AD-AII6 05 YORK UNIV OOWNSVIEW (ONTARIO) DEPT OF CHEMISTRY F/S 7/4NAPPING OF THE ENERGY LEVELS OF METALLOPHTHALOCYANINES VIA ELEC--TC(U)JUN 82 A B LEVER, S LICOCCIA, K MAGNELL NOOO 4-78-C-0592

UNCLASSIFIED TR-23 NL

mEhEEhEhIRM

4i'I OFFICE OF NAVAL RESEARCH

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iTECHNICAL REPORT NO. 23

' MAPPING OF THE ENERGY LEVELS OF METALLOPHTHALOCYANINES VIA

ELECTRONIC SPECTROSCOPY, ELECTROCHEMISTRY, AND PHOTOCHEMISTRY

BY

A.B.P. Lever , S. Licoccia, K. Magnell, P.C. Minor and B.S. Ramaswamy

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Mapping of the Energy Levels of Metallophthalo- Aug. 80 - April 81

cyanines via Electronic Spectroscopy, Electro- 6. PERFORMINGONG. REPCRTNUMBER

chemistry, and Photochemistry.

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A.B.P. Lever , S. Licoccia, K. Magnell, P.C. Minor N00014-78-C-0592and B.S. Ramaswamy.

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Phthalocyanine, Eleetrochemistry, Photochemistry, Photocatalysis

20. ABSTRACT (Contfou., an rover** oldsi $necesar n'pwd isionvtr &j black mmsbeo)4 The mapping of the energy levels In metallophthalocyanines is accomplished by acombination of electrochemistry, electronic spectroscopy, and photochemistry.This chapter reviews the elc( Lrochemi:al properthvs of metallophthalocyanintsand inrludes a large arm)unt ,f previously unpublished data. The results ar'rationalized in terms of thet nature of rhe electron transfer, that is, redoxat metal or ligand. Well-defined correlations are shown to exist between theease of oxidation or reduction of the l)phthalocyanine ligand and the oxidationste and/or Doar izin2 n,,wjvr. of the L, al ion. Solvent and riui suhst ut in

DO Jo 1473 EDITIoN OFI NOV45 1S O'SOLETE I "T 0

SIN 0102-014- 6603 IUnlassifiedSECURITY CLASSIFICATION Of THIS PAGE (men~ Date &nft-m

20. continued:-

effects also are presented and explained. Charge transfer transitionenergies can be calculated directly from these data, and agreementbetween experiment and theory is excellent. Finally, the data are usedto calculate both the photo-generated excited state redox energies andthe thermodynamics of quenching by donors and acceptors.

Accos~I:,,- r' .

NTIS (ata I

.Di r 1 '.at 4on

~ Av'. 4'..• , .-~ Codes

DLt special .....3

Mapping of the Energy Levels ofMetallophthalocyanines via Electronic

Spectroscopy, Electrochemistry, andPhotochemistry

A. B. P. LEVER. S. LICOC(IA. K. 1\(;NIti P. C. NIINOR, mid 13. S.WIMAS\VA MY

York Universit), Departmient of Chenistry, 4700 Keele Street, Downsview,Ontario, Canada, N13J 11):3

The napping of the energt levels in inetallo-

phthalocyalibiles is accoIn plished by a combination ofclectrochemistry, electronic spectroscopy, and photo-chemistry. This chapter rcie ws the e!ectrochemicalproperties of m etallophthalocyanines and includes alarge amount of previously unpublished data. The re-suits are rationalize'd iI t'rmins of the nature of the elrc-tron transfer, that is, redox (it metal or ligand. Well-defined correlatiows are shouwn to exist between the easeof oxidatior or reduction of the ph, thalocya ni te ligandand the oxidation state, andor polarizing power, of theiietal ion. Solvent and ring substitution effects also are

Pr'su ted and explai,'d. Char.,e tra ,sfer tra usitionencrgies 'au be calculated directly from these data, anldagreelielt between e'xperimettt and theory is excellent.Finally, the data are used to calculate both the photo-gencrated excited state redox energies and the ther-modynainics of quenching by donors and aeceptors.

P hthalocyanine (MPc) complexes have significant importance formany reasons, including their similarity to the biologically impor-

tant metalloporphyrins, their classical use as dyestuffs, and their de-veloping use as components of various solar energy conversion de-vices.

Of paramount importance in understaniding and predicting thephysics and chemistry of the metalloplithalocyaninues is knowledge ofthe energy levels therein. This knowledge can be gained throughstudy of the electronic spectroscopy (absorption and emission), elec-trocheinstry, photochemistry (and photophysics), and photoelectronspectroscopy of these species. In this chapter we address the first threeOf these techtli(ucs and consider the infhrniation obtained fromtheir use.

The lower excited states (m) to it least 35,000 cot ') of a widerange of iietallophithialocyaiiinec derivatives can he identified andmiapped. These states vary as a fitoftotimi of the environment (solvent,axial coordination, etc.), phithalocvanine ring substituent, nature andsize of metal ion, oxidation state, and electronic configuration.

The excited states may be classified as 7r-7r* and n-lr* transitions,primarily located oil the phthalocyanine ring. The ligand (Pc ring) tometal charge transfer (LNICT), illetal to ligarcd charge transfer(NILCT), and d-d transitions primarily occur on the metal atom. Thesetransitions liay occur in two spin manifestation|s. In addition, spincoupling can occur between metal ion ground state wave functions andexcited state wave functions of the phthalocyanine ring to yield arange of triplet-multiplets.

The following three sections review progress in the elec-trochemistry, electronic spectroscopy, and photochemistry of metal-[ophthalocyanines. The data reported for metallophthalocyanines maybe compared usefully with the data obtained for the porphyrins dis-jissed in Chapter XX in this volume.

Elect rochemis try

Many reports (1-18) discuss the electrochemical propertiesof the metallophthalocyanines, but only in recent literaturehave these data coalesced into a useful working body of know-ledge.

In general, oxidation and reduction are expected at the metalcenter and at the phthalocyanine ring. In each case, one ormore electron transfer processes may be observed. Indeed, twosuccessive ring oxidations and up to four successive ring red-uctions may occur. Ring reductions are generally electro-chemically reversible, but ring oxidations usually are not(at least on platinum). Generally, no more than two redoxprocesses were characterized at a given metal center.

Identification of the nature of a given redox product usuallyis based on electronic spectroscopy and, where relevent, electronspin resonance (ESR) spectroscopy (2,9,12). Generally, it ispossible to deduce unequivocally whether a reduction or oxidationoccurred at the metal center or phthalocyanine ring. Phthalocyanineanion and cation radical electronic spectra are quite distinct fromthose of the Pc(-2) species, (12,19) although some confusion mayexist when low-oxidation state transition-metal ions, such asiron(l), are involved (9).

The higher filled and low(r empty energy levels of a typicalmetallophthalocyanine are illustrated in Figure 1 (20-23). Thetwo highest filled orbitals of relevance to our discussion areTr orbitals with alu [highest occupied molecular orbital (HOMO)]anda 2 symmetry, respectively; and the lowest empty ringorbitals are the eI lowest unoccupied molecular orbital (LUMO))andb 2u r* -orbitals. Tie metal valence orbitals may be buriedinside the filled phthalocyanine levels, or filled and/or emptyvalence orbitals may occur in the HOMO-LUMO gap; in addition, emptymetal orbitals may lie at energies comparable to, or above, theLUMO phthalocyanine level.

Phthalocyanine redox chemistry may he classified conveniently intotwo sections: main groups and transition groups.Main Group Phthalocyanlne Electrochemistry. Redox chemistry in

the main groups is (ui-ml ly quite ';traightforward; generally themetal atom (center is unaffected, and all observable processesoccur on the plithalocyanine ring. For main group ions that liein the phthalocyanine plane, the first ring oxidation (from HOMO)is separated from the first ring reduction (to LUMO) by approximately1560 mV (Table I), which is the magnitude of the molecular bandgap.ThLs valm, n ;o't'n.s largo Iy tmn;,fffect-I. by thl' natirc of the main groupmetal, although some deviation may occur it the metal is too largeto be accommodated by the phthalocvanine center (14).

-3-

However, the absolute energies of first oxidation or reductionvary considerably, and depend on the size and charge of the metalion. The ease of oxidation or reduction of the phthalocyanineunit depends on the electric field generated by the central metalion. Indeed, there is a clear relationship between the polarizingpower of the central ion, expressed as (ze/r) and the redox energy.

The more polarizing is the ion, the easier it is to reduce thephthalocyanine ring, and the more difficult the ring is to ovidize.

A good linear relationship is observed between these quantiLies (14)

defined by Equations 1 and 2:

oxidation (ze/r)(E" - 1.410) = -0.01 (1)

reduction (ze/r)E" + 0.145) = -0.012 (2)

where the potentials are referenced to NHE and the radii usedoriginate from Shannon and Prewitt (24). The lines are essentiallparallel. These data imply a HOMO-LUMO separation of ahout 1.56 Vindependent of the central main group ion, provided it lies in th,plane (14). This treatment is discussed latr.

The second relutioi i~roiess ap)pears, oil the a\'eraaie, about 420ill%

' l, l. ll e I tati\( thall the fil'st itct till (faihle 1) hit tilec i)otelitials

tic lu1011 'catlt'i't .tntl and I .ss dilt'ltl dtelid(e t oil the polarizint!titwer l' tht. celti'al ioll; ute\ tltul(,cs ltt.llo|hthLlt'i(tutuh iC5 w'thlie- tinore poltr/.iog iois arc gtnural inore readily reduced to the

iing dianion. "'The third- and hourth-rcdtction processes occur near- 1.7 to - 1.8 and -2.0 V (Table 1) (12) (for h)oth inain and transition

Uj.) ions). Tlhese dtta are displayed in Figure 2. For the first twore' ,ction processes, the ease of r(duiction follows the sequence:

')c(. 2 -- Pl g(II. -\ 19( 1lrINh > l'cAI-(II)X

1i. SO(ettHllu _ lla'y hV ,atioidiZ d in terms of increasing negativecl'arge on the plidalhc\ anilne Je1and when passi)g from, the covalentNII boind il Pc 12, to it i fl\ (leC trOstatiCt interaction with the elec-tropositi , uiii.t ' igicsilltl, M id finally to a till] le ati e charge with the

fl) 1,1a 11t 61.1 m ,.tl slt. For aluinluiti( llI), the high polarizing powerI,|tti, lou leadis to a lre d islcerned.cIl any it rlte Ilegatie

oarge oil the placaluryanin lial. These potentials nioitor theweegati\e charge oi the prthalo(n (ie Iiga. u g t

Transition Goup Plithaloyanine Electrochemistry. The ires-one (or oa trasitist ioital io, :le ars to perturb the re(lox processes,-ccturritg at the lithalocyanlile ulilit however, a pattern simiflar to the

tain grotlp species cal ne discerned. l nany cases, ole or adorete lx processes daty ocit'r ait the e'etral iton t pItenltials l ighee-weell rien g oxi ( li an d redis ctioti, u lui' sp e,l t, .tlpurtin g electro-lyvle, or a ,oddhd ligtand can biaod to tte axial sites of the metal ion inIte (or olre) of its xiatiol levels, thel withe observed reIox potential0thit eods Ilarkud] o. the teoir., of Solvent, electrolyte, or addedhigand.

Table 11 I i\' vs data for tra,)sitioll llital ion lihthahocyanines where(Ol\'ent et)ordllatioll is no)t stloogly peltttrlhin , that is, ill wea'k donor

:'Olvtents, or il,\olv.ilg iletals tilat b~ind weakly alotig the axis. Thesedata shoidd bv, getier.ally inlk-rlprefabit, withouit tatkiiig into considera-

,Noll severe perttir!,ationb 1 '. Sovn "fi ... ...ILJ 'e d '] ''.,. . ..

Electronic and ESH spectroscopy demonstrated that the OTi(IV),

OV(IV), Ni(ll), Cu(II), and Zni(l) species do not undergo redox pro-ctsses at the metal at potentials between liganl oxidation and reduc-tion. Iron and cobalt, on the other hand, can frin M(1), M(II), andM(III) species at these intermediate poteiiial, that is, oxidation of tileIlhthalocyalihie ligand occurred alter the inlal was oxidized to M(III),and reduction of the pithalo-yanine ligand occurred only after reduc-tion of the metal to M(I). Chromium and manganese phthalocyaninesforn M(II) and M(III) oxidation states (5-13).

In parallel with main group phlihalocyanine chemistry, the abilityto reduce a netalloplithalocyaninc increases, that is, the potential be-comes more positive, as the oxidation state of the central ion increases.This ability can be seen froi the data collparisons abstracted fromlable 11 and shown in Tables III, 1 and Figure 2.

Not surprisingly, the potentials are similar to those of main groupions of tile same oxidation state and ap)roxinlate size, altholugh thislact a)j)artIiltly was niot rectOgimise(l clearly in the pii-st. licatuse thespread in potentials for a given itietal oxidation state is remarkably

iliall, aiid te ir isa clear eiiugliCgh SVMlr-tiOil ietween tite ranges far atleast the first and second reduction, the potentials generally call beused diagnostically to identify the oxidation state of the central transi-tion metal ion.

Most oxidations of the phthalocyanine ring in transition metalplhthalocyanines are electrochemically irreversible, obscuring oxida-tion state trends. Trivalent and tetravelent transition metal phthalo-cyanines generally oxidize a little above 1.0 V; this trend is alsotrue for the more polarizing divalent ions, nickel(II) and coppei(II).Pc(-2)Zn also oxidizes near 1.0 V while the earlier first row transitionmetal phthalocyanines oxidize at slightly more negative potentials.

Solvent effects on these ligand redox potentials are small. How-ever, solvent effects on metal redox process potentials can be extraor-dinarily large. Table V gives the ranges for various redox processes as afunction of solvent and/or supporting electrolyte. The effect of solventdepends clearly oin the electronic configurations of the species in-volved, and the effect of supporting electrolyte depends on whetherthere is binding of the anion to either component of the couple.

The iron(II)/iron(I)and cobalt(lll)/cobalt(II)couples both involvelow spin ci"/d- configurations. Strongly binding axial ligands (solventmolecules) destablize the z' electron (in d7 ) and favor oxidation to thelow spin dP species. Thui, ill both eases the potentials shift negativelywith increasing dlolor strength of the solhent, which follows tile order

i)MA I)MF )I)MS() < pyridine (3)

\vhere DIMA, l)\M, ,o id I)\IS() represent dimethylainine, diinethyl-lornamide. and di miethl\ I sulfOxide, respectively.

Indeed, there is a linear correlation of these potentials with Gut-inann Donicity Nn niber (2,5) of the solvent (8). The cobalt(ll)/cobalt(I)couple (low spin d'/d ") also shitis negatively with increasing donicityof tile solvent, probably because axial biniding to the square planar d'cobalt(l) is weak or nonexistent. The iron(lII)/iron(II) couple (low spinFdoP) however, shows the opposite trend shifting positively with in-creasing donicity of the solvelit (8, 9). This fact best can be explainedby synergism, where stroingl coordinating axial ligands favor backdollnation by the( low spill d'i ion into phthahocvanine ir-acceptor orbi-tals. Back doatit lol in iroli( 1 is weak.r becauI'se of the greater charae'lll the Ihetal h,le o t lt li t. oIlveut elicts call be quite

hmami.it( il. ib m .1iillmp . I. ullult i ol of i( ti il) p tlmialoyainine in~ mil iiu. onI mi~iiill huH 1(11 io as i sth. il'l,., itl a similar solution in

I)NA oi DMi with chiomc iiil rapidly air oxidizes to iron(Il Ihil ii ~ ~ ~ ~ ~1) i i * hK n i iiiii ii i" . .. o.....'....

Moreover, if cyanide ion is added to asolution of tetrasulfonated iron(lI) phthalocyanine, the iron(Ill) state is

stabilized tot a remarkable degree (Table V) through axial coordinationof cyanide ions. Indeed, unsubstituted iron(lI) phthalocvanine is solh-

ie in water if cyanide ions are present (5). Even more remarkablestabilization of iron(lt1) is seen when mi idazole is osed as an axialligand (Table V) (7).

When a series of substituted pyridines were used as solvents (8),both the cobalt(IlI)/cobalt(Il) and cobalt(li)/cobalt(I) couples shiftednegatively with increasing pK. of the solvent. This result may be ex-

plaill'd inl terms of the l)rago E and C niodel (26) givenl that for tinsseries tll(' eletnostatic o'llpoliclnt, E':, is (lillging wlil(h the covalentconipoilelit, C, re'llainls roughly constant, all observation ill agreeneltwith similar porph%'rin redox datat (27). An interesting solvent effect isobserved whi cobalt(II) P is oxidized to cobalt(ifl) Pc. i'cause thelatter species has a very strong propensity to lie hexacoordinated withtwo axially bound solvent molecules, the oxidation potential is clearlysolvent dependent. In a noncoordinating solvent such as dichloroben-

zene, a hexacoordinated cobalt(IlI) species cannot be formed. Underthese conditions, the Co(IlI) Pc/Co(li) Pc couple shifts positively to aconsiderable extent, such that the initial oxidation of the species is toform a cobalt(II) phthalocyanine cation radical (16). When pyridine isadded to such a solution, a cobalt(lIl) species apl~ently is formed.

Preliminary data for chro mium(IllI) phthaloc'yanines reveal in-creased stabili7ation of chrominio(l 1) %% ith strong donor solvents. Thiseflect could be d(Ic to stabilization of the low spin (4 chromium(II)through back donation to the plithalocyanine ring. However, furtherdata are necessary to inldrstand this phenomenon, especially as asimilar stabilization of h w-spi id ' d mai nmgnse(Hi) phthalocyanine ap-parently is 1)1o evident (Table V). SootC data eXist con1,r ning the effectof ring substihtim) on redox energies (,see Table 11). however, signific-antly less so thin) in the porphyrill series (28). G(enerally, electrondonors favor ring oxidation and disfavor ring reduction. An interestingcomparison exists for sulfonic acid substitution where the neutral acidform ofTsPcFe( 11) rednces at -0.40 V [vs. normal hydrogen electrode(NHE)I while the sodium salt, with four negative charges on theperiphery of the molecule, does not reduce until -0.67 V (5). Thesecond redtiction is similarly, but a little less markedly, effected. Forstibstihnents Sn01 is chloride, methyl, tcrtbntvl, sulfonic acid, car-Iboxylic acid, and olhvrs, the shiis ill redox energies (except for specialtlt.s sclh its jiist indict'd'(l whire the charnge oi the ring is modified)rarely (-.\.e( lot) 1 mY,V

Electronic Spe'ctr,ocopy,

fillllall 0t al. 629) first detailed the electronic structure o)fhlletailhm)iltll~itlcya.tli lu's, showing that t.: two principal bands (both'E. -'A,,) (for S = ) inetal ions) il the visible spectrun of allpllhalocyailtie (-2) s)cies cold )e assigned its as. (r) --* e, (7r*) (Qhand near 600t) inn) and a 2. (n") - e. (7r*) (Soret band near 350 nn).Unlike ti(, i)orphv rim, syste(m (28), tl, a,, and a2v orhitals are fairlywell-separate(l in cicrgy, ald these two transitions o!o not, therefore,mix appreciably. Emission data reveal that the spin triplet componentof tie t() bad lies alboult 51X)0-55M)(X ci below Ih' spin singlet (29).lI, systemns lacking chargre transfer absorption, such as most ofthe maingrotil) 'I)((' i., it is this spil triplel st;ih, thlt is likelh 1 )e photoactivewhenll t Illhthaloc.1 , is lfilizi.l a', . phmoo(.talyst. The '() state(fhiors(.('cll(-,) h .1. lhtill. o ooly a few lalinneOWCOllnls, while thetrilie( stll(- hiftllo, is Ill Ih lll'll , mIq u ll-ollillilis('old l gioll al liquid)litlliL''t 11 lll'll ll (29. 30t)} \ lt,'l/(. 11d.- llfilglli(i illlS such ;as cop)-

per(hI) are onllerled, Ills lowest state Illost lke'y is a triplet-.Iltiplet, lying at r.oglili te saint, energy its '() (28, 31).

lowcci, the sitoattin can change dranatically when chargetfalsiel" transitionis are present. Such transitions occur whenever metald-levcls lit- it energies inside the llOMO-LUM() bandgap of theplhthalocyamnie (or close to, but above tilt, LUM() energy). Such transi-tions %.,cre discussed in depth (31) and only a stmninary of these data isprovided here.

With nmoderatel. oxidizing ions such as manganese(Il1) andclironiu(till( II), charge traosfi," ab sorptioli from both the phthalo-cyan ic (11. aIid a2. olbitals uitlo eh'd) orbitals on the metal is allowedelectronically and observed readily, the former transition lying in thenear IR region (Figure 1, Table VI).

Consider the charge transfer transition:

Pc(-2)(a1 .YCr(II)d - Pc(- I)(a,.)Cr(II)d'

labeled LNICTI ill Figure 1. This reaction may be construed as thesum of two redox processes, viz:

I'c(-2)(aI,1 'Cr(lll)dt + e =Pc(- 2)(a )Cr(II)d (4)

Pc(-2)(a)J 2 Cr(II)d 4 = Pc(- l)(au))Cr(II)d 4 + e- (5)whose potentials [-0.40 and -(+0.70)\ j can be summed to yield a

::+usition energy of 1.10 eV, that is, a predicted charge transfer energyot 8870 ern', in satisfactory agreement with an observed transition at790) ci -'. Both chroiniunin(III) and manganese(IlI) exhibit chargetransfer hands in timte near III regioni, although for niaigaiese(lI) andchromium(ll) species these LMCT bands are llne-shilted to approxi-inately 11,000 in-'. A treatment similar to the one indicated in Equa-tions 3-5 allows prediction of the energies of these charge transfertransitions generally to within an accuracy of about 1000 cm -'/T .11111. VI)

Furtlher pr tloftlhc assignmnent is oblaii td Iby location of a second

cliaige traisle',r Ikamd (I ,C'I'2). ar-kinr frnin a2,, to e,(d) (Figure 1),kyillg betmwe Ille () amd Sol.l lhands. Nlost siginificalilly, the ellergyseparation hetwcel, IImes, two CllIrgC transfcr bands is almost exactlyetqual to the eniergy sCparation between the(- Q and Soret bands (31).Virtually all t' aiticipated LMCT hands ill time first-row traisifioulmetal plhthalocyanintes can be assigned and calculated by this simpleprocedure. MLCT band energies also can ie calculated, but appear tobe too weak to be ol)served. The energies of orbitally forbidden LMCTtransitions also can bt derived by this technique, allowing the pre-summed dctection of states that ctrunot be observed directly by elec-troic spectroscopy (Svc Tale VI).

Surprisingly, such t simple relationship between electronic ab-sorption bands and redoxi potctial energies is successfil. Evidently,the enitropy dillerencs bt-tw, cit tiet( varils components of the coupleare (.-IV small. Moreover. tlie I (l ('ransitions appear to be (0-0) invibrational character, elimniiatinig another possible source of disagree-Iment between calculated redox energy ,und observed data (31). Prece-dent for such agrcement between charge transfer energies and sums ofrcdox potentials exists (.32, 33). Because these LMCT transitions fre-

quently lie at energies beh)w the Q band, they (or higher spin versions)are likely to be photochemically active. However, lifetime data are notyet available.

Hence, a comibi iation of electronic absorption spectroscopy andclectroel'tniical illeaila"rement call Imlap tile energyv levels of a metal-lophthalocyaniite with considcrabl' atecracy, and provide a measureof the redox pottntials of the various cxcited states, information ofconsiderable value ill uiderstanding pliool-ctnical behavior. Al-though Ifw heavy transitioi metal ion pihthalocyanines have been in-vt.,lig vtl to dale, Ilicil belilavio is iiol .xpeechd to valy greatly frolllite dtils i.tieill m his chapter (provided tile Iletals lie inside theplithaloc alin e macrocyeh ring). In general, their d-levels will beIuricd below tIt pihtlalocyamine HOMO level and redox procems at

tIme metlaft -

-7-

Pholochvnji3I ry

A rangc of' igaiii groiij). first ro'% , atid( litter transition inetal ionswere scrCeIlcel recenitly for tin.ii ability to generate reduce(! iethylv'iologen (MV' ) when irradiated (into the Q hand) in the presence ofmmetliyl viologemi (MIV* ) and a donor suich ats trietliamolarnine (34).Species containing Mg. TFiO, Cr(I1), Fc(ll), Z1I), Ihl1.and Ru(II)generated reduced inethyl viologen, albeit in smrall yieldl (<.01%).Other mietal ions, specifically VO, Cr-(IIL), Nln(llI), Fe(lll), Go(Hl),Ni(1I), and Cui(1I) (lid not generate redulced mnethyl viologen undersimiflar conditions.

lolls withl low hyi ig (ncar. 114) LIMC], bands clem ly were photo-(Aclluill\ illactk\ bclflIsc liiiicil of tli excitation enlergy wits lost byvin tersysteil mmCrossing f ront tilme Q) hand to the low lying LMICT band.

Several iechamisimos are possib~le by which redumced imethyl violo-gell (MIV) mlight be produlcd. Speci ficillvN, reductive qulencing of thecexcited( state of' thle Iphotocatallyst (C*) bv time donor, viel(Ii og c oul~dreStilt ill form11ation (of, MV, by reaction of, MV2, with c il it followingthecrmail reamctionm. Alternattively, c' coliid be qimemielmed oxidatively byM\'2' yielding MV* directly, together withl C, , Vhichi couldl then returnto the ground State e by at tiiernal reactioni with the donor.

Detailed kinetic studlies, not yet undIertaken, are necessary to de-du1ce tmmeqivocally wvhichi mechianismi is occuirring.

Given the grounid state redox data discusse'd in) the section oncltcctrocmeiiiistry, together with the electronic absorption data giveni inltimI section onl electrom lic Npvct roscopy. thev redox potentials of the ex-

cited states. c*, can Ibe derived (35-377) 'Fius,, ifE, (ill electron volts)is tile eqtU ilil'iatt( eXCited Stalte- en ergy of th e lowest, jihotochemiicallyactive, excited state of the photocatalvst, then thle redox potentialsinvolimlg u* arc:

c'/ek- § c (6)

:/. c-/c -E ,., (7)

These equations should be fairly accurate providedthat the entropy differences between ground and excitedstates are small, as apparently is the case.

With these data, it is clearly possible to calculatethe thermodynamic driving forces for the various excitedstate and ground state reactions, discussed above. Whenthis is carried out f341 the complexes are readily dividedinto two sets, one in which the thermodynamics of one ormore processes make most unfavourable the production ofreduced methylviologen and one in which the thermodynamicsare less unfavourable or even slightly favourable. Exper-imentally the inactive metallophthalocyanines clearly belongto the former set, and the active species to the latterset. In no case, however, studied so far, are the thermo-dynamics for the formation of reduced methylviologenstrongly favourable explaining, in part, why the yieldshave not been high for these species.

Kinetic phenomena, that is, suppression of otherwisethermodynamically favourable back reactions, play adominant role in determining which catalysts are suit-able and which are not. It is equally clear from thisinvestigation [34] that studies such as those shown inthe "electrochemistry" and "electronic spectroscopy"can provide a sound basis for understanding photo-catalytic behaviour and can influence the design offuture catalysts. Growing interest in the use ofmetallophthalocyanines in solar energy conversionattests to the potential value of such catalysts (2,15,34,38-43). Such data are also of significant value inunderstanding biological photoredox behaviour, especiallyevents occurring during photosynthesis.

ACKNOWLEDGEMENTS

This research is part of a joint project withA.J. Bard (University of Texas at Austin) supportedby the Office of Naval Research (Washington). We arealso indebted to the Natural Sciences and EngineeringResearch Council (Ottawa) for financial support.

-9-

Literature Cited

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tsnkhsissi~li, 1979. 15. I -5h

5. Lever, A. It. 1'. Ads . 1I, ;, ( hr . jlibo ,pit. 1965, 7, 27.6. Lexa. 1).; leix 1. Clim. I'hp,. 1974, 71. 510, 517.7. Kadish, K. M.; Bolhomley, L. A.; Che'ng, J. S.J Am. Client. Sot. 1978, 100,

2731.8. Lever. A. B. P.; Minor, 1'. C. Adt. Mol. Relax. liter. Proe. 1980, 18, 115.9. Shepard, Jr. V. It.; Armstrong, N. I. J. Plis. ('hn. 1979, 83, 1268.

10. Wclberg, A.; Manassen, J. J. Am. Chem. Soc. 1970, 92, 2982.11. 0ll1esal, IL. D.; IaLnelss, It. T. J. Act. C/seen. Sie. 1968, 90, 1455.12. Clack, 1). W.; I lsh, N. S.; \\ools , I S. Isis. C/him. Acta 1976,19, 129.13. l)olphij I).; Jamses, B. . Mmray. A. j.; Thsrnbakck, J. It. Can. J. Chesee,

1980,3,S. 112514. lyver, A. It. ', Mie,. I' ( . Itg. C/hrst. 1981, 20, ()000.15. Loutiv. It. 0.; Chetg, 1'. C.., Pl~.s. C/s tit. 1980, 74, 2902.

'a 16. Giraudeau, A.. Fas, 1F-1. F.; BaId A. J. I. Am. ('ient.Soc. 1980, 102. 5137.17. Fanning. C Park, C. B., Jamss',, (, .; Ih Icatley. Jr., W. IJ. lNtor,. Nuc,

Chem. 1980, 12. 34:318. Beck, F. 1h'r. /l-sscu'ts. i'lts (:/et .1973, 7. 3519. Myers, J. F.; Ilaylsel-(taltlani, (. \V. I.'\ r, A. It. 1). /isng. Clhees. 1975, 1-1.

461.20. McIltigh, A. J.; Gotiterman, M., %V'iss, Jr. C. Th,'or. Chim. Arta 1972,24,

.3A6.

21 Edwvards, A. M.; (;ollhl Me . J, .I Mot Sin ertostc. 1970, 33, 292.22. Schall!r, A. %I.; Uoutereman, M. T'/tor. Chim Acta 1972, 25. 62; 1973,30,

9.23. Goutermaii, M. hit '".c PFeorplvriots"; Deolplius. I).; Ed. 1977; Vol. III.24. Shannon, It. 1.; Prewitt. C. T. Ata Crystallogr., Sect. 13, 1969, 25, 925.25. Gutnann, V. "'Tie )onor Acceptor Approach to Molecular Interactions";

Plenum: N(' York; 1978.26. Drago. It. S Strout n d Btd. 197:3, 15, 1:327. Katdli. K NI . BHtlhelt \. I, \ (t;. (1, t 1940, 1 . 832.28. )a\ is, 1) C Its .'.I'l P hoiis tet', Diljhite. I) : Ed. 1978; Vol. V. ). 127.;

I'tel t. R, II ibid p) 5:3.2). Vite.ll, V S . s sigi, E. \ NI I cckhll'. I. F I seess. I'/sc,. 1971,55. 4131.30 Yoshiesl K.; Ktests, K., 1etIeh. PlI. I/lips See. Jpis. 1973,3.5, 121.31. L'v sr, A. B. 1.; ictia, S. MIisseer, I'. C., l3stasswesaiy, B. S.; Pickens, S.

ItJ- Ate / s'l t (. 1 1981, It)t, (4 (N . (, 7 .' ., -!t c xAt , t1

-32:-I aintos,, I. S, J sss , I). C; 1,. / C/im. '/itqs. 1951, 48, C18.33. Barne.'. C.; I), , 1 .I .('lts. Sit. 1964, 3.58f.3-. Lever, A. 1. P.; I , (t i. S.: I ,ttsi.wan y, It S ; Kat lil, A.; Stylles, 1). V.

Istor. (him. A( (i 198 I1. I16.35. Meyer, T. J 1sr. J. Cci,'t 1976, 1.5, 2(M1.36. Whitte, I). ,. Ac',t. tC/esm. Reis. 1980. 13. 8337. Weller, A., ic, e. Ptst Apjl& C/ht. 1968, It6. 115.38. Jaeger, C. I) ; Fal, I"-R.; Bard, A. J.J. Am. Cients. Soc. 1980, 102, 2592.39. Lotitff It ().; Sharpl J. I. J. Ajtpl. lhetteehs't. 1977, 7, 315.40. Fan, "- ; IatUlker, 1,. IHJ. Am. ChJem. Sme. 1979. 101, 4779.41. Iarvet J.H (/', Che ts. Cossee. 1980, 805.42. larriman, A., ishos. M. Cj.. Phtol'hes. 1980, /., 253.43. Talmo, T.; Voirle. I)., Kaneko, M.; Yamada, A Be,. Bulesencsv. Phys.

(let. 1980,84, 10:32.44. Louts6s , H. 0., peersotal e'omeet sicattistee. 1981..15. Li, C.; Chin, D. Anal. Lett. 1975,8, 291.

'Th' eIshhaloc'yae iitle' ioelll t it't I izs d e l t' i's plesntelth its I11 19.

Ati energy lt\cl of s'eUett'lly (ec) localize is l the Iht' I)Itcral nitrogei atoms isorilitttid he'seests dc is Il,) evids, e'that it iplas a role is the slpe'troscopy.

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