Name: color and spin in coordination complexes
Uploaded: 2020-09-24T15:06:53.000Z
Duration: 3 min 2 s
Channel: JoVE

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 &#8212; 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&#8217;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.

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.
<img alt="Absorption spectrum diagram, representing spectral distribution and light absorption across wavelengths." src="/files/ftp_upload/11463/11463_Image_1.png" style="text-align: center; height: 300px;" data-altupdatedat="1754904545000" data-alt="Absorption spectrum diagram showing light wavelength distribution in spectroscopic analysis." />
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:
&#8226; Different aqueous metal ions can have different colors. 
&#8226; Different oxidation states of one metal can produce different colors. 
&#8226; 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 &#916;). 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 (&#916;).
<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&#8315; &#60; Br&#8315; &#60; Cl&#8315; &#60; F&#8315; &#60; OH&#8315; &#60; H&#8322;O &#60; NH&#8323; &#60; en &#60; CN&#8315; &#60; CO" />
Figure 2. Spectrochemical series.
For example, 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]&#8722;, 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&#8212;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&#8722; that contain no unpaired electrons are diamagnetic. Diamagnetic substances have a slight tendency to be repelled by magnetic fields.
<img alt="Static equilibrium diagram showing various block arrangements illustrating balance principles." src="/files/ftp_upload/11463/11463_Image_3.png" style="text-align: center; height: 300px;" data-altupdatedat="1754904545000" data-alt="Binary fission diagram illustrating prokaryotic cell division stages with DNA replication." />
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&#8722; 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 <a href="https://openstax.org/books/chemistry-2e/pages/19-3-spectroscopic-and-magnetic-properties-of-coordination-compounds">Openstax, Chemistry 2e, Section19.3: Spectroscopic and Magnetic Properties of Coordination Compounds.</a>

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 ...

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Transition Metals and Coordination Complexes

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color and spin in coordination complexes

Color in coordination complexes comes from how metal electrons absorb light. When atoms or molecules absorb light at the right frequency, electrons move to higher-energy orbitals. For many main-group atoms and molecules, the absorbed light is ultraviolet and cannot be seen. In coordination compounds, the energy gap between d orbitals often matches visible light, so the compound can absorb and emit colors that the human eye detects.

Small changes in orbital energy can cause large shifts in the color that is absorbed. That is why coordination compounds can change color depending on the metal ion in water, the oxidation state of the metal, and the ligands attached to the metal center. Ligands are molecules or ions that bind to the metal. The iron(II) complex [Fe(H2O)6]SO4 is blue-green because the high-spin complex absorbs red light. The low-spin iron(II) complex K4[Fe(CN)6] is pale yellow because it absorbs higher-energy violet light.

Ligand strength also affects color through crystal field splitting, the size of the energy gap between d orbitals. Strong-field ligands create a large split, and these transition metal complexes often look yellow, orange, or red because they absorb violet or blue light. Weak-field ligands create a smaller split, and these complexes often look blue-green, blue, or indigo because they absorb yellow, orange, or red light. The spectrochemical series lists ligands in order of increasing crystal field splitting energy.

Some coordination compounds do not absorb visible light at all. A Cu+ complex has a d10 electron configuration, and the eg orbitals are filled. To move an electron to a higher level such as a 4p orbital, the compound would need very high-energy ultraviolet light. Since no visible light is absorbed, the compound appears white or colorless. A solution of [Cu(CN)2]- is one example. In contrast, octahedral Cu2+ complexes have a vacancy in the eg orbitals, so electrons can be excited there. These complexes are usually colored, often blue, blue-green, violet, or yellow.

Magnetism gives another way to study coordination complexes. Molecules such as O2 with unpaired electrons are paramagnetic, which means they are attracted to magnetic fields. Many transition metal complexes are also paramagnetic because they contain unpaired electrons. Molecules such as N2 and ions such as Na+ and [Fe(CN)6]4- have no unpaired electrons, so they are diamagnetic and show a slight repulsion by magnetic fields.

Magnetic measurements can support the idea of high-spin and low-spin complexes. This comparison is useful for d4, d5, d6, and d7 octahedral systems, but not for d1, d2, d3, d5, d8, d9, and d10 systems. When an electron is unpaired, its spin creates a magnetic moment, and the size of that moment depends on how many unpaired electrons are present. The measured magnetic moment can therefore help determine electron pairing. For example, low-spin d6 [Fe(CN)6]4- is diamagnetic, while high-spin d6 [Fe(H2O)6]2+ has four unpaired electrons and a magnetic moment that matches that arrangement.

Color and Spin in Coordination Complexes

properties of transition metals

Transition Metals in the Periodic Table

Transition metals are the d-block elements and the inner transition metals in the f-block. They are defined by partially filled d orbitals. In the periodic table, the transition elements are in groups 3 through 11, and the inner transition metals appear in the two rows below the main body of the table. The d-block is often grouped into transition series. The first transition series runs from Sc through Cu. The second series runs from Y through Ag. The third series includes La and Hf through...

coordination compounds and nomenclature

How Transition Metal Complexes Are Named

Transition metal complexes are named by following rules for ligands, metal ions, and oxidation state. These compounds form when a central metal ion or atom, often a transition metal or inner transition metal, accepts electron pairs from ligands. A ligand is a Lewis base that can be an atom, molecule, or ion with one or more lone pairs to donate. In these complexes, the bond is a coordinate covalent bond. That means both electrons in the bond come from the donor, such as ammonia, and are...

metal-ligand bonds

Coordination Complexes and Chelation

Coordination complexes form when a metal ion bonds to ligands. These compounds appear in familiar examples such as hemoglobin in blood, chlorophyll in green plants, vitamin B-12, and the catalyst used to make polyethylene. Transition metal ions are especially likely to form these complexes. In a coordination complex, the metal acts as a Lewis acid. The ligand acts as a Lewis base and donates both electrons in the coordinate covalent bond. A ligand can be an atom, molecule, or ion, as long as...

coordination number and geometry

Transition Metal Shapes from Coordination Number

Transition metal complexes are often identified by their coordination number and the shape around the metal ion. The coordination number is the number of donor atoms bonded to the central metal. In many complexes, the most common shapes are octahedral, tetrahedral, and square planar. Coordination number and molecular geometry are linked in a clear pattern. A coordination number of 2 gives a linear shape, as in [Ag(NH3)2]+. A coordination number of 3 gives a trigonal planar shape, as in...

structural isomerism

Linkage and Ionization Isomers

Isomers are different chemical species that share the same chemical formula. In coordination compounds, structural isomerism includes linkage isomers and coordination-sphere isomers. These two forms show how the same atoms can be arranged in different ways around a metal center. Linkage isomers form when a ligand can attach to a transition metal through two different atoms. The cyanide ligand, CN−, can bind through carbon or nitrogen. The thiocyanate ligand, SCN−, can bind through sulfur or...

stereoisomerism

Isomer Shapes in Coordination Complexes

Isomer shapes in coordination complexes can change a compound’s properties even when the chemical formula stays the same. Isomers are different chemical species with the same formula. In transition metal complexes, the atoms stay connected through the same types of bonds, but their arrangement in space can differ. One common form is geometric isomerism. In these complexes, ligands (the ions or molecules attached to the metal) can sit in cis or trans positions. The octahedral [Co(NH3)4Cl2]+ ion...

valence bond theory

Metal Ligand Bonds in Coordination Complexes

Coordination complexes form metal-ligand bonds through valence bond theory. In this model, Linus Pauling explained bonding with hybridization and overlap of atomic orbitals. The central metal atom or ion acts as a Lewis acid. It uses empty orbitals of suitable energy to accept electron pairs from ligands, which act as Lewis bases. These coordinate covalent bonds help explain the color, geometry, and magnetic behavior seen in coordination compounds. The type of hybridization helps determine the...

crystal field theory - octahedral complexes

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Octahedral d-Orbital Splitting in Crystal Fields

Crystal field theory explains how ligands affect the d orbitals in transition metal complexes. This electrostatic model helps predict the colors, magnetic behavior, and some structures of coordination compounds. It focuses on the nonbonding electrons on the central metal ion, not on the metal-ligand bonds. In crystal field theory, the metal ion and ligands are treated as very small point charges. Because electrons are negative, electrons donated by the ligands repel the electrons in the...

crystal field theory - tetrahedral and square planar complexes

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Crystal Field Splitting in Non-Octahedral Shapes

Crystal field theory also explains d orbital splitting in tetrahedral and square planar complexes. These shapes do not follow the same pattern as an octahedral complex, so the d orbitals respond to ligand placement in different ways. In tetrahedral complexes, the four ligands sit between the axes. None of the d orbitals points directly at the ligands, but the d x2−y2 and d z2 orbitals lie along the Cartesian axes and overlap less with the ligands than the d xy, d xz, and d yz orbitals. The d...

Color in coordination complexes comes from how metal electrons absorb light. When atoms or molecules absorb light at the right frequency, electrons move to higher-energy orbitals. For many main-group atoms and molecules, the absorbed light is ultraviolet and cannot be seen. In coordination compounds, the energy gap between d orbitals often matches visible light, so the compound can absorb and emit colors that the human eye detects. Small changes in orbital energy can cause large shifts in the...

key terms and concepts

Learn how coordination complexes get their colors, why ligands change light absorption, and how magnetic measurements reveal unpaired electrons in transition metal complexes.

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Color and Spin in Coordination Complexes

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