Metal Complexes as Color Indicators for Solvet Parameters

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Metal Complexes as Color lndicatorsfor Solvent ParametersRudolf W. Souk up and Roland SchmidInstitute of Inorganic Chemistry, Technical University of Vienna. Getreidemarkt 9, A-1060 Vienna, Austria

"Illo res passim opporere eolores facit"(Turba philosophorum, SermoXI11 [ArtisAuriferae I],about 9th century AD.)Color indicators are numbered amone the oldest. simde st,

and most impressive tools of chemistry. ~ c t u a l l yheir usereaches hack to ancient times, when, in Plinius' days, oak gallsserved as a check for the purity of alum and for the adultera-tion of verdieris (1.2). Nowadavs the kev word "color indi-cator" makesone think of the large color box of pH indicatorsused in titratine Brdnsted acids with Brdnsted bases inaqueous media. Less known, however, is the fact t ha t thereare also nonprotic Lewis acid-base reactions accompanied bycolor changes, as is the case when particular metal &mplexesare dissolved in various solvents. For instance, in 1939, Bos,Lifschitz, and Dijkema recognized such color changes uponthe transfer of certain nickel complexes through solvents (3).Similar effects are displayed by some cobalt complexes (4).The strategy, however, of employing metal complexes as colorindicators for solvent properties (5) does not appear to bemuch older than a decade. Nevertheless, metal ions are ex-cellent color indicators for both the a-donor (nucleophilic,basic, cation-solvating) and u-acceptor (electrophilic, acidic,anion-solvating) abilities of solvents, and this phenomenonis reviewed here.Basicity Indicators

In 1982. Sone and Fukuda. in a vew interestine review (6).have pointed to the solvent effects on certain copper(I1)complexes dissolved in nonaaueous media. They have shownthat in the case of ~u(t me n) (a ca c) ~1 0~I ) , where tmen =N,N,N',N1-tetramethylethylenediaminend acac = acetyla-cetonate, the color of the solutions follows the donor abilityof the solvent, as is quantita tively revealed (7)by t he linearrelationship between the wave number of the maximum ab-sorption band in the visible range and the solvent donornumber (DN) (8) see appendix). T he color changes are sodramatic that even with the naked eye a rough order of donorsolvents is afforded, which is useful, for instance, for a teachingdemonstration. Thus, Cu(tmen)(acac)+ is reddish violet inweakly basic nitrobenzene (DN= 4), changes to a bluish violetin the medium basic ketones, e.g., acetone (DN = 17), s azur

Solv. -

in formamide (DN= 27). turquoise in strongly basic pyridine(D N = 331, and finally becomes green in extremely basic pi-peridine (DN = 51). In t his way a color chart is obtained en-abling the ordering of other donor solvents. One merit of thisindicator is tha t i t straightforwardly demonstrates that allcompounds involving free electron pairs are good donors,whereas solubility limits are reachedin highly inert solventssuch as hexane and carbon tetrachloride, or eenerallv in vervpoor acceptor solvents of acceptor number;, AN 2 . ~ h kother limit is reached in extreme acceptors like formic acid(AN = 84), where the indicator is no longer stable.

The effect of donor solvents on the electronic structure ofthe metul complex under runsiderution is well undrrstood inthe light of simple molecular orbital iheury (Fig. 1). Th echelate cation is axially solrated a) hat the original squareplanar complex is transformed tu a te~rugonal,and ventuallyLOw approximatels ocmhedrnl confirmration accordine to thestrength of solvatibn. Let us sta rt from the octahedron andconsider the change in the d orbital splitting when the ligandson the z axis are gradually moved out and those in the X Yplane moved in. As a result, the interactions become less withorbitals have a z component, i.e., d,z, d,,, and d,,. These or-bitals are thus stabilized (whereupon the others are enereet-ically raised). The color of th e sq"are planar complex, &en,is due to transitions to a "hole" in d,2-,2 from the other fourorbitals. Of these, the highest energy transition d,,, d,, -dX2-,2 at about 550 nm (9) can be assumed to be the mostimportant m e, followed hy dl ?- ,>-,z. C'onsequ~ntly, nincrease inz ligand haliic~ty hould producea blueshii t as thearrows in the figure, denoting energy differences, becomegradually shorter (6). In fact, this feature is found experi-mentally.

CoverSolvent Baslcltles and Acidities Visualized by theColors of Cu(tmen)(acac)CIOa and Fe(phen)*(CN)~,Respectively

The donor numben (DN) and acceptor numbers (AN) have beenadopted as empirical qu antities to scale th e electron pair donating andelectron pair accepting properties of me solvents. Admittedly,mere areother more w less equivalent scales which may be u sedas well. Thusthe D N scale is in close agreement with, e.g., the Lewis basicity pa.rameter B. whereas in place of the acce pt a numbers theh values canbe uiiiized.Care,however, is advised in me case of lhe highly sbucturedsolvents since their coord ination chemical power strongly dependsonthe properliesof me ubstrate. For an overview see references (35)and( 3 0 .~~~,Slope and position of the iines connecting DN an d A N values in thediagram are pertinent to the characterization of solvents: Predominan tdonon (acceptors)have a steep negative (positive) slope. Elevated fiatiines w i d dentify highly amphateric solvents like water and fwmamidewhich areoutstanding far their excellent capabilily of dissolving salts.As has been shown recently ( 15 ) . siop eand position of the iines allowmyh stimates to be madeeven o me olvent dieiecb ic constants sincethe latter can be represented by a linear combination of acceptornumbers and donor numbers (35, 36).

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square Soh Solvplanar tetragonal octahedral

Figure 1.'Relative orbital energy levels for Cu2+ in Square planar, tebagonal.and octahedral environment( 6 ) .Theconfigurationalchanges includea length-ening of the original bonds accompanying the increase in the coordinationnumber. The subfigure is far Ni2+.

I t may be noted in passing tha t the third transition, dry-l%-p, is much less affected by changing degrees in solvation.Therefore. if this transition also contributes sienificantlv tothr spec trkn , as is the case fur unmixed c op pe r~ ll )crtylar-etonatr ( 1 0 ) .solvenl etfrcts leav to the eve less. but can I wmade obvious by mathematical'spectral analyses (11). Theadvanrage oiCu(acacJ?,huwever, is its surprisingly good sol-uhilny wen in highly inert cyclohexane, carbon tetrachloride,and in very weak acceptors like esters, ethers, etc.From a historical point of view, it may he mentioned th atthe engagement of copper(I1) as central metal proved to hea nice enlargement asregards the play of colors of the solu-tions. Apart from the rather insensitive vanadvl acetylace-tonate (12-14), whose color changes merely from turquoise(in benzene) to green (in piperidine), attention has been paidas well to Ni(tmen)(acac)+ (15). In this instance the colorchanges abruptly from red to green on varying solvent basicity,without any intermediate shades, because of the existence oftwo discrete complex forms, one octahedral and the otherquadratic planar (16),

Ni(tmen)(acac)++ 2 solvent+ Ni(tmen)(acac)(solvent)z'red green

The differina solvatochromism between CuZf and Ni2+ followsfrom the diifr rent elrctronic confiyuratims. As already cor-rrct ly noted by Lifschirz 131, ortahedral nick~! l(l l) s wara-magnetic high-spin, whereas the planar form is diamagneticlow-spin (inset, Fig. 1). Therefore a configurational changeis connected with a loss or gain of the discrete pairing en-ergy.Thus , the importance of the proper choice of both centralatom and chelate ligand to yield a suitable color basicity in-dicator is emphasized. It should be mentioned that simplecrystal field theory cannot account for the spectral differencesbetween Cu(tmen)(acac)+ and Cu(acacb because of the. ~ ~ ~~~owrsimplification of tht: point-charge mu dd An nd\,ancementis afforderl bv differrntmtinr lrrween o and a lieand; thrnurh

increase in solvent acceptor strengthFe(phen12(CN12, octahedral, distorted,

distorted ~elphen~F- like Fe(phen12(NCBHJ2-like(compressed) (elongated

Figure2. Simplifiedschematic orbital arrangements rationalizing the contrarysolvent effects n the spectraofFe(phenb(CNh nd Fe(~hen)dCN)~+.he di-agram. not drawn to scale, is adiusted so that ?r.is constant.Therefore the~~~-center ofgravityofthe dlevels varieswith the ligard field strength. The splittingof the dorbitals isaflerPurceliet a1.(29).Full amws:the relevant bdhansitionin iron(ll1). dotdash arrows: Fe(l1) charge transfer transition. Botlom: configu-rational changes following salvation.

bonding ligands, as acetylacetonate is, the customary d orbitalsequence is upset with the orbitals having r character, viz.,tze, being raised. Upon this, the unpaired electron comes tohe located in d,, (17). In the above mono-acetylacetonatecomplex this effect is not yet dominant .Acidity Indicators

Nearly a dozen metal complexes are known whose spectraare reipnnsivt: to solvent-acceptor properlien. A prominentfeature is the inclusion of ligancls providing a lone electron paircawahle of electro~hilicolvation. viz.. cvanide. carhon i o n -, "oxide, and in one case thiocyanate; specific examples areCI(NCS)~~- , o(bipy)(CO)a, W(hipy)(CO)s Fe(bipy)z(CN)z,and Fe(hipy)(CN)?, where bipy = 2,2'-bipyridine (18,19) .In another instance variations in color can occur if not sostrongly hound ligands (absence of appreciable back-dona-tion) dissociate in stronelv anion-solvatine solvents. Such asituation is met with wceeh the perchloraG ion in the abovementioned Cu(tmen)(acac)ClOa is substituted with halo-. . .genides (6).Among all these the complex F e( ph en )~ (C N) ~111, wherephen = 1,lO-phenanthroline, is outstanding and, moreover,can be readily purchased. T he wave number of the maximum

molecular orbital consider%ions. In the caseuof strong ?r460 Journal of Chemical Education


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Figwe 3. Solvent dependent spctra of Fe(phenMCN)r (A)andFe(phen)ACN)z+(6). (A) [Fe(phen)p(CN)2]= 1 X lo-' Meach in dimethylformamide I ), nitr-methane (21, dilute hydrochloric acid (3). and concentrated sulfuric acid (4) .Broken line: [Fe(phet~)~~+]5 X lo@ Mi n acetonitrile. Path length= 1 cm.(0 ) Fe(phen),(CN),+] = 1 X Meach in nitromethane (l' ), ormic acid (2').70% perchloric acid (3'). and concentrated sulfuric acid (4'). Broken line: [Fe-(phenh3+] = 5 X Min acetonitrile. Path length = 1 cm.

of the charge transfer band (tz. - *) 20) correlates excel-.lently with'iht. solvent acceptor number ( 21 .22 ) (see appen-dix).The com~lexs blue in solvents with feeble acidic prop-erties (e.g., hexamethylphosphoric triamide, AN = l l ) ,changes to red in more acidic ethanol (AN = 37) and finallyturns yellow in formic acid (AN = 84), without any decom-position. Between, all conceivable color shades are encoun-tered, reddish violet, bluish violet, orange, etc.Remarkably, the solvatochromism met with the corre-sponding trivalent compound, Fe(phen )~(CN)z+ ,s just th econverse of tha t for Fe( ~he n)z (CN )z18,20): The color is redin nmr accentors (e.e.. nitromethane, Ah = 20), reddish violetin medium acceptors like water (AN = 55), and blue to greenin concentrated mineral acids (AN > 100). This unexpectedfeature becomes intelligible when one remembers th at (1) thecolor of spin-paired iron(III), as in Fe(phen)z(CN)z+, s notdue to charge transfer bu t rather t o d-d transitions, namelyt 5- 4e (23.24). (2) the cvanide ions are cis-bonded in bothcomplexes (25-27) in contras t to former report s (28), (3) th eorder of d levels for divalent Fe(phen)z(CN)z is d,z-,z >d,,,d,, > d,, (29), he same as that of th e trivalent complex(25) ,and (4) solvation effects a lowering of the otherwise ex-treme ligand field produced by th e cyanide ligand which lieson the high end of th e spectrochemical series.In these terms, the oppositely directed spectral changes arereadily visualized (Fig. 2). Consider the shortening andlengthening of the arrows pertinent to th e respective trans i-

tions: T he d-d transitions are energetically lowered (in par-ticular due to a lowered dr2-,,* level) by a decrease in ligandfield strength shown by the energy differences between t2g ande, orbitals (i.e., 10Dq ) ,whereas the charge transfer transitionis raised (29) because of the descending barycenter of the dorbitals.Alterations in the electron distribution within a moleculewould always be accompanied by changes in the geometry. Inour case. the comnressed trieonal distortion. modified bv asuperimposed rh imbi c splitting (25, 30) in th e unsolvaiedoctahedral com~lexe sboth divalent and trivalent iron) tendsLO iw removed hy acceptor-solvent attack. Thereby the n ba-sicitv of the cvanide liaand is reduced. resulrinc in Fe-C bondlenGhenings.-(However, the Fe-C distance increase expectedfrom the cvanide o basicitv decrease might be dampened.since, in addition, cyanide is a a acid and s&ation rei ders itmore acidic.) Very strong acceptors, finally, renew the dis-tortion as is depicted in Figure 2. In between, when the ligandfield strength has become just that of phen, a regular octa-hedron might be expected in which case the Fe( ph en )~ (CN) ~+spectrum should mimic that of the tris complex Fe(phed33+,and, analogously, the spectrum of Fe(phen)z(CN)z that ofFe(p hen)~Z+. his is, in fact, borne out and is observed fordivalent and trivalent iron in solvents of acceptor numbersof about 70 and 100, respectively (Fig. 3). (Th e molar absor-bancies. interestinelv. are about half that of the tris hena an-.,throline complexes.)Finallv. the extraordinarv stabilitv of F e ( ~ h e n ) d C N )toward Ggand exchange is worth mentioning. i ts so1;bilitihowever, is limited in solvents of feeble coordination prop-erties. If need be, one can switch over to th e tetramethylsubstituted comuound which is somewhat more soluble. Wewould like to conclude by noting tha t the Fe(phen)z(CN)z+cation is reduced bv strong donor solvents raisinn difficultiesin distinguishing bitween redon and solvatwhro~&sm ffecls.[Trends between donor and reducing ca~a bil iti es fsolvenware discussed in a recent paper (32).i

ReagentsThe metal complexes can be easily prepared accord-ing to literature methods: Cu(tmen)(acac)CIOa (or, pre-ferably, the nitrate) (71, Fe(tmph en)z (CN) ~2Hz O 33),Fe(phen)z(CN)zN034Hz0 33) and Ni(tmen)(acac)B(CsH&(16). A convenient source of the latter is Ni(OHz)6(N03)2(16) ,instead of the commonlv used ~e rchl or at ealt du rine whoseuse a somewhat dangerous side-product develops. kinally,Fe(phen)z(CN)~.2H~Os available commerciallv (Alfa Ven-tron, 12189), but can he prepared similarly tb the tetra-methyl-substituted analog.The solvents to be used in the following experiments shouldbe of the highest purity available.

Experiment 1: Scaling the Electro n Pair Donating Ability ofSolvents,\ spatula-ttpful ,of Cwrmen)(acarJSOx i s diriulvtd in the vnrroussdvents dcscrtlwd in thr trxr. Srr ruwr and bu x drwrilring it in thisarticle.

Experiment 2: Scaling the Electron Pair Acceptor Ability ofSolventsFe(phen)n(CN)z s dissolved in the following solvents (for alter-natives see box describing cover), gradually with warming:hexamethylphosphoric riamide (AN= 11,blue solution)N,N-dimethylformamide AN = 16,bluish violet)diehloromethane (AN = 20 , violet)ethanol (AN = 37, ruhy-colored)glacial acetic acid (AN= 53, bright red)formic acid (AN = 84, yellow)trifluoroacetie acid (AN = 105,pale yellow).

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Th us by simply dissolving th e indicators and putting the emergingcolor shades in their places, bath DN qand AN values can be evalu-ated for a large variety of solvents including water, methano l, ethanol,butanal, acetaldehyde, acetone,' dioaane, tetrahy drofu ran,propylenecarbonate, acetic anhydrid, cblorobenzene,' bromobenzene,* pyri-dine, aniline, benzylmethylamine,* formam ide, dimethylform amide,nitromethane,nitrobenzene, acetonitrile, sulfalan e, rim ethyl phos-phate, and hexamethylphosphoric riam ide. (The asterisk denotesthat Fe(tmphen)n(CN)~s used for solubility reaso ns.)Experiment 3: Scaling the Acidityof Extremely Strong Acids

By means of Fe(phen)z(CN)nNOswe get the ord er of concentratedmineral acids

HC I < H N 0 3

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Metal Complexes as Color Indicators for Solvet Parameters