Name: transition metal shapes from coordination number
Uploaded: 2020-09-24T15:06:53.000Z
Duration: 2 min 57 s
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HS-PS2-6.

Coordination Number and Geometry

De nombreux métaux de transition présentent plusieurs nombres d'oxydation contribuant à leurs propriétés uniques, telles que les couleurs. Mais comment est déterminé l'indice d'oxydation du métal? Les composés de coordination sont des espèces électriquement neutres consistant en un complexe de coordination et contre-ions avec une valence primaire et secondaire.La valence primaire est le nombre d'oxydation de l'ion métallique. Pour trouver le numéro d'oxydation, commencez en identifiant les charges apportées par les ligands et contre-ions. Ensuite, additionnez les charges et déterminez le nombre d'oxydation de l'ion métallique.Si tous les ligands sont neutres, la charge ionique complexe devient le nombre d'oxydation de l'ion métallique. La valence secondaire fait référence au nombre de ligands directement liés à l'ion métallique central, également appelé le numéro de coordination. Ici, le numéro de coordination du rhodium est de six.Certains ions métalliques ne possèdent qu'un seul numéro de coordination. Le cobalt et le platine ont un numéro de coordination de 6 et 4. Cependant, pour de nombreux ions métalliques, le nombre de coordination varie de 2 à 6.La taille relative des ligands et des ions métalliques influence le numéro de coordination. Par exemple, des ligands plus petits comme le fluor coordonne six fois de fer par rapport au chlore plus gros, qui coordonne seulement quatre fois. Les charges négatives transmises par les ligands à l'ion métallique influencent également le numéro de coordination.Le nombre de coordination du nickel avec de molécules de l'eau neutre est de 6, qui est réduit à 4 avec d'ions de chlorure anionique. La forme géométrique de l'ion complexe dépend en partie du numéro de coordination d'un ion métallique. Un complexe avec un numéro de coordination de deux a une géométrie linéaire, où deux ligands sont distants de 180 degres de chaque côté de l'ion métallique.Un complexe avec un numéro de coordination de 4 présente deux types de géométrie basée sur l'électron de valence dans la sous-classe d. Les ions métalliques avec huit électrons d, comme le palladium sont plans carrés. Alors que les ions métalliques avec dix électrons d, comme le zinc présentent une géométrie tétraédrique.Un complexe avec un nombre de coordination de 6 est octaédrique. Les six ligands occupent six sommets, quatre ligands forment les coins d'un carré, et les deux autres les étages supérieurs et inférieurs à une distance équivalente. Ainsi, un octaèdre apparaît comme deux pyramides avec une base carrée commune et huit faces.

nombre de coordination et géométrie

Pour les complexes de métaux de transition, le nombre de coordination détermine la géométrie autour de l'ion métallique central. Le tableau 1 compare les nombres de coordination à la géométrie moléculaire. Les structures les plus courantes des complexes dans les composés de coordination sont octaédrique, tétraédrique et carrée plane. 
 <table style="margin-left: auto; margin-right: auto;">
 <tbody>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;"> Nombre de coordination  </td>
 <td style="color: black; border: 1px solid black; text-align: center;">Géométrie moléculaire </td>
 <td style="color: black; border: 1px solid black; text-align: center;">Exemple </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;">Linéaire </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;">Plane trigonale </td>
 <td style="color: black; border: 1px solid black; text-align: center;">[Cu(CN)3]2− </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;"> Tétraédrique (d0 ou d10), faibles états d'oxydation de 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;">Plane carrée (d8) </td>
 <td style="color: black; border: 1px solid black; text-align: center;">[NiCl4]2−</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;">Bipyramide trigonale </td>
 <td style="color: black; border: 1px solid black; text-align: center;">[CoCl5]2− </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;">Pyramide carrée </td>
 <td style="color: black; border: 1px solid black; text-align: center;">[VO(CN)4]2− </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;">Octaédrique </td>
 <td style="color: black; border: 1px solid black; text-align: center;">[CoCl6]3− </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;">Bipyramide pentagonale </td>
 <td style="color: black; border: 1px solid black; text-align: center;">[ZrF7]3− </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;">Antiprisme carré </td>
 <td style="color: black; border: 1px solid black; text-align: center;">[ReF8]2− </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;">Dodécaèdre </td>
 <td style="color: black; border: 1px solid black; text-align: center;">[Mo(CN)8]4− </td>
 </tr>
 <tr>
 <td style="color: black; border: 1px solid black; text-align: center;">9 et plus </td>
 <td style="color: black; border: 1px solid black; text-align: center;">Structures plus compliquées </td>
 <td style="color: black; border: 1px solid black; text-align: center;">[ReH9]2− </td>
 </tr>
 </tbody>
 </table>
 Tableau 1. Nombres de coordination et géométrie moléculaire. 
 Contrairement aux atomes du groupe principal dans lesquels les électrons liants et les électrons non-liants déterminent la forme moléculaire, les électrons d non-liants ne modifient pas la disposition des ligands. Les complexes octaédriques ont un nombre de coordination de six, et les six atomes donneurs sont disposés aux coins d'un octaèdre autour de l'ion métallique central. Des exemples sont présentés dans la figure 1. Les anions chlorure et nitrate dans [Co(H2O)6]Cl2 et [Cr(en)3](NO3)3, et les cations potassium dans K2[PtCl6], sont à l'extérieur des crochets et ne sont pas liés à l'ion métallique. 
 <img alt="Image1" src="/files/ftp_upload/11457/11457_Image_1.svg" style="text-align: center; height: 200px;" />
 Figure 1. De nombreux complexes de métaux de transition adoptent des géométries octaédriques, avec six atomes donneurs formant des angles de liaison de 90° autour de l'atome central avec des ligands adjacents. Notez que seuls les ligands dans la sphère de coordination influencent la géométrie autour du centre métallique. 
 Pour les métaux de transition ayant un nombre de coordination de quatre, deux géométries différentes sont possibles : tétraédrique ou carrée plane. Dans des complexes tétraédriques comme [Zn(CN)4]2− (figure 3), chacune des paires de ligands forme un angle de 109,5°. Dans les complexes plans carrés, comme [Pt(NH3)2Cl2], chaque ligand a deux autres ligands à des angles de 90° (appelés positions cis) et un autre ligand à un angle de 180°, en position trans. 
 <img alt="Image2" src="/files/ftp_upload/11457/11457_Image_2.svg" style="text-align: center; height: 200px;" />
 Figure 2. Les métaux de transition avec un nombre de coordination de quatre peuvent adopter une géométrie tétraédrique (a) comme dans K2[Zn(CN)4] ou une géométrie plane carrée (b) comme indiqué dans [Pt(NH3)2Cl2]. 
 Ce texte est adapté de <a href="https://openstax.org/books/chemistry-2e/pages/19-2-coordination-chemistry-of-transition-metals"> Openstax, Chimie 2e, Section19.2 : </a><a href="https://openstax.org/books/chemistry-2e/pages/19-2-coordination-chemistry-of-transition-metals"> Chimie de coordination des métaux de transition</a><a href="https://openstax.org/books/chemistry-2e/pages/19-2-coordination-chemistry-of-transition-metals">. </a>

Nombre de coordination et géométrie

Pour les complexes de métaux de transition, le nombre de coordination détermine la géométrie autour de l’ion métallique central. Le tableau 1 compare les numéros de coordination à ...

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Transition Metals and Coordination Complexes

Métaux de transition et complexes de coordination

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KEY TERMS AND CONCEPTS

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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

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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 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

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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