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. F…
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.
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Q1: Why do transition metal complexes appear in different colors?
Transition metal complexes absorb specific wavelengths of visible light when photons have enough energy to excite electrons between d orbitals. The absorbed light depends on crystal field splitting energy (delta), which determines which colors are absorbed and thus which colors appear to the human eye. For example, hexafluorocobaltate(III) absorbs red light and appears green, while hexaamminecobalt(III) absorbs blue light and appears yellow.
Q2: What is the relationship between ligand strength and complex color?
Strong-field ligands create large crystal field splitting (large delta), causing complexes to absorb high-energy violet or blue light and appear yellow, orange, or red. Weak-field ligands produce small delta values, causing complexes to absorb lower-energy yellow, orange, or red light and appear blue-green, blue, or indigo. The spectrochemical series ranks ligands by their ability to split d orbital energies.
Q3: How does crystal field splitting energy determine whether a complex is high-spin or low-spin?
When delta is smaller than the spin-pairing energy, electrons occupy higher-energy orbitals singly before pairing, creating high-spin complexes with unpaired electrons. When delta exceeds spin-pairing energy, electrons pair in lower-energy orbitals first, forming low-spin complexes with fewer unpaired electrons. Hexafluorocobaltate(III) is high-spin with four unpaired electrons, while hexaamminecobalt(III) is low-spin with zero unpaired electrons.
Q4: What makes a complex paramagnetic or diamagnetic?
Complexes with unpaired electrons are paramagnetic and attracted to magnetic fields. High-spin complexes like hexafluorocobaltate(III) are paramagnetic due to their unpaired electrons. Complexes with no unpaired electrons are diamagnetic and repelled by magnetic fields. Low-spin hexaamminecobalt(III) is diamagnetic because all electrons are paired in lower-energy orbitals.
Q5: Why do different oxidation states of the same metal produce different colored complexes?
Different oxidation states alter the number of d electrons available for excitation and change the crystal field splitting energy. For example, Cu+ complexes with d10 configuration require ultraviolet photons to excite electrons, appearing colorless. Cu2+ complexes have vacancies in d orbitals that allow visible light absorption, appearing blue, blue-green, violet, or yellow depending on the ligands present.
Q6: How does the number of unpaired electrons relate to magnetic moment?
The magnetic moment of a complex is directly proportional to the number of unpaired electrons present. Complexes with more unpaired electrons exhibit larger magnetic moments. Measuring the magnetic moment experimentally confirms the spin state: high-spin d6 [Fe(H2O)6]2+ shows four unpaired electrons with a corresponding magnetic moment, while low-spin d6 [Fe(CN)6]4− is diamagnetic with zero unpaired electrons.
Q7: What determines whether a coordination complex absorbs visible or ultraviolet light?
The energy gap between d orbitals determines the wavelength of absorbed light. When all d orbitals are filled (like Cu+ with d10 configuration), electrons must be excited to much higher orbitals, requiring ultraviolet photons and producing colorless complexes. When d orbital vacancies exist, visible light photons can excite electrons, producing colored complexes. The specific ligands coordinated to the metal center influence this energy gap.