J.C.S. Faraday I, 1980,76,2587-2603Some Reactions at a Mercury@) Sulphide PhotoanodeBY R. STEPHEN DAVIDSON"? AND CHARLES J. WILLSHERDepartment of Chemistry, University of Leicester, Leicester LE 1 7RHReceived 9th July, 1979The pho t oelec t rochem ical react ions of pigmentary mercury(I1) s ul p hide have been investigatedusing electrochemical cells having a platinum electrode coated with the sulphide. E.m.f. and currentmeasurements show that the sulphide behaves as an n-type semiconductor and when sodium nitrateis used as electrolyte the photoanode is stable. Use of other electrolytes can lead to solubilisationand a change of colour of the sulphide. The darkened form of the sulphide so produced is morephotoreactive than the red form. The sulphide photoassists the electrolysis of water and this isrationalised on the basis of an energy level diagram drawn up from experimental data.We have previously described the preparation of platinum electrodes coveredwith pigmentary titanium dioxide and the photoelectrochemical properties of suchelectrodes.The ease of preparation of such electrodes and the fact that a pigmentarysample of the semiconductor rather than the more expensive single crystal can beused to carry out photoelectrochemistry suggested that techniques developed in thiswork could be applied to other semiconductors with the hope that a material maybe found which is suitable for transducjng visible light into electrical energy or asource of fuel. The requirements for such a material are (a) it absorbs visibleradiation, (b) it is photostable and (c) it either has a good power output or electrolysesor photoassists the electrolysis of water.We now report upon the photoelectro-chemistry of mercury(I1) sulphide which fulfils, in part, the above criteria.EXPERIMENTALMercury(@ sulphide (vermilion) (Koch Light) was used as received. Other sources ofmercury(i1) sulphide included B.D.H., Fisons and May and Baker.In nearly every case the sulphides showed a negative photo-e.m.f. and displayed a photo-anodic current. Meta-cinnabar (Alfa Products) was used as received.Electrolyte solutions were made up using de-ionised water and the salts were of the highestpurity available.The loading of mercury(I1) sulphide is approximately 0.025 g. Experiments were carriedout with a single compartment cell fitted with a cover and provision for ckoxygenation of theelectrolyte by purging with nitrogen.The cell contained three electrodes, the coated platinumelectrode, a platinum counter electrode and a reference electrode (saturated calomeleIectrod.e, s.c.e.). Potentials were measured with a Philips high impelmcc voltmeter(niodel PM 2434) and currents with either a Heath polarography module (EUA-19-2) or aWenking potentiostat (LB 75L).The lightbeam was focused on one face of the covered platinum gauze. For studying the spectralresponse of the electrodes, the light was filtered by use of broad band interference filters(Balzers) and for studying the effect of light intensity, neutral density filters were interposedbetween the cell and light source.t Present address : Chemistry Department, City University, St.John Street., London EC4 4PB.2587The magnitude of these values varied from source to source.The prcparation of the electrodes has been previously described.2*The light source and solution filter system have been previously de~cribed.~In these cases, water formed the heat filter2588 REACTIONS AT MERCURY(I1) SULPHIDE PHOTOANODEIrradiation of suspensions of mercury(I1) sulphide in electrolytes was carried out with aPyrex jacketed water-cooled 125 W medium pressure mercury lamp (Hanovia).Analysis of electrolyte solutions for solubilised mercury was carried out with a Perkin-Elmer 360 atomic absorption spectrometer. The lowest reliable detection limit was7.5 p g ~ r n - ~ .RESULTSRED MERCURY(II) S U L P H ~ D E I N SODIUM NITRATE SOLUTIONIrradiation of platinum electrodes coated with red mercury(I1) sulphide produceda negative photo-e.m.f., the magnitude of which depended upon the source of thesulphide.Usually it was between - 100 and -200 mV. The photoresponse wassimilar to that reported for titanium dioxide, i.e., on commencement of irradiationthe e.m.f. rapidly built up to its maximum value which was maintained until termina-tion of irradiation when it slowly decayed until the dark e.m.f. was attained. Themagnitude of the photocurrents and the onset of relative anodic photocurrentsdepends upon the source of the sulphide. Fig. 1 illustrates photocurrent againstapplied potential at pH 7.The onset of photocurrent usually lay between -0.1 5and -0.35 V (against s.c.e., pH 7) ; -0.2 V was taken to be the flat-band potential.The effect of change in pH of the electrolyte upon the potential required to causecurrent to flow and upon the photo-e.m.f.s is shown fig. 2 and 3. The onset ofrelative anodic photocurrents, the dark potential and potential upon illurnination(against s.c.e.) all vary by z -0.06 V per pH unit although the relative photo-potential is pH-independent between pH 2 and pH 12. In very acidic and veryalkaline electrolytes, dark currents are large and relative photopotentials diminished.The spectral response of the photovoltage and photocurrent are shown in fig. 4.The effect of light intensity upon photovoltage and photocurrent was determined byinterposing neutral density filters between the light source and the cell and the resultsare shown in fig.5.In order to test for photoinduced decomposition of the sulphide, electrodes wereirradiated for several hours under bias and the electrolyte analysed for mercury ions(Hg2+) by means of atomic absorption spectroscopy. Within the pH range 2-12no solubilisation could be detected. electrons were passed anodicallythrough a sulphide electrode on illumination. If two electrons can solubilise onemercury atom, according to1.2 x 10'HgS 4 Hg2+ + S +2ethis would lead to M 2 x g of mercury in solution. The volume of electrolyteemployed was 150 cm3, so theoretically a mercury concentration of 13 pg ~ m - ~should be detected.A value of 0-1 , u g ~ m - ~ was observed. Usually the mercurysulphide electrode showed a 1-2 % weight loss after irradiation. This small loss isattributed to handling.Occasionally, irradiation of red mercury sulphide electrodes in nitrate solutioncaused the sulphide to change to a brown-black col~our. This colour change isdiscussed in more detail later on. When an anodic bias was applied to the sulphideelectrode gas bubbles were sometimes formed at the sulphide and platinum counterelectrode. Gas evolution was more noticeable if the sulphide electrode had becomedarkened on illumination. Gas evolution was also favoured by increasing the pHof the solutions. Since varying the sodium nitrate concentration between 10 tomol dm-3 had little effect upon the photocurrents, the photocurrent is attributedto the photoassisted oxidation of hydroxyl ions at the surface of the sulphide anR .S . DAVIDSON A N D C . J . WILLSHER 2589Ix x-XXXX XI I I I I I I I- 0.8 -0.4 8.0 0.4applied po tential/V against s.c.e.FIG. 1.-Relative photocurrent against applied potential for red HgS in neutral 0.1 mol dm-3NaN03 (N,-purged, unbuffered). The bias setting is manually adjusted and the dark current allowedto settle before illumination. The flat-band potential is considered to be -0.15 to -0.35 V, fromthe photocurrent onset.0 4 8 12PHFIG. 2.-Approximate onset of anodic relative photocurrent at red HgS against pH of N2-purged0.1 mol dm-3 NaN03 electrolyte.The pH is varied by NaOH or HN03 addition and the voltageat which no net photocurrent flows was determined by sweeping the potential at 1 mVs-l. Thegraph has a slope of M -0.07 V pH-'. The form of the photocurrent-bias plot at all pHs investi-gated is similar to that shown in fig. 3. Reference slope at -0.059 V pH-l2590 REACTIONS AT MERCURY (11) SULPHIDE PHOTOANODEPHFIG. 3.-Variation with electrolyte (0.1 mol dm-3 NaNO,) pH in dark potential (0) and potentialupon illumination ( X ). The relative photovoltage is pH-independent although its magnitude variesfrom sample to sample. Reference slope at -0.059 V pH-l.- 6 0 1 Ic- -60 -I-20 -400 500 6 00wavelength/nm400 I 500FIG. 4.-Wavelength response of photocurrent (solid line) and photovoltage (dashed line) for redHgS in 0.1 mol dm-3 NaN03 obtained by irradiation with 1.8 kW xenon lamp with various broadband interference filters and 11 cm of H20 in the light pathR .S . DAVIDSON AND C . J . WILLSHER 259 1the reduction of protons at the counter electrode. The minimum applied potentialto achieve gas evolution at the platinum counter electrode (with respect to thesaturated calomel electrode) was found to be +0.2 V.RED MERCURY@) SULPHIDE I N ELECTROLYTES OTHER THANSODIUM NITRATEIt was found that darkening of red mercury(I1) sulphide leads to an increase inphotocurrents. The use of electrolytes other than sodium nitrate was investigatedto see if the darkening process was affected by ions in the electrolyte. Use of thefollowing salts as electrolytes induced darkening : potassium ferrocyanide, iodate,iodide, thiocyanate, bromide, chloride, fluoride, cyanide and lead nitrate.Potassiumsulphate, like sodium nitrate, behaved capriciously. The use of many electrolytescaused solubilisation of mercury, e.g., potassium oxalate, ferrocyanide, bromide,iodide, iodate, thiocyanate, sodium thiosulphate, hydrogen-phosphate, stannouschloride, ferrous sulphate, manganous nitrate, thallous nitrate and cerous nitrate.Solubilisation occurred when mercury sulphide electrodes were irradiated in theelectrolyte solution and also when suspensions of mercury sulphide in these electro-lytes were irradiated (see table 1). When some ions having Eo values betweenz -0.5 V and + 1.0 V (against s.h.e.) are present as electrolyte ions, they competewith hydroxyl ions and/or water for reaction with the photogenerated holes.Largerphotocurrents were noted in these electrolytes than for a nitrate electrolyte of thesame pH. This is taken as evidence of successful Competition for holes. Someresults are presented in table 1.Usually the photo-response of the mercury sulphide electrode was independent of the type of electrolyteand was similar to that for 0.1 mol dm-3 sodium nitrate provided the pH of theIn all cases the electrolyte solutions were 0.1 mol dm-3.(a) dintensity (arb. units) (b) light intensity (arb. units)FIG. 5.-(a) The dependence of relative photovoltage of red HgS in 0.1 mol dm-3 NaN03 on lightintensity. At high intensity the photopotential saturates.(b) The dependence of relative photo-current (at 0.0 V against s.c.e.) of red HgS in 0.1 mol dm-3 NaN03 on light intensity2592 REACTIONS AT MERCURY(II) SULPHIDE PHOTOANODEsolution is taken into account. Potassium ferrocyanide and thiocyanate proved tobe the exceptions in that the use of these compounds as electrolytes led to the pro-duction of positive photo potentials (see table 2). That this peculiar behaviour isdue to the electrolyte and not to a change in the mercury sulphide was shown by thefact that when the electrodes were removed from the ferrocyanide or thiocyanatesolutions, washed with distilled water, dried and then irradiated in 0.1 rnol d ~ n - ~sodium nitrate solution they showed the normal photoproperties.(a) [KI]/niol dm-350 0400mI300DuLon 1 c200100000000 0000I0 ' 0XXXXXx xXXI x I 1 I I I I I I n20 40 60 80FIG.6.-(a) Solubilisation of mercury upon irradiation (solid line) and in the dark (dashed line) fromred HgS in various KI concentrations. Irradiation is by a 125 W medium pressure Hg lamp(Hanovia) and the suspension is 2 g of HgS in 500 cm3 of stirred KI (open to the air), for 5 h.(b) 2 g of red HgS in 500 cm3 of 0.1 mol dm-3 KI ; solubilised mercury from irradiation of stirredsuspension with 125 W Hg lamp against time. The solubilisation in the dark was FZ 6 pg ~ m - ~ foreach run. The solubilisation rate is M 4 p g C M - ~ h-ITABLE SOME RESULTS FOR RED HgS ON PLATINUM IN SELECTEDThe effective irradiated area is 2.25 cm2 and the light intensity m 2.5 xrelativephoto-relative currentphoto- at 0.0 V photocurrent no.of electronsb [Hg]electrolyte (aqueous voltage against onset passed anod.ically electrolyte0.1 mol dm-3) pHa /mV s.c.e./A against s.c.e. upon irradiation / p7,77110.711.3599.512.27- 75- 200- 200- 125- + 80+ 120- 180- 150- 170-115- + 5- 200+ 70 + 15+3+4- + 100- 1+2f 5+4+6 + 50+2- 0.35- 0.25- 0.25(negative+ 0.05+ 0.05of - 0.45)- 0.65-0.13- 0.35- 0.25- 0.75- 0.252 . 6 ~5 . 8 ~ 1 0 I 81.2x 10185.5 x 101 74.5 x 10' anoc'.ically5.8 x 10'' cathodically2.1 x anod.ica1ly2.6 x loi6 cathodically2 . 6 ~ 1 0 1 73.1 x 10178 .5 ~ 1 O I 68.9 x 101 73 . 2 ~ 10"w 1 0 I 8mmmG3G3a Electrolyte is unbuffered. This column does not include electrons contributing to any currentHgS in 500 cm3 of electroIyte, stirred and irradiated for 240 h by a 125 W medium pressure Hg lamp.corresponding suspension stirred in the dark for 240 h. Readings below reliable detection h i t oTABLE 2.-cURRENTS AND POTENTIALS FOR A RED HgS ELECTRODE IN AN ELECTROLYTE OF VARIEDindicates intial dark potential and " after " means dark potential afterpotassium potential upon relative relative onset potentialt hiocyanate illumination photo- photocurrent of photo-/mol against s.c.e. against s.c.e. /mV against s.c.e. against s.c.e.- 270 (black)1 .o +255 (red) + 290 + 460 - 9.5 > +0.3 V- 170 (black)The sign of the photoeffect depends upon [SCN-] and mercury solubilisation is noted for [KSCN]coccentration dark potential /mv potential /FA at 0.0 V current/V-5 + 155 (red) - 150 + 120 - 3.5 w +0.2s 0-1 + 255 (red) - 105 + 35 - 0.8 + 0.05+235 (red)- 60 (darkened) - 125 - 65 + 2.4 -0.12s 0-3 + 145 (before) - 140 - 150 + 3.2 w -0.1 + 10 (after)10-4 + 180 (before) - 105 - 175 + 1.3 w -0.1 + 70 (afterR .S . DAVIDSON A N D C. J . WILLSHER 2595TABLE 3.-vARIATION IN SOME PARAMETERS (IRRADIATION TIME, [KI]) FOR DARKENING REDHgS AND SOME RESULTS OF DARKENED ELECTRODESOptimum blackening occurs for [KI] = 0.1 mol dm-3 and time of irradiation = 30-180 min,with the HgS held at 0.0 V against s.c.e.~~ _ _ _ ~ _ _ _ ~ ~ _ _ _ _relative relativephotocurrent photocurrentsrelative blackened of blackenedphotopotential form/pA, form/pAconditions for blackening of blackened at 0.0 V bias at +0.4 V biasred HgS form/mV against s.c.e.against s.c.e. remarksirradiated" for 30 rnin atopen-circuit in 0.1 mold~n--~ KIirradiated" for 30 min at0.0 V (against s.c.e.) in0.1 mol dm-3 KIirradiated" for 60 rnin at0.0 V (against s.c.e.) in0.1 mol dm-3 KIirradiated" for 180 min at0.0 V (against s.c.e.) in0.1 rnol dm-3 KIirradiated" for 300 min at0.0 V (against s.c.e.) in0.1 rnol dm-3 KIirradiated" for 30 min at0.0 V (against s.c.e.) inmol dnr3 KIirradiated" for 30 rnin at0.0 V (against s.c.e.) inlo-' rnol dm3 KIirradiated" for 30 rnin at0.0 V (against s.c.e.) in1.0 mol drr3 KIirradiated" for 60 rnin at0.0 V (against s.c.e.) in1.0 mol dmW3 KI- 165 f 5- 230 + 18- 195 + 18- 140 + 27- 65 +9- 155 + 5- 155 + 6.5- 230 +3- 150 + 3.5+ 13+ 36+ 28+ 12.5+ 35+ 15+ 31+ 22+ 15HgS darkbrownHgS darkbrownHgS blackHgS blackHgS blackHgS slightlydarkenedHgS dark redHgS blackHgS black" Irradiation is by the xenon lamp uia CuCl, filter.Performed in pH 11.5 0.1 mol dmA3NaN03 (N,-purged, unbuffered with xenon lamp irradiation through CuCl,. Lightintensity w 2.5 x lov2 W cm-2)2596Use of potassium iodide as electrolyte led to rapid darkening of the sulphide andalso to solubilisation. The rate of solubilisation as a function of irradiation timeand concentration of potassium iodide was investigated using stirred suspensions ofthe sulphide.The results are shown in fig. 6(a) and (b). The effect of potassiumiodide upon mercury sulphide electrodes was also investigated. To do this, theelectrodes were irradiated in potassium iodide solution and then removed, washedand dried and their photoresponse measured using 0.1 mol dm-3 sodium nitrate aselectrolyte. The results are shown in table 3. By controlling the irradiation timeand concentration of potassium iodide, electrodes can be prepared which are farmore photoresponsive, as judged by the value of the photocurrent, than the redmercury(r1) sulphide electrodes. The parameters which influence optimisation of thephotoresponse also include the wavelength of the irradiating light, the number ofelectrons passed across the electrolyte-mercury sulphide interface and iodide con-centration.A comparison was made of the effectiveness of halide ions in darkeningthe sulphide and it was found that iodide > bromide > chloride.REACTIONS AT MERCURY(II) SULPHIDE PHOTOANODEBLACKENED RED MERCURY@) SULPHIDE I N ALKALINE0.1 mol dm-3 SODIUM NITRATE SOLUTIONFor these experiments a blackened electrode was prepared in the following way.A red mercury(r1) sulphide electrode was irradiated for 60 min with the xenon lampusing copper(I1) chloride as filter solution and 0.1 mol dm-3 potassium iodide aselectrolyte. The voltage of the mercury(r1) sulphide electrode was held potentio-statically at 0.0 V with respect to a s.c.e.The treatment caused a weight loss in theelectrode. The electrode was washed with distilled water and dried. The photo-response of electrodes prepared in this way was examined in 0.1 mol dm-3 sodiumnitrate solution at pH 11.5. Usually the photovoltage of such electrodes is similaror larger than that of the red mercury(I1) sulphide electrode. The blackenedmercury(r1) sulphide electrodes attain level photovoltage faster than the red mercury(I1)sulphide electrodes. The photocurrents of the blackened electrodes can be up toten times greater than those that are attained with red mercury(I1) sulphide electrodes.Introduction of the darkened electrodes to the alkaline sodium nitrate solution causesno mercury solubilisation (within the limits of detection by atomic absorptionspectroscopy).Plots of the variation in photopotential with pH and of the change inapplied potential required to cause an anodic photocurrent to flow with pH aresimilar to those shown in fig. 2 and 3. However, in both cases the slopes of the linesare -0.03 V pH-l, i.e., half that observed for red mercury(I1) sulphide electrodes.The spectral response of the blackened electrodes differs from that of the redmercury(1r) sulphide electrodes in that it extends beyond 700 nm, there being nopositive photovoltages at sub band-gap wavelengths.Irradiation of the blackened mercury(i1) sulphide electrodes and application ofa biassing potential 3 +0.2 V (against s.c.e.) causes gas evolution at the platinisedplatinum counter electrode. That the gas evolved is hydrogen was shown by thefact that it almost immediately precipitates palladium from dilute aqueouspalladium(I1) chloride solution^.^ Gas evolution from the mercury(I1) sulphideelectrode is very slow.This may be due to the fact that mercury(r1) sulphide adsorbs~ x y g e n . ~ The adsorption of oxygen may account for the fact that irradiation of thedarkened electrodes for long periods leads to a decline in photocurrent. If suchpassivated electrodes are removed from the electrolyte, dried and then put back inthe cell they behave in the normal way after re-establishment of the electrolytR. S . DAVIDSON A N D C. J . WILLSHER 2597junction, i.e., the passivation is not caused by chemical change of the surface of themercury(I1) sulphide.AUTHENTIC B LACK MERCURY ( I I) S ULP HI D E (metf2-C I N NAB AR)I N 0.1 mol dm-3 SODIUM NITRATE SOLUTIONIrradiation of platinum electrodes coated with meta-cinnabar did not generate ane.m.f.Application of a biassing potential of +0.1 V (against s.c.e.) producedminute photoanodic currents. Stirring a suspension of meta-cinnabar (2 g) in500 cm3 of 0.1 mol dm-3 potassium iodide solution led to mercury solubilisation( 3 0 0 ~ g c m - ~ after 5 h). Illumination of the suspension with a 125 W mediumpressure mercury lamp increased the rate of solubilisation (1000 pg ~ m - ~ in 5 h).Solubilisation was observed when sodium nitrate was used as electrolyte. Irradiationof a suspension of the sulphide (2 g) in the electrolyte (0.1 mol dm-3, 500 cm3) for5 h gave a solution containing mercury (200 pg ~ m - ~ ) .DISCUSSIONIt has been shown that platinum electrodes can be covered with pigmentarytitanium dioxide and that illumination of the oxide injects electrons into the platinum.6Electrodes covered with mercury(I1) sulphide in powder form behave in a similarfashion, i.e., the sulphide is acting as an n-type semiconductor.A disadvantage ofelectrodes prepared in this way is the inhomogeneous nature of the covering. Theuncovered regions of platinum give rise to short circuiting and reduce the poweroutput of the cells. This short circuiting made it impossible to perform capacitancemeasurements and so obtain a direct measurement of the flat-band potential anddonor density of the red mercury(r1) sulphide. Another problem with these electrodesis that the contact between the sulphide and platinum may not ohmic.The build up of e.m.f.to a constant value on irradiation and its decline to thedark potential on termination of irradiation followed a similar pattern to thatobserved with titanium dioxide. As with titanium dioxide the build up of photo-currents on commencement of irradiation was dependent on the type of electrolyte.'Consistent results were obtained with sodium nitrate as electrolyte. That this saltdid affect the photoelectrochemistry of mercury(i1) sulphide was shown by the factthat the photoresponse of the electrode was not affected when the concentration ofthe nitrate solution was varied from 10 to 1 xInspection of fig.1 shows that the onset of the anodic photocurrent in 0.1 mol dm-3sodium nitrate solution occurs when the applied voltage is in the range -0.15 to-0.35 V (against s.c.e.). This onset potential is linearly related to the pH of thesolution (fig. 2) as is the photo-e.m.f. of the sulphide electrode. These two parameterschange in a similar way with change in pH. This suggests that illumination of thesulphide electrode under open circuit conditions leads to band flattening and con-sequently the potential required to cause the onset of the anodic photocurrent canbe taken as the flat-band potential.From fig. 3 it can be seen that the dark potential and the potential produced uponillumination vary by x -0.06 V pH-', i.e., the Nernst equation is obeyed.Therelative photopotential is invariant with pH. The fact that the dark potential andpotential produced upon illumination are dependent upon pH to the same extentshows that the conduction and valence bands at the semiconductor-electrolytejunction shift equally to more negative potentials as the electrolyte pH increases.Since the dark potential is a measure of the Fermi level in the dark (against s.c.~.)'the position of the Fermi level at any pH can be potential calculated and put on anmol dm-32598energy scale. The photopotential attained with intense radiation should correspondto the flat-band potential.In order to utilise these values to construct an energy level diagram for redmercury(I1) sulphide the energy of the conduction band was to be determined.Butlerand Ginley have described a method which is based on a knowledge of the solid-stateproperties of the constituent atoms.' Thus the energy of the conduction band ofmercury(1r) sulphide may be calculated as follows.For sulphur : 1st ionisation potential = 10.36 eV and electron dlinity = 2.1 eV.9so,REACTIONS AT MERCURY(II) SULPHIDE PHOTOANODEx(Hg) = +(1.54+ 10.43) = 5.98xq) = 3(2.1+ 10.36) = 6.23a * * X(HgS) = JxfHgS)X(S) = 6*11where x is the electronegativity.level.The undoped Fermi level of mercury(I1) sulphide lies 6.1 1 eV below the vacuumSince the suiphide (red) has a band gap of 2 eV loEA(HgS) = 6.1 1 - +2= 5.11 eVwhere EA is the electron affinity.controlled by the band gap.one uses the equationThe photoaction spectrum (fig.4) confirms that the wavelength response isTo relate the electron affinity value to solution energy levels (redox potentials)E = a-qVwhere E is the energy on the vacuum scale in eV, a is a constant, q is the electroniccharge and V is the potential in Volts in the standard hydrogen electrode (s.h.e.)scale. The constant a has not been definitively fixed and values of -4.48 and-4.73 I 2 have been used. Use of the latter value gives the energy of the conductionlevel (EcB) as +0.37 against s.h.e.Fig. 7 is an energy level diagram drawn up from the values given in table 4.Owing to the uncertainty in the value for the electron affinity for sulphur andthe constant relating the vacuum and s.h.e. scales, these values could vary by kO.3 V.The energy level diagram satisfactorily explains why only a small biassing potentialis necessary to cause the photoassisted electrolysis of water and indicates thatreducing agents with Eo redox negative of the valence band top can be thermo-dynamically oxidised on illumination.TABLE 4.-vALUES OF CONDUCTION BAND BASE (EcB), VALENCE BAND TOP (&IS), FLAT-BANDPOTENTIAL (EFB) AND FERMI LEVEL (&(DARK)) (FOR SEMICONDUCTOR AQUEOUS ELECTROLYTEJUNCTION IN THE DARK) ON VARIOUS SCALESscale ECB EVB EFB(approx) EF(DARK)(approx)vacuum + 5.10 eV +7.10 eV +5.18 eV + 5.38 eVs.h.e., pH = 0 +0.37 V +2.37 V +0.45 V +0.65 Vs.c.e., pH = 7 -0.29 V +1.71 V -0.20 v 0.0 R .S . DAVIDSON AND C. J . WILLSHERtED(cathodic 1E; E; - - - - - -t‘D(anodic12599EC-FIG.7 . 4 ~ ) Postulated band edges at the flat-band condition for red HgS at pH 7 (against s.c.e.)and their relative positions to some solutions redox couples and &(cathodic) and EC(anodic), thethermodynamic potentials for HgS+ 2e + Hg+ S2- and Hg2- + S + 2e + HgS, respectively. 0.15 Vis the energy gap between the semiconductor Fermi level and solution redox level, taken as half waybetweenE(H+/H2) and E ( o ~ / H ~ o ) for an aqueous solution containing no added redox species. (b)Bandbending [0.15 V from (a)] of red HgS. 0.35 V corresponds to the energy needed to raise the conditionband surface edge to the H+ reduction level and 1.15 V is the difference between the 02/H20 levelon the valence band top.Photosolubilisation of HgS, as indicated in table 1, takes place in a number ofelectrolytes, but probably does not occur by anodic oxidation :HgS = Hg2++S+2e.Passage of two electrons should solubilise one mercury atom and z 10l9 electronsneed to be passed to produce a detectable 10 pg ~ m - ~ by atomic absorption spectro-metry.Some electrolytes show [Hg] > 10 pg ~ m - ~ for < l O I 9 electrons passed anda lot of mercury solubilised in suspension where only photochemical and not photo-electrochemical reactions can occur. The “ inert ” sodium nitrate electrolyte withred and blackened sulphide has been employed to pass sufficient current to givedetectable mercury concentrations in solution, but none has been seen. We thereforeconclude that solubilisation is photochemical and any contribution from anodicoxidation is slight, if occurring at all.The relative positions of decompositionlevels l3 with band edges on fig. 7 confirm this, cathodic decomposition is not likelysince ED(cathodic) < Ec, but anodic dissolution is possible as ED(anodic) < E,, butredox couples with Eo < &)(anodic) can compete with the decomposition rea~ti0n.l~Photochemical solubilisation occurs when the sulphide lattice is ruptured by thereaction of a photogenerated positive hole with a reducing agent, but no current ispassed through the lattice. This is the situation for an irradiated suspension.Photoelectrochemical solubilisation (anodic oxidation) is the process of lattice ruptureby passage of electrons through it into an external circuit.The use of potassium ferrocyanide and thiocyanate as electrolyte produces positivephotopotentials.A similar situation arises when platinum electrodes covered withtitanium dioxide are irradiated in potassium iodide, potassium thiocyanate an2600 REACTIONS AT MERCURY(II) SULPHIDE PHOTOANODEpotassium ferrocyanide solution^.^^ It would appear that in all these cases theanions are being oxidised at the semiconductor and the oxidised form of the ions arereduced at the base platinum sites of the semiconductor covered electrode. Thisprocess has been termed " short-circuiting ". The presence of reducible ions in thevicinity of the inhomogeneously covered platinum covered electrode will give riseto a positive photovoltage.The use of potassium iodide as electrolyte caused solubilisation of the redmercury@) sulphide. This was accompanied by the production of high currentsand also a change in the colour of the sulphide to a brown-black colour.The degreeof darkening is dependent upon the duration of the irradiation, the light intensity andthe wavelength of the light. As yet the mechanism of the darkening and the chemicalcomposition of the darkened material are not known. This is a problem which hasintrigued chemists,16 mineralogists and artists.'However, the darkened form of mercuryfrr) sulphide is far more photoreactivethan the red form as shown by its capacity to deliver higher photocurrents. Anothergreat advantage as far as utilisation of the semiconductor for harnessing solar energyis concerned is that it responds to light of A > 700 nm.So far we have not beenable to find the wavelength at which the photoresponse commences. An unexplainedfeature of these blackened electrodes is the effect of changes of pH upon their per-formance. The finding that the variation in photopotential with pH is -0.03 VpH-l is most unusual and suggests different ion adsorption on the blackened formof red mercury(I1) sulphide.The observation that authentic black mercury(I1) sulphide (meta-cinnabar)generates only a slight photoeffect and undergoes considerable photosolubilisationin sodium nitrate electrolyte leads to the conclusion that very little " meta-cinnabar "is formed on darkening. The darkened form is not unstable in a nitrate electrolyteand its nature remains uncertain.X-ray powder photographs and opto-acousticspectra of blackened red mercury(I1) sulphide suggest some " meta-cinnabar " maybe present, although mercury is also solubilised in the iodide-induced darkening, butwe cannot state the exact nature of the blackened form.We thank the S.R.C. for a maintenance grant to C. J. W.R. S. Davidson, R. R. Meek and R. M. Slater, J.C.S. Faraday I, 1979, 75,2507.* R. S. Davidson and C. J. Willsher, British Patent 7913420.R. S. Davidson and C. J. Willsher, Nature, 1979, 278, 238.Comprehensive Inorganic Chemistry, ed. A. F. Trotman-Dickenson (Pergamon Press, Oxford,1973, vol. 1, p, 8.L. I. Grossweiner, J. Phys. Chem., 1955, 59, 742.R. M. Slater, Ph.D. Thesis (University of Leicester, 1975).(a) M. A. Butler and D. S. Ginley, Chem. Phys. Letters, 1977, 47, 319 ; (b) M. A. Butler andD. S. Ginley, J. Electrochem. SOC., 1978, 125, 228.V. I. Vedenyer, L. V. Gurvich, V. N. Kondrat'yev, V. A. Medvedev and Y . L. Frankevich,Band Energies, Ionisation Potentials and Electron Afinities (Edward Arnold, London, 1966).F. Lohmann, 2. Naturforsch. A, 1967, 22, 843.R. Gomer and G. Tryson, J. Chem. Phys., 1977, 66,4413.' S. R. Morrison, The Chemical Physics of Surfaces (Plenum, New York, 1977), p. 269.lo W. H. Strehlow and E. L. Cook, J . Phys. Chem. Ref. Data, 1973, 2, 163.I3 W. M. Latimer, Oxidation Potentials (Prentice Hall, Englewood Cliffs, N.J., 1938), p. 166.I4 (a) H. Gerischer, J . Vac. Sci. Technol., 1978, 15, 1422 ; (b) A. J. Bard and M. S . Wrighton,J. Electrochem. SOC., 1977, 124, 1706 ; (c) H. Gerischer, J. Electroanalyt. Chem., 1977, 82, 133.H. H. Chambers, R. S. Davidson, R. R. Meek and R. M. Slater, J.C.S. Faraday I , 1979, 75,2517R . S . DAVIDSON A N D C . J . WILLSHER 2601l6 C. Brosset, Naturwiss., 1936, 24, 813.l7 W. H. Cropp, Proc. Austral. Inst. Mining Met,, 1923, 52, 259.l8 R. L. Feller, Nat. Gallery of Art, Report and Studies in History of Art (U.S. Govt. PrintingOffice, Washington. D.C., 1967), p. 99.(PAPER 9/1072