Feh., 1953
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THE ANODIC BEHAVIOR OF ANTIMONY BY S. E. S.ELWAKKAD AND A. HICKLING Department of Inorganic and Physical Chemistry, Liverpool University, Liverpool, England Received April 17, 1968
The study of the anodic polarization of antimony in acid and alkaline solutjons, mainly using the cathode ray oscillograph, has revealed the following features: (1) In weak acid and alkaline electrolytes an antimony anode can be rendered passive at moderate current densities; the initial stage in this process is the formation of a relatively thick layer of Sb203 on the electrode, underneath which a thick film of Sb20sis later produced. (2) At very low current densities intermediate.formation of SbZO, can be detected. (3) In strong acid and alkaline electrolytes in which the trioxide is soluble no passivation occurs. (4) Reduction of the anodically formed oxides does not take place on cathodic polarization prior to hydrogen evolution, although deposition of antimony frcm solution may occur.
In a previous studyi it has been shown that the irregularities noticed by various workers in the behavior of the antimony electrode may be explained by the gradual formation of a higher oxide on the surface of the electrode and by its exceedingly slow rate of attainment of equilibrium, which depends upon the supply of oxygen and on the surface of.meta1exposed. In the present investigation the nature of the different oxides formed on the surface of antimony has been further clarified by studying the anodic behavior of the metal in acid and alkaline solutions. This also serves to ascertain which of the numerous oxidation-reduction systems for which potentials can be calculated2 are, in fact, established at working antimony electrodes. No previous work on antimony from the present standpoint has been published and very little information on its anodic polarization is available from other fields. Grube and Schweigardta studied the anodic behavior of antimony in potassium hydroxide solutions by voltammetric methods. They found that at low current densities antimony dissolved to give antimonite (KSbOz) solutions, but that at higher current densities passivity occurred, the electrode becoming coated with a layer of white oxide; antimonite solutions could be further oxidized anodically to antimonate (KSbOa) a t a platinum anode, but the reaction apparently took place indirectly by the action of the oxygen liberated a t the electrode. In the present study it is shown that in weak acid and alkaline electrolytes an antimony anode can be rapidly rendered passive at moderate current densities. The initial stage in this process is the formation of a relatively thick layer of Sb203 on the electrode underneath which a thick film of SbzOs is later produced. Formation of the tetroxide Sbz04can only be detected at very low current densities. In strong acid and alkaline electrolytes in which the trioxide is soluble no passivation occurs. Reduction of the anodically formed oxides does not take place on cathodic polarization prior to hydrogen evolution, although deposition of dissolved antimony from the solution may occur. Both the tetroxide and pentoxide are unstable in contact with the metal, as on interrupting the current after their formation the potential drops very quickly to that of the trioxide. (1) El Wakkad, J . Chsm. Soc.. 2894 (1950). (2) Latimer, “The Oxidation States of the Elements and their Potentials in Aqueous Solutions,” Prentice-Hall, Inc., New York, N. Y., 1988, p. lQ8. ’ (8) Qrube and Sohweigsrdt, Z. ElakCocAim., 98, 257 (1928)~
Experimental The main part of the work was carried out oscillographically a t a moderate curfent density (c.d.) of 0.02 a./cm.2 and the electrical circuit and electrolytic cell employed were essentially as previously described.4-10 The antimony electrodes used were made by deposition of antimony under standard conditions onto a platinum wire electrode of area 0.05 cm.2 sealed into the end of a glass tube. The deposition was carried out by the method of Schoch and Brown11 using as electrolyte a solution made by dissolving 0.1 g. of pure Sb201 2 g. of N2H4.HC1in 20 ml. of concentrated HCl and diluting to 250 ml.; electrolysis was carried out at 60-80” with a c.d. of 0.005 a:/cm.2 for 2 hours using an antimony anode. A fresh plating was used for each experiment. Observations have been made mainly with HC1 and NaOH solutions of various concentrations as electrolytes. Except where otherwise stated, all experiments were made a t 19” in air; the presence of air had prev~ouslybeen found to have no influence on the oscillograms. The results were obtained as photographed oscillograms showing directly the variation of potential with quantity of electricity passed when the antimony electrode was polarized. A 45.4-pf. condenser was used in series with the electrolytic cell and the quantlty of eleotricity normally passed in a single polarization pulse was about 3000 microcoulombs, the time of a single sweep across the screen of the cathode ray tube occupying approximately 3 seconds. Suitable reference lines were photographed immediately after recording the polarization. tracks permitting the construction of accurate potential/quantlty of electricity graphs. Polarization at very low c.d.’s was studied in 0.1 N HCl both by the oscillographic method and by a direct otentiometric method in which the variation of p o t e n d with time was followed over a prolonged period. In the former method the procedure was essentially as described above but the polarizing c.d. was 0.0005 .a./cm.2. I n the direct potentiometric method a large antimony surface was prepared by depositing antimony in a spongy form on a platinum electrode of area 1.8 cm.2; this was polarized‘at a constant current of 100 pa. and the variation of potential with time over a prolonged period observed; the electrical circuit and electrolytic cell have been previously described .I2
+
Results Behavior in Acid Solution.-In Fig. 1 is shown the characteristic curve for the anodic polarization of antimony in 0.1 N HCI at 19” with a polarizing c.d. of 0.02 a./cm.2. The oscillogram changed somewhat at fist with the duration of the pulsating electrolysis, but became steady after two minutes, indicating that the antimony surface had attained a stable effective area! and this time was therefore adopted in recording all the oscillograms. It may be seen from h g . 1 that the potential of the antimony electrode at first. rises very rapidly from the hydrogen evolution value whlch is (4) El Wakkad and Hickling, Trans. Faraday Soo., 46, 820 (1950). (5) El Wakkad and Salem, THIS JOURNAL, 56, 621 (1952). (6) Hickling, Trans. Faruday Boe., 41, 333 (1945). (7) Hickling, ibid., 40, 518 (1945). (8) Hickling and Spice, ib4d.. 48, 762 (1947). (9) Hickling and Taylor, Disc. Faraday SOC,,No. 1, 277 (1947). (10) Hickling and Taylor, Trans. Faraday Soc., 44,261 (1948). (11) Schoch and Brown, J . Am. Chem. SOC.,88, 1660 (lQl6). (19) El Wakksd and Emun, J . Chrm. Soe., 461 (1961).
S. E. S. EL WAKKADAND A. HICKLING
204
the starting point of the track; this rapid rise is followed by a short definite step and this in turn by a second arrest, the electrode then becoming passive with slow rise of potential toward that of oxygen evolut,ion. By measurement of a large number of oscillograms with different series capacities, the commencing potentials of the two steps were found to 0.37 v., respectively, be fairly well defined a t +0.10 and on the hydrogen scale. Increase of c.d. was found to shorten the steps and to decrease the quantity of electricity required for passivation; this is shown in Fig. 2 where polarization tracks at 0.01 and 0.05 a./cm.z in 0.1 N HCl are given. Increase of temperature to 70' had no appreciable influence on the anodic polarization track in 0.1 N HCI. I n Fig. 3 are shown the anodic and cathodic polarization tracks in 0.1 N HC1 a t 19" with a c.d. of 0.02 a./cm.2; the very unusual feature is here found that on cathodic polarization the potential drops almost at once from the oxygen to the hydrogen evolution value, and there is no sign of any reduction processes corresponding to the steps in the anodic polarization track. I n more concentrated hydrochloric acid solutions the antimony electrode did not become passive so readily, as is shown in Fig. 4 which records the anodic and cathodic tracks at a c.d. of 0.02 a./cni.Zin N HC1. Here it is seen that some anodic process sets in a t ca. 0 v.,, the potential slowly rising during a prolonged arrest to about +0.15 v., and then the track resembles the beginning of that in 0.1 N HCl albhough wibh the quantity of electricity passed the potential does not reach any limiting value. In contrast to the observation in 0.1 N HCI, the cathodic track here shows a definite arrest at -0.25 v. before the potential drops t,o that of hydrogen evolution. With increase of acid concentration these features became more marked; thus in 5 N HCl at the same c.d., an anodic process occurred continuously at a potential slightly higher than 0 v., the electrode showing no sign of passivity during the polarization pulse. and the cathodic track showed a reduction process occurring a t only a slightly more negative potential, this being ultimately followed by hydrogen evolution,
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s
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F
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d
vi-1.5-
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-OS-
-1.
1.5-
I
,2
3
COULOMBS 10:~ Figs. 5-8.-Oscillograms for antimony in alkaline solution.
-
1.0
0.5-
us -
5 0-
-054 'I
'2
3
I
'2
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COULOMBS Figs. 1-4.-0scillograms for antimony in acid golution. Behavior in Alkaline Solution.-In Fig. 5 are shown the anodic and cathodic tracks for the polarization of antimony at 0.02 &./em.* in N NaOH; the anodic track indicates some oxidation process occurring continuously in the vicinity of -0.75 v., while the cathodic track shows a reduction step in the region of -1.25 v. followed by hydrogen evolution. With decrease of alkali concentration the antimony electrode showed signs of passivity on anodic polarization. Thus in Fig. 6, which ahows the anodic polarization curve in 0.1 N NaOH at 0.02 a./cm.S, it ia seen that there are two
Measurement of the Oxide Potentials of Antimony.Schumann18 measured the potential of the Sb/SbzOa system and obtained the value of +0.152 v. a t unit hydrogen ion activity. Direct measurements on the other oxide systems appear to be lacking although some values have been deduced from thermal data.2 I n the present work direct measurements of the steady potentials set up by all the stable oxide systems in 0.1 N HC1 and N NaOH have been made so as to facilitate comparison with the arrests noted in the oscillograms. B.D.H. pure samples of powdered Sb, SbzOs and SbzOs were used in making up the systems; SbzO4 was prepared by heating Sb205a t 50?-60O0 t? constant weight, the decrease in weight on heating a eeing with that for conversion of Sb205 to SbzO! within 1%. The general procedure in making the potential measurements was as follows: the two components of each system were agitated with the electrolyte by means of a continuous stream of nitrogen, and the potential ofia short platinum wire electrode immersed in the system was measured against a saturated calomel reference electrode by means of a valve potentiometer; the system was maintained at 20' in a thermostat, and the variation of potential with time was observed until a steady value was attained. The results are given below, expressed on the hydrogen scale. System
Sb/SbsOa SbzOa/SbnOp SbzO4/SbzOs SbzOs/Sb&s
E uilibrium potential v 0.1 ?€IC1 N kaOH fO. 09 (f.48)
-0.63
(+.53) +.61
(-.17) .07
( - .24)
-
Addition of powdered Sb to any of the higher oxide systems rapidly brought the potential to the Sb/SbnOo value. The (la) Sohumann, J. Am. Chsm. Boo., 4 % 5 2 (1824).
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THEANODICBEHAVIOR OF ANTIMONY
valucs for the systcins containing SbzO4 are enclosed in pttrtmtheses, since, although t,he potent.ials recorded were t tie steady values obtained experimentally, it seems daubtful whether they can represent true equilibrium potentials for the systems specified. Thus it would be expected that the ShzOa/SbzOspotential should be the arithmetic mean of the Sb203/Sb204 and Sbz04/Sb20s values, which is not the case, and the values found for systems containing SbzO, are not in good agreement with calculations from thermal data. I t seems possible that thest: discrepancies may be due to antimony t,et,roxide failing t80give in solution any directly corres onding ions, since its chemical nature is still not clearly dekned."
205
to the initial oxide layers a t anodes such as platinum, gold and nickel which have been previously studied.6J~8 Thus it would appear that the antimony trioxide film has only poorly protective properties and must grow to a very appreciable thickness before the anode potential can rise substantially. In strong acid and alkaline solutions, in which the trioxide is soluble and gives the antimonyl SbO+ and antimonite SbOz- ions, continuous dissolution of the antimony anode occurs; i n circumstances where the oxide is less soluble, Discussion initial dissolution may be followed by formation Under conditions such that the antimony eler- of the oxide layer. The nature of the second arrest in the anodic trode can become passive the anodic polarization curve shows two arrests both in acid and i n alkaline tracks is of great interest. This begins a t average solut,ions. The first step commences at +O.lO v. values of +0.37 v. in 0.1 N HCI and -0.37 v. in 0.1 N HC1 and a t -0.57 v. in 0.1 N NaOH, the i n 0.1 N NaOH, the uncertainty in these values uncertainty in these mean values being not greater being about -10.02 v. By analogy with the pre, - lwould ~ . ~ ~be expected than hO.01 v.;. the equilibrium potentials for the vious metals s t ~ d i e d ~ , ~ it that this would correspond to an equilibrium posystem Sb/SbzOBin these two solutions are +0.09 a i d -0.57 v., respectively. The close agreemelit) tential between higher and lower antimony oxides. of these two sets of values leaves little doubt that Reference to the table of experimentally determined the first step corresponds to the formation of equilibrium potentials, however, shows that there .antimony trioside upon the anode. The quantity is no agreement between these potentials and any of electricity passed in the first step varies some- of those listed, and the discrepancies are much what with experimental conditions, but a t the usual greater than any likely experimental error. The c.d. of 0.02 a./cm.z it is approximately 13000 potentials of the unstable systems Sb/SbzOc microcoulombs/cm.2 in 0.1 N HCl, 0.1 N NaOH and Sb/SbzOb have therefore been calculated from aiid N NazCO3 electrolytes; this would correspond thermal data. From the heats of formation and to the liberation of 41 X loL5atoms of osygen. estimated entropies of the two oxides, their free From the specific gravity of antimony of 6.68, energies of formation have been calculated2 as the diameter of the antimony atom may be cal- - 165,900 and - 195,500 cal., respectively. From d a t e d to be 3.1 X em., and hence there would these values the potentials (in volts) are readily be 1.03 X 10I6 atoms/t>rue cm.2 of metal surface. derived and are 0 1NHCl 0.1 N NaOH The real surface area of the antimony electrode will undoubtedly be greater than its apparent area but $0 27 -0.42 Sh/SbzOi no very satisfactory method exists of finding the f O . 32 -0 37 Sb/SbzOa relation between these two quantities. To obtain a rough idea of the ratio accessible area/apparent The second pair of values are in fair agreement with area for the antimony electrodes used, the double the potentials for the commencement of the second layer capacity of the metal prepared in the stand- step in the anodic tracks, and this suggests therefore ard way when used as a cat'liode in acid solution that the second step corresponds to the formation was measured and compared with the value for of :Ltitimony pentoside i n direct contact with the mercury under the same conditions; these measure- niitimony surface. This state of affairs might be ments suggested that the real area of the antimony e\pec*ted if the antimony trioside film initially electrode was about 3 times its apparent area. formed is a poor electrical conductor; electrolysis This figure may be substantially in error since the may then proceed through cracks in the film, which basic assumptions of this method of obtaining as has been shown is of a poorly protective characcessible areas are not entirely free from criticism acter, with consequent increase in the effective and the capacity measurements present consider- c.d. at the antimony surface leading to oxidation able difficulties, but if it is tentatively adopted it of the metal to the higher state. Some support for would suggest that there are about 3 X 10'5 atoms this view was afforded by the study of the decay of of antimony/apparent cmS2 on the electrode. anodic polarization in 0.1 N HC1 which was carried Hence the ratio atoms O/atoms Sb in the first out with a 0.5-cm.2 antimony electrode polarized step would be about 13, and since the oxide formed at 1 ma. for periods varying from 10 seconds t o 10 has been shown from potential measurements to be minutes, the current then being switched off and Sbz03, this would indicate that the oxide layer is the variation of potential with time followed oscillosome 9 molecules thick. No precise quantitative graphically. The results were in all cases substansignificance can be attached to this calculated tially the same and are illustrated by one example i n thickness, since as previously pointed out the Fig. 9. On interrupting the current it would be quantity of electricity passed in the step varies expected that the pentoxide should react with with experimental conditions and the ratio real/ metallic antimony to give the trioxide, and the apparent, area may be substantially in error, but the decay curve shows a step in the general region of the general inference is quite definite that the anti- Sb2O3/Sb2O5potential before the potential finally mony trioside layer is relatively thick in contrast drops to somewhat below the Sb/SbsOs value. The quantity of electricity passed in the anodic (14) Terrey, Ann. Rep. Chem. Soc., 36, 119 (1938).
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S. E. 5. EL WAKKAD AND A. HICKLING
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antimony tetroxide had been observed in the I
,
,
IO
20
, 30
,
,
,
40
50
(rD
, TO
,
80
MINUTES.. Fig. 10.-Anodic
polarization of antimony at very low current density.
oxide reduction steps, but on the contrary the potential drops immediately to that of hydrogen evolution. This again would fit in with the vie4 that the oxide layer has little intrinsic electrical conductivity, and that on cathodic polarization hzdrogen ions pass through the pores in the layer and are discharged on the metal underneath; the oxide layer may then strip off or be reduced concurrently with hydrogen evolution. I n more strongly acid electrolytes, where the anodic tracks indicate dissolution of antimony, then reduction steps are found in the cathodic tracks, and the potentials a t which these occur suggest that they correspond merely to deposition of antimony from the solution. In the alkaline solutions some dissolution of the oxides appears to occur in all cases and the cathodic tracks show reduction steps at very negative potentials; these steps were found experimentally to be extended by addition of antimonite to the electrolyte, and hence probably correspond to reduction of the antimonite ion to the metal.
