Name: metal ligand bonds in coordination complexes
Uploaded: 2020-09-24T15:06:53.000Z
Duration: 2 min 42 s
Channel: JoVE

Coordination complexes exhibit different geometries like octahedral, tetrahedral, and square planar, depending on the metal orbitals participating in the coordinate covalent bonds.
Valence bond theory, or VBT, explains this based on the metal&#8217;s hybridization of s, p, and d orbitals. Hybridization provides vacant orbitals of equivalent energies that can coordinate to the filled ligand orbitals.
Consider an octahedral complex, hexafluorocobaltate(III). The Co3+ ion has partially filled 3d6 orbitals and vacant 4s, 4p, and 4d orbitals. 
To form the complex, the metal&#39;s six empty orbitals, 4s, 4p, and two 4d, are hybridized to six sp3d2 orbitals, which are pointing towards the corners of an octahedron. The hybrid orbitals accept lone pairs from the six fluoride groups to form the paramagnetic complex.
In the hexamminecobalt(III), the amino groups cause the 3d electrons of cobalt to rearrange, creating two vacant 3d orbitals. 
These combine with the 4s and 4p orbitals to generate six d2sp3 hybrid orbitals, which then coordinately bind to the six ammonia groups to form the diamagnetic complex. 
Now, consider a tetrahedral complex &#8212; tetrachloronickelate. Ni2+ ion has a 3d8 configuration. The empty 4s and 4p orbitals hybridize to give four sp3 orbitals that point towards the corners of a tetrahedron. 
Here, four pairs of electrons, one from each chloride group, occupy the hybrid orbitals to form the paramagnetic complex. 
Lastly, examine a square planar complex like tetrachloroplatinate. The Pt2+ ion has a d8 configuration.
The chloride groups force the metal&#8217;s 3d electrons to rearrange, creating a vacant orbital. The vacant orbital combines with the 4s and two 4p orbitals to give four dsp2 hybrid orbitals directed towards the corners of a square.
Upon accepting the electron pairs from the chloride ligands, the diamagnetic complex is formed. 
Although the coordination complexes vary in their colors and magnetic behavior, VBT does not explain their electronic spectra and the variation in their magnetic behavior with temperature.

valence bond theory

Coordination compounds and complexes exhibit different colors, geometries, and magnetic behavior, depending on the metal atom/ion and ligands from which they are composed. In an attempt to explain the bonding and structure of coordination complexes, Linus Pauling proposed the valence bond theory, or VBT, using the concepts of hybridization and the overlapping of the atomic orbitals. According to VBT, the central metal atom or ion (Lewis acid) hybridizes to provide empty orbitals of suitable energy. These orbitals accept electron pairs from filled ligand orbitals (Lewis bases) to form coordinate covalent metal-ligand bonds. The type of hybridization and the number of hybrid orbitals determine the geometry of the complex.
<table style="margin-left: auto; margin-right: auto;">
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			<td style="color: black; border: 1px solid black; text-align: center;">Geometry</td>
			<td style="color: black; border: 1px solid black; text-align: center;"> Hybridization </td>
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			<td style="color: black; border: 1px solid black; text-align: center;">Linear</td>
			<td style="color: black; border: 1px solid black; text-align: center;">sp</td>
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			<td style="color: black; border: 1px solid black; text-align: center;">Tetrahedral</td>
			<td style="color: black; border: 1px solid black; text-align: center;">sp3</td>
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			<td style="color: black; border: 1px solid black; text-align: center;"> Square planar </td>
			<td style="color: black; border: 1px solid black; text-align: center;">dsp2</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Octahedral</td>
			<td style="color: black; border: 1px solid black; text-align: center;"> d2sp3 or sp3d2 </td>
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	</tbody>
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In a tetrahedral complex, three vacant p orbitals and one vacant s orbital on the metal hybridize to form four sp3 hybrid orbitals, which overlap with the filled ligand orbitals to form the covalent coordinate bond. Similarly, six hybrid orbitals are created for the octahedral complexes by mixing the vacant atomic orbitals on the central metal ion (d2sp3 or sp3d2 hybridization). In the case of linear complexes, the one s and one p orbitals overlap, leading to the formation of two sp hybrid orbitals.
Inner and Outer Orbital Complexes
The strength of the approaching ligands influence the hybridization of the atomic orbitals on the central metal ion. Consider the example of an octahedral complex such as [Co(NH3)6]3+. The Co3+ ion contains six electrons in the 3d orbitals and has vacant 4s and 4p orbitals. The incoming NH3 ligands, which are strong field ligands, force the unpaired 3d electrons to rearrange and pair up with the other 3d electrons. This creates two vacant 3d orbitals, which combine with one 4s and three 4p orbitals to form six equivalent d2sp3 hybrid orbitals. The six hybrid orbitals overlap with the filled atomic orbitals of the ammonia ligands to form the octahedral complex. Since the inner d (3d) orbitals on the metal participate in hybridization, [Co(NH3)6]3+ is an inner orbital complex. Due to the absence of unpaired electrons, the complex is diamagnetic, or called a low spin complex.
<img alt="Static equilibrium equations &#931;Fx=0 &#931;Fy=0; structural analysis diagram with force vectors." src="/files/ftp_upload/11460/11460_Image_1.png" style="text-align: center; height: 100px;" data-altupdatedat="1761830439000" data-alt="Image1" />
 In another octahedral complex like [Co(F)6]3+, since the fluoride ligand is a weak field ligand, the 3d6 electrons of the metal do not rearrange. To provide vacant orbitals for hybridization, two of the outermost empty 4d orbitals combine with the one 4s and three 4p orbitals to form six vacant hybrid orbitals. Since the outermost d orbitals are used, the hybridization is referred to as sp3d2 hybridization, and the complex is called an outer orbital complex. The presence of unpaired electrons makes the complex paramagnetic, and hence these complexes are also known as high spin complexes.
<img alt="Capacitance-voltage characteristic diagram with semiconductor interface and ideal MOS capacitor curve." src="/files/ftp_upload/11460/11460_Image_2.png" style="text-align: center; height: 100px;" data-altupdatedat="1761830439000" data-alt="Image2" />
High-spin or outer orbital complexes are more labile and less stable (due to the higher energies of sp3d2 orbitals) compared to the low-spin or inner orbital complexes.

Valence Bond Theory

Coordination compounds and complexes exhibit different colors, geometries, and magnetic behavior, depending on the metal atom/ion and ligands from which ...

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Transition Metals and Coordination Complexes

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KEY TERMS AND CONCEPTS

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 shape of the complex. A linear complex uses sp hybridization. A tetrahedral complex uses sp3 hybridization. A square planar complex uses dsp2 hybridization. Octahedral complexes use either d2sp3 or sp3d2 hybridization. In a tetrahedral complex, three vacant p orbitals and one vacant s orbital on the metal mix to make four sp3 hybrid orbitals. These overlap with filled ligand orbitals to form the bond. In a linear complex, one s orbital and one p orbital overlap to form two sp hybrid orbitals. Octahedral complexes form six hybrid orbitals from the vacant orbitals on the metal ion.

Ligand strength can change which orbitals the metal uses. A strong field ligand can force electrons to pair up, while a weak field ligand may not. The octahedral complex [Co(NH3)6]3+ is a good example. The Co3+ ion has six electrons in the 3d orbitals and vacant 4s and 4p orbitals. The NH3 ligands are strong field ligands, so the unpaired 3d electrons rearrange and pair. This creates two vacant 3d orbitals. These join with one 4s orbital and three 4p orbitals to make six d2sp3 hybrid orbitals. The ammonia ligands then overlap with these orbitals to form the octahedral complex.

Because the inner 3d orbitals take part in hybridization, [Co(NH3)6]3+ is an inner orbital complex. It has no unpaired electrons, so it is diamagnetic. It is also called a low spin complex. A similar octahedral complex, [Co(F)6]3+, behaves differently because fluoride is a weak field ligand. In that case, the 3d6 electrons do not rearrange. Two outer 4d orbitals combine with one 4s and three 4p orbitals to form six sp3d2 hybrid orbitals. Since outer d orbitals are used, this is an outer orbital complex. The unpaired electrons make it paramagnetic, so it is a high spin complex. High spin, or outer orbital, complexes are also more labile and less stable than low spin, inner orbital complexes.

Metal Ligand Bonds in Coordination Complexes

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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...

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

crystal-field-theory-tetrahedral-and-square-planar-complexes

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...

colors and magnetism

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...

key terms and concepts

Learn how valence bond theory explains metal-ligand bonding, hybridization, complex shape, and the difference between inner and outer orbital complexes.

Metal Ligand Bonds in Coordination Complexes | JoVE K12

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Metal Ligand Bonds in Coordination Complexes

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