20.10: Colors and Magnetism

Color in Coordination 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 and emitted, which is seen as colors by the human eye.

Figure 1. Electromagnetic spectrum of visible light and absorbance.

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 the coordination compounds depend on many factors, like:

Different aqueous metal ions can have different colors.
• Different oxidation states of one metal can produce different colors.
• Specific ligands coordinated to the metal center 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. In contrast, the low-spin iron(II) complex K₄[Fe(CN)₆] appears pale yellow because it absorbs higher-energy violet photons.

In general, strong-field ligands cause a large split in the energies of d orbitals of the central metal atom (large Δ). 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. The strength of the ligands to split the d orbitals are listed in the spectrochemical series. Here the ligands are written in the increasing value of crystal field splitting energy (Δ).

<img alt="Ligand field strength series diagram, showing increasing field strength from I- to CO." src="/files/ftp_upload/11463/11463_Image_2.png" style="text-align: center; height: 100px;" data-altupdatedat="1754904545000" data-alt="Ligand field strength series diagram; illustrates I⁻ < Br⁻ < Cl⁻ < F⁻ < OH⁻ < H₂O < NH₃ < en < CN⁻

Figure 2. Spectrochemical series.

For example, a coordination compound of the Cu⁺ ion has a d¹⁰ configuration, and all the e_g 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.

Magnetism in Coordination Complexes

Experimental evidence of magnetic measurements supports the theory of high- and low-spin complexes. Molecules such as O₂ that contain unpaired electrons are paramagnetic. Paramagnetic substances are attracted to magnetic fields. Many transition metal complexes have unpaired electrons and hence are paramagnetic. Molecules such as N₂ and ions such as Na⁺ and [Fe(CN)₆]⁴⁻ that contain no unpaired electrons are diamagnetic. Diamagnetic substances have a slight tendency to be repelled by magnetic fields.

Figure 3. Orbital diagrams of the octahedral complexes in the high and low spin state for the d⁴, d⁵, d⁶, and d⁷ systems. This distinction can not be made for d¹, d², d³, d⁵, d⁸,d⁹ and d¹⁰ systems.

When an electron in an atom or ion is unpaired, the magnetic moment due to its spin makes the entire atom or ion paramagnetic. The size of the magnetic moment of a system containing unpaired electrons is related directly to the number of such electrons: the greater the number of unpaired electrons, the larger the magnetic moment. Therefore, the observed magnetic moment is used to determine the number of unpaired electrons present. The measured magnetic moment of low-spin d⁶ [Fe(CN)₆]⁴⁻ confirms that iron is diamagnetic, whereas high-spin d⁶ [Fe(H₂O)₆]²⁺ has four unpaired electrons with a magnetic moment that confirms this arrangement (Figure 2).

This text is adapted from Openstax, Chemistry 2e, Section19.3: Spectroscopic and Magnetic Properties of Coordination Compounds.

Transition metal complexes exhibit a variety of different colors, attributed to the absorption of specific wavelengths of visible light by these compounds.

Light is absorbed when it has the needed energy to excite an electron from a lower energy level to a higher one. Consequently, transition metal complexes generally absorb light matching the crystal field splitting energy, or delta, of the complex, which is typically in the visible light range.

For example, hexafluorocobaltate(III) strongly absorbs red light but minimally absorbs green light, leading it to appear green in color.

Hexaamminecobalt(III), which has a higher delta, strongly absorbs high-energy blue light but minimally absorbs yellow light. Accordingly, hexaamminecobalt(III) appears yellow in color.

The effects of the smaller delta of hexafluorocobaltate(III) are not limited to its color. When delta is low enough, like in hexafluorocobaltate(III), electrons singly occupy the higher-energy orbitals before pairing in the lower-energy orbitals.

Here, delta is smaller compared to the spin-pairing energy — the energy of the electrostatic repulsion between electrons in the same orbital. As such, it is more energetically feasible for electrons to overcome delta and occupy high-energy orbitals than to overcome the spin-pairing energy to pair in the low-energy orbitals.

In contrast, in hexaamminecobalt(III), delta is greater than the spin-pairing energy. Accordingly, electrons pair in the lower-energy orbitals, leaving the higher-energy orbitals vacant, as expected from Hund’s rule.

As a consequence of this difference in electronic distribution, while the Co(III) ion has four unpaired electrons in hexafluorocobaltate(III), it has zero unpaired electrons in hexaamminecobalt(III). Accordingly, the former is classified as a high-spin complex and the latter is labeled as a low-spin complex.

In general, weak-field ligands, which are associated with small values of delta, lead to high-spin complexes, while strong-field ligands, which promote high values of delta, form low-spin complexes.

High-spin and low-spin complexes can exhibit very different magnetic properties. For example, the high-spin hexafluorocobaltate(III) is attracted by a magnet owing to its unpaired electrons and is called paramagnetic.

Meanwhile, the low-spin hexaamminecobalt(III) is repelled by a magnet and labeled as diamagnetic.