Chemistry and Artists' Colors
Mary Virginia Oma, O.S.U. College of New Rochelle New Rochelle, NY 10801
Part 11. Structural features of colored compounds
In Part I of this article the nature of color has been discussed with respect to the three componentn of color production and the molecular basis of light modification by colored compounds. It now remains for us to examine the kindsof chemical substances which exhibit color. We know that many different materials, both natural and synthetic, exhibit color. Among the colored materials taken from living things are chlorophyll from green plants and hemoglobin from blood. The world of minerals and gems is full of ckor. Many of the materials that are used as a&sts' colors consist of these inorganic materials. A cursory glance through the table of physical constants of inorganic compounds in the "Handbook of Chemistry and Physics" (18) indicates that most of the compounds ofelemen& of thed-block (transition elements) exhibit color, and since the elemenu are arranged in the Periodic Table on the basis of atomic structure. this fact suggests that certain structural features might he responsible for the observed colors. However. other elements which fall outside the transition group suchas lead, antimony, arsenic, and sulfur exhihit a wide variety of cohrs in their compounds. Thus, we might conclude that~morethan a single strktural characteristic is caoahle of givine rise to colored com~ounds. In any event, one fact is clear: sibstances exhibit color, because they are capable of absorbing selected wauelength ranges of white light. The absorbed light promotes the electrons within these species to higher energy levels. Such transitions lead to changes in the colors observed. Changes in energy level differences are fundamentally environmental chanees: the atoms in the snecies are sensitive to the forces surr&nding them so that virtually any change in the atoms' electronic environments will ~ r o d u c ecolor chanees. This is " a working principle which we can use to examine the origin of color in numerous substances. Althnugh the structural differences among artistd colors have been discussed in greater detail hv this author elsewhere (10, 19), a brief review-of the various-types of colored compounds follows. Conjugated2 Organlc Compounds We saw in the discussion in Part I that a molecule must selectivelv absorb electromaenetic radiation in the visible region of the spectrum in order to exhibit color. In the case of colorless organic molecules. the electrons which constitute the chemical hinds occupy bonding molecular orbitals with relatively low energies with respect to the higher energy antibonding orbitals. As a result, it takes a great deal of energy to promote such a species to an excited state, and the molecule will absorb energy only in the ultraviolet. This is true of saturated compounds and of unsaturated compounds possessing only a few double bonds. However, as the degree of conjugation increases, the energy gap between electronic energy levels tends to decrease. and thus less enerw is reauired to achieve an excited state. If the degree of conjugation is great enough, selective absorption will take place in the visible region. In addition, if nitrogen or other heteroatoms are present in the conjugated system, the required degree of conjugation is greatly diminished because of the presence of lone-pair electrons on the heteroatoms.
-.
2A
conjugate organic compound is one which contains two or more
double bonds with a single bond positioned between them. 264 1 Journal of Chemical Education
Several theoretical models can account for this observed trend in organic compounds, namely, molecular orbital theory, valence bond theory, and the free electron theory. Although the application of the latter model is limited to ?r-electron systems, most dyes and organic pigments fall into this category. This model assumes that the conjugated r-electrons are in a well of constant notential e n e r w ~ i h o s eboundaries are roughly a little longer ihan the lengtrof the carbon chain, and that the electrons are free to move within the system. These assumptions reduce the problem to solving the SchrBdinger equation for a particle in a one-dimensionalbox. This amounts to the calculation of the theoretical absorption energy by E = (n 1) h/8mL2 where n is the number of conjugated ?r-electrons, h is the Planck constant, rn is the mass of the electron, and L is the length of the potential energy well. Although the required transition energy increases as the number of ?r-electrons increases, this is more than offset hy the fact that the enerw decreases as the sauare of the enerev well leneth (which i&oughly equal to the'length of the mz&de). T ~ S as the deeree of coniueation . .. increases. the leneth of the molecule increases, and the required transition energy decreases. This model has vielded calculated values verv close to the ohserved values for several series of dyes (20,21). The properties of several organic artists' pigments are listed in Table 3. All of them exhibit color, because of extended conjugation and the presence of heteroatoms. The vhthalocyanine pigments also exhibit color, because of the presence of copper(I1); this effect will he discussed in the following
+
Coordination Compounds Toward the end of the nineteenth century, compounds that defied the usual rules of valence occupied the thoughts of many inorganic chemists. In 1893, a Swiss chemist by the name of Alfred Werner published his famous paper, "Contribution to the Constitution of Inorganic Compounds" (22). In this, he set forth his ideas of coordination sphere and molecular geometry as they applied to inorganic compounds. He postulated that coordination compounds consisted of a central metal ion around which were arranged either neutral molecules or anions (the lieands of modern terminolow) and that each metal ion had acharacteristic number of &dination sites, which he termed the coordination number. He also proposed that the bonds to the ligands were fixed in space and could therefore he treated hv the aoolication of structural principles. Although ~ e r n ~ r the&y 's accounted for the structures of such compounds as lCo(NH~)sH~01Cls. it had .. nothing a t all to say about the nature of the hondingwithin the coordination sphere. It was Linus Pauling who made the first successful application of what is now called valence bond theory to coordination compounds. This theory, however, had one verv serious shortcomine. As we observed previoush, the d-block elements of the fourth oeriod exhibit color in most of their comoounds. Nickel salts are often green, copper salts are hlue to blue-green, and chromium com~ounds(as the element's verv name indicates) may be red, green. yellow, or hlue. These rolurs result from absorption of visible light caused bv electronic transitions from the ground state to excited states. The inability of valence bond theory to account for these transitions and the resulting colors was probably a key factor in the ascendancy
,
Table 3. Representative Organic Artists' Pigments Piament
Shucwre
Hansa Yellow i0G (ca. 1910-11)
a&
Vhlhalocyanine "Thalo" Blue (1937)
Colour Index No.
FTooertiss
Pigment Yellow 3 (11710)
Insoluble in water; stable to dilute acid 8 base; soluble in most organic solvents; insoluble in linseed oil: Stable to 14OoC.;good permanency to light
Pigment Blue 15 (74160)
Insoluble in water, organic solvents. 011s;Stable to dilute acid and base; stable lo 200°C.: excellent permanency to light
Pigment Violet 19 (46500)
Insoluble in organic salvems and water; stable in acidic 8 basic SOIUtiOnS; Stable to 165% excellent permanency to light
"-"\,FaN
N
'& L C/ I
I
QuinacridoneViolet (1956)
of other views on bonding. In the late twenties and early thirties, H. Bethe and J. H. VanVleck developed the crystal field theory of coordination compound honding. This theory considers the effect of the lieand on the d-orbitals of the central metal ion. The end res& of the imposition of the ligand field is a splitting of the ground state of the metal ion into several different levels, the number of which depends upon the nature of the field and the number of d-electrons. The existence of d-orbitals with several different energies allows for excitation of electrons from one level to another. These transitions are often referred to as crystal field, or d-d-transitions, and the energies required often corres~ondto the wavelengths of visible liaht. These transitions 'annot occur if the metal ion has no d~electrons, as in the case of Sc3+, or if the d-orbitals are completely filled, as in the case of Zn2+. Most of the compounds of the intermediate d-block elements of the fourth period exhibit colors due to d-d transitions, and the colors of most minerals and gems can be partially explained by this model (23). Several general articles on the crystal field theory have appeared in this Journal (24), so the discussion here has been deliberately brief. Numerous artists' pigments exhibit color due to crystal field transitions. Among them are manganese pyrophosphate (MnNHdP207); chromium oxide green (Cr203); Thhard's Blue (CoA1204); smalt (KzCoSi40lo); the many "earth pigments" based on Fez03and hydrous iron oxide, FeO(0H). The color of many minerals and gems such as emerald, ruby, olivine, turquoise, azurite, and malachite are the result of these transitions. Also, the copper(I1) in the phtbalocyanine pig~ i.m e nhv t. crystal field ments contributes to the color of the . . transitions. The presence of color in the metal-free pigment indicates a contribution due to the coniugated organic moiety as well. The properties of several of this; materials are listed in Table 4. Compounds Exhlbitlng - Color Due to Charge . Transfer Although the colon arising from crystal field transitions are sornetimr.i striking nnd brilliant, they nrc, in general, not as intense as the colors of pigments like chrome (PhCrOJ, or antimony vermillion (SbzSa). However, in neither of these cases does the metal undergo crystal field transitions, because it does not meet the incompletely filled d-orbitals requirement. However, the energy required to transfer an electron from one atom to another within these compounds has been found to correspond to the energies of visible light. Such
Figure 5. Charge transfer transition i n s molecular complex
compounds are said to exhibit color due to charge transfer transitions. Charge transfer can occur when two atoms in close proximity to one another have enerw levels similar to those illustrated in Figure 5. A level ichtmc for a typical donur-acreptor pair is shown on the left side of the d~agram(251,and absorption of electromagnetic radiation can>esult in promotion of an electron from the highest occupied orbital of the donor to the lowest unoccupied orbital of the acceptor species. Such a transition results in the formation of an ionic structure, (D+-A-), and if the energy required for the transition is of the order of the energies of visible light, then color will be seen. Although the diagram itlustrates a transition in a molecular complex, the same type of transition can take place in many inorganic compounds. Some examples of such transitions in artists' pigments are the following: Chrome Yellow: PbC$+042-Bhg_PbCr5+O-032Light
Prussian Blue: Fe3+(NHJFe2+(CN)s Light
Fe2+INHJFe3+ (CNh
In the first example, transfer of an electron from an oxwen ligand tochromiim partially reduces the rhromium to C;;VJ. In the second example, Fetll) partially reduces Fe(lll),while itself becoming oxidized to Ferlllt. The propertics of several of these compounds that are used as artists' pigments are also listed in Table 4. VoIume 57, Number 4, April 1980 1 265
Exclhd Slmh
In the formation of crystalline or semicrystalline materials, numerous atoms with only definite allowed energies are arraneed in such a wav as to multiolv . . the number of allowed energy states. These states are so closely spaced that a collertion of such states is often referred to as a "band," and the spacing hetween bands, that is the energy forbidden to the crystal's elertronp, is called the-hand gap." In metalliccrystala, the valence electrons are very loosely held, and electromagnetic radiation of almost all wavelenahs can be absorbed. Non-metallic crystals, however, contain very few loosely bound electrons, and a photon will be absorbed only if it has enough energy to overcome the hand gap and free an electron from a hnnding orbital. If the hand gap falls within the visible reeion. the crvstal u,ill absorb visihle liaht and exhihit color. M & ~semiconducting crystals such as^^^, AgI, CuBr and ZnSe have band gaps between 1.7 and 3.1 eV, and are colored because of band transitions. Two notable artists' pigments which undergo band transitions are vermillion (HgS) and cadmium yellow (CdS). It is important to note that because of the presence of many closely spaced energy levels in the conduction band, i.e., the level to which electrons are promoted from the valence band, semiconducting crystals will absorb all wavelengths of electromagnetic radiation equal to and greater than the band gap energy. For example, the hand gap of.CdS is about 2.4 eV, so CdS will absorb all electromagnetic radiation with an energy of 2.4 eV or greater. CdSe, with a band gap of 1.74 eV, will absorb all electromagnetic radiation with an energy of 1.74 eV, or ereater. This behavior results in unex~ectedcolors, but a comparison of the spectrophotometric behavior of a semiconductor with a c m n ~ o u n dwhich exhibits color by charge transfer will readily make the difference clear. In E'jgure6, KMnOa. with an enerm level difference of about 2.3 eV, undergoes a charge transfer transition by absorbing electro-
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