Chemistry » Transition Metals » Spectroscopic and Magnetic Properties of Coordination Compounds

Colors of Transition Metal Complexes

Colors of Transition Metal Complexes

When atoms or molecules absorb light at the proper frequency, their electrons are excited to higher-energy orbitals. For many main group atoms and molecules, the absorbed photons are in the ultraviolet range of the electromagnetic spectrum, which cannot be detected by the human eye. For coordination compounds, the energy difference between the d orbitals often allows photons in the visible range to be absorbed.

The human eye perceives a mixture of all the colors, in the proportions present in sunlight, as white light. Complementary colors, those located across from each other on a color wheel, are also used in color vision. The eye perceives a mixture of two complementary colors, in the proper proportions, as white light.

Likewise, when a color is missing from white light, the eye sees its complement. For example, when red photons are absorbed from white light, the eyes see the color green. When violet photons are removed from white light, the eyes see lemon yellow. The blue color of the [Cu(NH3)4]2+ ion results because this ion absorbs orange and red light, leaving the complementary colors of blue and green (see the figure below).

This figure includes three diagrams. In a, red, orange, yellow, green, blue, indigo, and violet incident arrows point from the upper left to the lower right toward one black, one white, and two yellow rectangular surfaces. No arrows are indicated leaving the black surface. On the white surface, arrows of all included colors extend from the surface of the rectangle extending from just right of the tips of the incident arrows that approach the rectangle to the upper right. For the first yellow surface, only a yellow arrow extends from a point just past the tip of the yellow incident arrow to the upper right. For the second yellow surface, all colors of arrows except indigo extend from points just past the tips of the incident arrows to the upper right. In b, a circle shaded in red at the upper left blends to orange, yellow, green, blue, indigo, and violet moving clockwise about the circle. The leftmost side of the red region has a radius that is labeled, “800 n m,” at the edge of the circle. A radius drawn to the top of the circle in the red-orange region is labeled, “620 n m.” A radius drawn to a point near the center of the first quadrant of the circle in the orange-yellow region is labeled, “580 n m.” A radius drawn to a point near the center of the second quadrant of the circle in the yellow-green region is labeled, “560 n m.” A radius drawn to the bottom of the circle in the blue region is labeled, “490 n m.” A radius drawn to a point near the center of the third quadrant of the circle in the indigo region is labeled, “430 n m.” An unlabeled radius is drawn to the leftmost point on the circle in the violet region. The violet region ends where the red region began on the circle. This radius is labeled, “400 n m,” just to the left and below the, “800 n m,” label associated with the red region. In c, a test tube containing a blue substance is shown. To the left of the test tube, incident colored arrows are shown pointing to the test tube. The colors of the arrows in order from bottom to top are red, orange, yellow, green, blue, indigo, and violet. Above these arrows is the label, “White light.” To the right of the test tube, yellow, green, blue, indigo, and violet arrows point right. These arrows are positioned at the same level on the test tube as their matching incident arrows on the left of the test tube. Above these arrows is the label, “Blue-appearing light.”

(a) An object is black if it absorbs all colors of light. If it reflects all colors of light, it is white. An object has a color if it absorbs all colors except one, such as this yellow strip. The strip also appears yellow if it absorbs the complementary color from white light (in this case, indigo). (b) Complementary colors are located directly across from one another on the color wheel. (c) A solution of [Cu(NH3)4]2+ ions absorbs red and orange light, so the transmitted light appears as the complementary color, blue.

Example

Colors of Complexes

The octahedral complex [Ti(H2O)6]3+ has a single d electron. To excite this electron from the ground state t2g orbital to the eg orbital, this complex absorbs light from 450 to 600 nm. The maximum absorbance corresponds to Δoct and occurs at 499 nm. Calculate the value of Δoct in Joules and predict what color the solution will appear.

Solution

Using Planck’s equation (refer to the section on electromagnetic energy), we calculate:

\(v\;=\;\cfrac{c}{\lambda }\phantom{\rule{0.4em}{0ex}}\text{so}\phantom{\rule{0.4em}{0ex}}\cfrac{3.00\;×\;{10}^{8}\text{m/s}}{\cfrac{\text{499 nm}\;×\;\text{1 m}}{{10}^{9}\text{nm}}}\;=\;6.01\;×\;{10}^{14}\;\text{Hz}\)

\(E=hnu\;\text{so}\;6.63\;×\;{10}^{-34}\;\text{J}\text{·}\text{s}\;×\;6.01\;×\;{10}^{14}\;\text{Hz}=3.99\;×\;{10}^{-19}\;\text{Joules/ion}\)

Because the complex absorbs 600 nm (orange) through 450 (blue), the indigo, violet, and red wavelengths will be transmitted, and the complex will appear purple.

Small changes in the relative energies of the orbitals that electrons are transitioning between can lead to drastic shifts in the color of light absorbed. Therefore, the colors of coordination compounds depend on many factors. As shown in the figure below, different aqueous metal ions can have different colors. In addition, different oxidation states of one metal can produce different colors, as shown for the vanadium complexes in the link below.

This figure shows three containers filled with liquids of different colors. The first appears to be purple, the second, orange, and the third red.

The partially filled d orbitals of the stable ions Cr3+(aq), Fe3+(aq), and Co2+(aq) (left, center and right, respectively) give rise to various colors. (credit: Sahar Atwa)

The specific ligands coordinated to the metal center also influence the color of coordination complexes. For example, the iron(II) complex [Fe(H2O)6]SO4 appears blue-green because the high-spin complex absorbs photons in the red wavelengths (see the figure below). In contrast, the low-spin iron(II) complex K4[Fe(CN)6] appears pale yellow because it absorbs higher-energy violet photons.

Two photos are shown. Photo a on the left shows a small mound of a white crystalline powder with a very faint yellow tint on a watch glass. Photo b shows a small mound of a yellow-tan crystalline powder.

Both (a) hexaaquairon(II) sulfate and (b) potassium hexacyanoferrate(II) contain d6 iron(II) octahedral metal centers, but they absorb photons in different ranges of the visible spectrum.

Optional Video:

Watch this video below of the reduction of vanadium complexes to observe the colorful effect of changing oxidation states.

In general, strong-field ligands cause a large split in the energies of d orbitals of the central metal atom (large Δoct). Transition metal coordination compounds with these ligands are yellow, orange, or red because they absorb higher-energy violet or blue light. On the other hand, coordination compounds of transition metals with weak-field ligands are often blue-green, blue, or indigo because they absorb lower-energy yellow, orange, or red light.

A coordination compound of the Cu+ ion has a d10 configuration, and all the eg orbitals are filled. To excite an electron to a higher level, such as the 4p orbital, photons of very high energy are necessary. This energy corresponds to very short wavelengths in the ultraviolet region of the spectrum. No visible light is absorbed, so the eye sees no change, and the compound appears white or colorless.

A solution containing [Cu(CN)2], for example, is colorless. On the other hand, octahedral Cu2+ complexes have a vacancy in the eg orbitals, and electrons can be excited to this level. The wavelength (energy) of the light absorbed corresponds to the visible part of the spectrum, and Cu2+ complexes are almost always colored—blue, blue-green violet, or yellow (see the figure below). Although CFT successfully describes many properties of coordination complexes, molecular orbital explanations (beyond the introductory scope provided here) are required to understand fully the behavior of coordination complexes.

Two photos are shown. Photo a on the left shows a small mound of a white crystalline powder on a watch glass. Photo b shows a small mound of a bright blue crystalline powder.

(a) Copper(I) complexes with d10 configurations such as CuI tend to be colorless, whereas (b) d9 copper(II) complexes such as Cu(NO3)2·5H2O are brightly colored.

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