the derivative curve (hence, the second process) was due to the oxidation of uncomplexed Mn(I1). An analysis of the behavior of the peak current of curve B in Figure 5 as a function of the voltage scan rate, however, still indicates only an uncomplicated charge transfer. Cyclic triangular wave experiments were also carried out on this system and a typical result is shown in Figure 6. A significant characteristic of these curves, at all concentrations and scan rates, was the presence of a small increase in the oxidation current at about +1.0 V us. SCE on the cathodic portion of the scan. Additional Controlled-Potential Electrolysis Experiments. During regular controlled-potential coulometric runs with the larger amounts of manganese, it was found that, if the electrolysis was interrupted while the electrolysis current was

still above about 10 mA, and the electrode was left at open circuit in the stirred solution, upon continuing the electrolysis, the rate of current decay was accelerated. The log i us. t characteristic during such an experiment is shown in Figure 7. Here the electrolysis current first decays to a virtually constant 15 mA, and a complete electrolysis would normally require over 30 minutes. Interruption of the electrolysis reestablishes the higher current, recapitulating the original current decay, thus decreasing the total electrolysis time. Periods of suspended electrolysis from a few seconds to 20 minutes were identical in effect, and for short hold periods, even the overall elapsed time was decreased. On the other hand, except for the relatively minute charging current, the total charge transferred for a complete electrolysis was not significantly increased. Obviously a chemical reaction occurs during the open-circuit period which permits a more rapid electrolysis of the remaining Mn(I1). The controlled-potential electrolysis behavior of the reduction of Mn(II1) in this medium was also investigated and found to be normal in every respect. At 0.0 V us. SCE, 5 mg of Mn(II1) can be quantitatively reduced in 12 minutes, and the log i us. t curves are linear for all concentrations and applied potentials. Because the characteristics of the oxidation of Mn(I1) indicate possible secondary reactions, voltammetric and spectrophotometric measurements were made on the solution during electrolysis. No transient or intermediate species could be detected in solution; however, in agreement with Buck (25) (who studied the system in HPSOJ it was noted that Mn(II1) is capable of oxidizing the platinum electrode. Stationary electrode curves for the reduction of Mn(II1) showed a small prewave due to the reduction of the oxide film, when the electrode, reduced at 0.0 V us. SCE, had been in contact with the solution for a few minutes. Measurement of the visible absorption spectrum of the Mn(II1) pyrophosphate complex as a function of time after the current interruption, and after complete electrolysis, indicated no slow transformations involving this species. Equilibration of Mn(I1) from the stock solution with the supporting electrolyte for periods of a few seconds to 43 hours revealed no differences in electrochemical behavior, thus it

(24) R. S. Nicholson and I. Shah, ibid., 36, 706 (1964).

(25) R. P. Buck, ibid., 35,692(1963). VOL. 41,NO. 6, MAY 1969

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appears that equilibrium in the Mn(1I)-pyrophosphate complex system also is rapidly established. Although the stability constants for this system are not accurately known, the Mn(I1)-pyrophosphate complex is rather weak. Estimates based on data in the literature (14,26) and ESR measurements in this laboratory yield a value of about 104. Thus at the concentrations of interest in this work, several per cent of the total Mn(I1) exists uncomplexed by pyrophosphate. Electrode Surface Examination. Very early in this work it was realized that the Mn(I1)-Mn(II1) system in pyrophosphate solution might involve the formation of Mn(IV), possibly as MnOz on the electrode. A reaction to be considered is the disproportionation reaction :

2 Mn(II1)

Mn(I1)

+ MnOV)

(2)

Manganese dioxide and Mn(1V) formation are favored by an increase in the pH of the solution and the absence of complexing ligands which stabilize Mn(II1) (2, 27). Accordingly, the deposition of MnOz on a platinum electrode from neutral NaC104 solution has been made the basis of a stripping analysis procedure for manganese (28). Experiments in which Mn(I1) was oxidized at the platinum gauze electrode in Na2S04solution at pH 2, using a control potential of $1.10 V us. SCE, resulted in the formation of a brown film on the electrode. This film, which undoubtedly was MnOz, was not soluble in Na4PZ07 solution alone, but was soluble when the solution contained Mn(II), and gave Mn(II1) via Reaction 2. The only previous investigators of the electrolytic oxidation of Mn(I1) in pyrophosphate medium, Gorbachev and Belyaeva (26, 29), also reported the existence of a brown film on the platinum electrode after their experiments at pH 7. Utilizing a rotating disk electrode, and solutions containing both oxidation states at total manganese concentrations of 0.01 to 0.024M, these authors proposed that the limiting current for the oxidation of Mn(I1) at pH 7 corresponds to a twoelectron process. Even though the coulometric and stationary electrode voltammetric data obtained here certainly indicate that the principal electron transfer reaction at pH 2 is a one-electron process, the novel behavior of the system still strongly suggests electrode surface effects. In an attempt to detect the presence of manganese on the platinum electrode as a result of electrolysis of Mn(II), experiments were carried out in which platinum sheet electrodes were used as the working - electrodes, and then were examined by ion-microprobe mass spectrometry. The electrodes were removed from the solution, without switching off the potentiostat, after a significant fraction of the Mn(I1) had been oxidized. The electrodes were then rinsed with water, analyzed for manganese, and the results compared with those from platinum sheet that had only contacted the electrolysis solution. The spectrometry clearly revealed a gradient in manganese concentration, from a significant amount at the surface of the electrode, to a negligible amount in the bulk of the platinum. Conclusions. Among the numerous cases of current-time behavior treated in the literature, one possible mechanism that has been shown to exhibit characteristics similar to those (26) S. V. Gorbachev and V. A. Belyaeva, Russ. J. Phys. Chem., 37,97 (1963). (27) R. G. Selim and J. J. Lingane, Anal. Chim. Acta, 21, 536 (1959). (28) C. 0. Huber and L. Lemmert, ANAL.CHEM., 38, 128 (1966). (29) S. V. Gorbachev and V. A. Belyaeva, Russ. J. Phys. Chem., 36, 114 (1962). 764

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observed for the oxidation of Mn(I1) is the electrolysis of polymeric species (30). This mechanism can be rejected on the basis of the voltammetric data obtained here, and because of the potentiometric work of Watters and Kolthoff ( I d ) , which shows either that mononuclear species only are involved-or at least that the Mn(I1) and Mn(II1) complexes must involve the same number of manganese atoms. The influence of solution impurities on the electrochemical results cannot be discounted entirely (31, 32), although no change in the observed results was noted when high purity distilled water, recrystallized sodium pyrophosphate, and different sources of Mn(I1) were used. Moreover, the effect of electrolyte impurities would be expected to be felt at low as well as high manganese concentration. Some previous examples of reactions having voltammetric characteristics resembling those of the oxidation of Mn(I1) are the oxidation of various organic substances (33), the oxidation of oxalate (34, 3 3 , and the oxidation of iodine to iodate (36). In each system the reaction has been said to be inhibited by the formation of the oxide film on the electrode. Because the oxide coverage increases with anodic potential, the current as a function of potential goes through a maximum and then decreases (33). The oxidation of Mn(I1) in pyrophosphate medium appears to be a qualitatively similar case of involvement with the oxide film. The single scan voltammetric data and the characteristics of the electrolysis as a function of applied potential show that the rate of the electrolysis reaches a maximum at 1.1 V us. SCE, then decreases slightly at more anodic potentials. This is also manifested by the hump in the cathodic portion of the cyclic scan of Figure 6. It is conceivable that the platinum oxide film also is the cause of the unusual current behavior in the controlledpotential electrolysis. For example a reaction of the type 2 Mn(I1)

+ 2 H+ + PtO

2 Mn(II1)

+ Pt + HzO

(3)

may play a role and Mn(I1) reduction of the film at open circuit may regenerate the electrode surface condition when the current is interrupted. The results of several experiments, however, argue against this mechanism. First, as noted earlier, the reverse reaction has been shown to occur. Second, voltammetric measurements of the cathodic stripping of the oxide film formed at f1.10 V us. SCE were not influenced by the presence of Mn(I1). Third, the current-time behavior in the electrolysis was the same, whether the manganese is added before the control potential is applied, or vice versa. Thus, although the platinum oxide layer is certainly involved in the detailed mechanism of the process, the primary cause for the characteristics of the Mn(I1) oxidation appears to be that a product of the electrolysis, such as MnOz, forms on the electrode during electrolysis, decreasing the rate of the overall reaction. The current interruption behavior could then be due to the dissolution of the MnOz by the Mn(I1) and the supporting electrolyte. It is reasonable that the (30) Y.Israel and L. Meites, J. Electroanal. Chem., 8, 99 (1964). (31j Reference 13, pp 112-117. (321 B. B. Baker and W. M. MacNevin, J. Amer. Chem. Soc., . 75, 1473 (1953). (33) Reference 13, pp 161-189. (34) Yu. A. Korostelin, S. V. Gorbachev, and Z . N. Ryantseva, Rum. J. Phys. Chem., 40, 1023 (1966). (35) J. J. Lingane, J. Electroanal. Chem., 1, 379 (1960). (36) I. M. Kolthoff and J. Jordan, J. Amer. Chem. SOC.,75, 1571 (1953).

higher initial concentrations of manganese, with their higher concentrations of uncomplexed Mn(II), would have a greater tendency to form MnO2. However, the fraction of total manganese present as Mn02 on the electrode cannot be large. An experiment in which the solution was removed from one coulometric cell after 50% of the Mn(I1) had been oxidized and then the remaining Mn(I1) determined in a second cell yielded within 0.1 % of the calculated amount, based on no Mn02 formation. Assuming that an adsorbed product or electrodeposit does inhibit the oxidation of Mn(II), and that part of the deposit is lost at open circuit, the shape of the log i us. t curve at high concentrations of Mn(I1) can be explained as follows. Before the control potential is applied to the electrode, the coverage is such that a relatively rapid electrolysis is possible. When a potential of $1.10 V us. SCE is imposed on the electrode, oxidation of Mn(I1) commences along with the formation of the deposit, and the current decreases more rapidly than normal because product build-up is concurrently inhibiting the Mn(I1) reaction. Eventually the current reaches a value at which the rate of Mn(I1) consumption does not rapidly decrease the Mn(I1) concentration, and a steady state is approached. Finally, the concentration of Mn(I1) decreases to the region where the limiting current density is less influenced by a species on the electrode surface, and at this point the log i us. t curve becomes concave downward, finally becoming linear as in an uncomplicated electrolysis. Interruption of the electrolysis allows dissolution of the surface deposit, again permitting more rapid electrolysis and causing the current-time behavior of Figure 7. Just as the disproportionation Reaction 2 is probably involved in the current interruption, so also must it be a factor in the electrode surface reaction, if Mn(1V) is present. As mentioned earlier, the characteristics of the oxidation of Ce(II1) to Ce(1V) in Na4P20, solution at pH 2 are similar to those of the Mn(I1) oxidation, with an even more pronounced enhancement of the electrolysis current upon interruption and even more distinctive peaking of the voltammetric curves. The maximum rate of electrolysis of Ce(II1) is reached at +0.9 V us. SCE and is considerably slower at $1.10 V. Examination of the electrode by ion microprobe mass spectrometry after electrolysis at +1.10 V [as was done with Mn(II)], also revealed significant amounts of cerium on its surface. On the other hand, vanadium(1V) was oxidized at +1.10 V us. SCE in Na4P2O7solution at pH 2, and the

electrode analyzed by mass spectrometry, but no vanadium was found. The oxidation of V(IV), also totally irreversible, exhibits none of the unusual voltammetric and controlledpotential electrolysis characteristics observed for Mn(I1) and Ce(II1). Although it was beyond the scope of the present work, a further investigation of the Mn(I1)-Mn(II1) system in pyrophosphate medium should include a study of the effect of electrode material, the use of fast potentiostatic measurements, and a more complete study of the reaction as a function of pH. The newer techniques such as the in situ optical examination methods and the use of ring-disk electrodes would probably also be useful. ACKNOWLEDGMENTS

Fred B. Stephens assisted in obtaining some of the voltammetric data. Ronald K. Stump performed the analyses with the ion microprobe mass spectrometer, and Raymond L. Ward carried out the electron spin resonance measurements. RECEIVED for review November 25, 1968. Accepted March 3, 1969. This work was performed under the auspices of the U. S. Atomic Energy Commission. Presented at the 157th Meeting, ACS, Minneapolis, April, 1969.

Correction Identification of Surface Functional Groups on Active Carbon by Infrared Internal Reflection Spectrophotometry In this article by James S. Mattson, Harry B. Mark, Jr., and Walter J. Weber, Jr. [ANAL.CHEM.,41,355 (1969)], the following errors appeared. In the title the word “Spectrophotometric” was substituted for “Spectrophotometry”. On page 356, column one, lines 15 and 18 under Experimental, the words “absorbs” and “absorbed” were substituted for “adsorbs” and “adsorbed”. On page 357, column one, line 16 under Results and Discussion, “H202”was printed for “H20”.

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