Name: transition metal shapes from coordination number
Uploaded: 2020-09-24T15:06:53.000Z
Duration: 2 min 57 s
Channel: JoVE

K12 Standards

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New York State Science Learning Standards

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Physical Science

HS. Structure and Properties of Matter

HS-PS2-6.

Coordination Number and Geometry

Many transition metals exhibit multiple oxidation numbers contributing to their unique properties, such as colors. But how is the metal&#8217;s oxidation number determined?
Coordination compounds are electrically neutral species consisting of a coordination complex and counterions with a primary and secondary valence.
The primary valence is the oxidation number of the metal ion. To find the oxidation number start by identifying the charges contributed by the ligands and counter ions. 
Next, sum up the charges and determine the oxidation number of the metal ion. If all ligands are neutral, the complex ion charge becomes the oxidation number of the metal ion.
The secondary valence refers to the number of ligands directly bonded to the central metal ion, also called the coordination number. Here, the coordination number of rhodium is six.
Some metal ions possess one coordination number only. Cobalt(III) and platinum(II) have a coordination number of 6 and 4. However, for many metal ions, the coordination number varies ranging from 2 to 6.
The relative size of ligands and metal ions influences the coordination number. For instance, smaller ligands like fluorine coordinate six times to iron(III), compared to the larger chlorine, which coordinates only four times. 
Negative charges imparted by ligands to the metal ion also influence the coordination number. The coordination number of nickel(II) with neutral water molecules is 6, which is reduced to 4 with anionic chloride ions.
The geometric shape of the complex ion partially depends on the coordination number of a metal ion. A complex with a coordination number of two has a linear geometry, where two ligands are 180&#176; apart on either side of the metal ion.
A complex with a coordination number of 4 exhibits two types of geometry based on the valence electron in the d-subshell. Metal ions with eight d-electrons, such as palladium(II), are square planar. While metal ions with ten d-electrons, such as zinc(II), exhibit a tetrahedral geometry. 
A complex with a coordination number of 6 is octahedral. The six ligands occupy six vertices, four ligands form the corners of a square, and the remaining two the planes above and below at an equivalent distance. Thus, an octahedron appears as two pyramids with a common square base and eight faces.

coordination number and geometry

For transition metal complexes, the coordination number determines the geometry around the central metal ion. Table 1 compares coordination numbers to molecular geometry. The most common structures of the complexes in coordination compounds are octahedral, tetrahedral, and square planar.
<table style="margin-left: auto; margin-right: auto;">
	<tbody>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;"> Coordination Number </td>
			<td style="color: black; border: 1px solid black; text-align: center;">Molecular Geometry</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Example</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">2</td>
			<td style="color: black; border: 1px solid black; text-align: center;">linear</td>
			<td style="color: black; border: 1px solid black; text-align: center;"> [Ag(NH3)2]+ </td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">3</td>
			<td style="color: black; border: 1px solid black; text-align: center;">trigonal planar</td>
			<td style="color: black; border: 1px solid black; text-align: center;">[Cu(CN)3]2&#8722;</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">4</td>
			<td style="color: black; border: 1px solid black; text-align: center;"> tetrahedral(d0 or d10), low oxidation states for M </td>
			<td style="color: black; border: 1px solid black; text-align: center;">[Ni(CO)4]</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">4</td>
			<td style="color: black; border: 1px solid black; text-align: center;">square planar (d8)</td>
			<td style="color: black; border: 1px solid black; text-align: center;">[NiCl4]2&#8722;</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">5</td>
			<td style="color: black; border: 1px solid black; text-align: center;">trigonal bipyramidal</td>
			<td style="color: black; border: 1px solid black; text-align: center;">[CoCl5]2&#8722;</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">5</td>
			<td style="color: black; border: 1px solid black; text-align: center;">square pyramidal</td>
			<td style="color: black; border: 1px solid black; text-align: center;">[VO(CN)4]2&#8722;</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">6</td>
			<td style="color: black; border: 1px solid black; text-align: center;">octahedral</td>
			<td style="color: black; border: 1px solid black; text-align: center;">[CoCl6]3&#8722;</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">7</td>
			<td style="color: black; border: 1px solid black; text-align: center;">pentagonal bipyramid</td>
			<td style="color: black; border: 1px solid black; text-align: center;">[ZrF7]3&#8722;</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">8</td>
			<td style="color: black; border: 1px solid black; text-align: center;">square antiprism</td>
			<td style="color: black; border: 1px solid black; text-align: center;">[ReF8]2&#8722;</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">8</td>
			<td style="color: black; border: 1px solid black; text-align: center;">dodecahedron</td>
			<td style="color: black; border: 1px solid black; text-align: center;">[Mo(CN)8]4&#8722;</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">9 and above</td>
			<td style="color: black; border: 1px solid black; text-align: center;">more complicated structures</td>
			<td style="color: black; border: 1px solid black; text-align: center;">[ReH9]2&#8722;</td>
		</tr>
	</tbody>
</table>
Table 1. Coordination Numbers and Molecular Geometry.
Unlike main group atoms in which both the bonding and nonbonding electrons determine the molecular shape, the nonbonding d-electrons do not change the arrangement of the ligands. Octahedral complexes have a coordination number of six, and the six donor atoms are arranged at the corners of an octahedron around the central metal ion. Examples are shown in Figure 1. The chloride and nitrate anions in [Co(H2O)6]Cl2 and [Cr(en)3](NO3)3, and the potassium cations in K2[PtCl6], are outside the brackets and are not bonded to the metal ion.
<img alt="Coordination chemistry formulas; diagram showing complex ions with ligands and counterions." src="/files/ftp_upload/11457/11457_Image_1.svg" style="text-align: center; height: 200px;" data-altupdatedat="1761830439000" data-alt="Image1" />
Figure 1. Many transition metal complexes adopt octahedral geometries, with six donor atoms forming bond angles of 90&#176; about the central atom with adjacent ligands. Note that only ligands within the coordination sphere affect the geometry around the metal center.
For transition metals with a coordination number of four, two different geometries are possible: tetrahedral or square planar. In tetrahedral complexes such as [Zn(CN)4]2&#8722; (Figure 3), each of the ligand pairs forms an angle of 109.5&#176;. In square planar complexes, such as [Pt(NH3)2Cl2], each ligand has two other ligands at 90&#176; angles (called the cis positions) and one additional ligand at a 180&#176; angle, in the trans position.
<img alt="Coordination chemistry diagrams showing zinc cyanide complex (a) and platinum complex (b) structures." src="/files/ftp_upload/11457/11457_Image_2.svg" style="text-align: center; height: 200px;" data-altupdatedat="1761830439000" data-alt="Image2" />
Figure 2. Transition metals with a coordination number of four can adopt a tetrahedral geometry (a) as in K2[Zn(CN)4] or a square planar geometry (b) as shown in [Pt(NH3)2Cl2].
This text is adapted from <a href="https://openstax.org/books/chemistry-2e/pages/19-2-coordination-chemistry-of-transition-metals">Openstax, Chemistry 2e, Section19.2: </a><a href="https://openstax.org/books/chemistry-2e/pages/19-2-coordination-chemistry-of-transition-metals">Coordination Chemistry of Transition Metals</a><a href="https://openstax.org/books/chemistry-2e/pages/19-2-coordination-chemistry-of-transition-metals">.</a>

For transition metal complexes, the coordination number determines the geometry around the central metal ion. Table 1 compares coordination numbers to ...

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Transition Metals and Coordination Complexes

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KEY TERMS AND CONCEPTS

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 [Cu(CN)3]2−. Four-coordinate complexes can be tetrahedral for d0 or d10 metals and for metals in low oxidation states, as in [Ni(CO)4], or square planar for d8 metals, as in [NiCl4]2−.

A coordination number of 5 can lead to two shapes. These are trigonal bipyramidal, as in [CoCl5]2−, and square pyramidal, as in [VO(CN)4]2−. A coordination number of 6 gives an octahedral geometry, as in [CoCl6]3−. Higher coordination numbers can form more complex shapes, including a pentagonal bipyramid for 7, square antiprism and dodecahedron for 8, and more complicated structures for 9 and above.

Only ligands in the coordination sphere affect the geometry around the metal center. In octahedral complexes, six donor atoms sit at the corners of an octahedron and make 90° angles with adjacent ligands. Examples include [Co(H2O)6]Cl2, [Cr(en)3](NO3)3, and K2[PtCl6], where the ions outside the brackets are not bonded to the metal ion.

Unlike main group atoms, the nonbonding d electrons in transition metals do not change ligand arrangement. For coordination number 4, tetrahedral and square planar geometries are both possible. In tetrahedral complexes such as K2[Zn(CN)4], the ligand pairs form 109.5° angles. In square planar complexes such as [Pt(NH3)2Cl2], ligands have cis positions at 90° and a trans position at 180°.

Transition Metal Shapes from Coordination Number

transition-metal-shapes-from-coordination-number

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...

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

crystal-field-theory-octahedral-complexes

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 coordination number predicts the shapes of transition metal complexes, including linear, tetrahedral, square planar, octahedral, and other geometries.

Transition Metal Shapes from Coordination Number | JoVE K12

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Transition Metal Shapes from Coordination Number

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