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

High School

Physical Science

HS. Structure and Properties of Matter

HS-PS2-6.

Coordination Number and Geometry

Viele Übergangsmetalle weisen mehrere Oxidationszahlen auf, die zu ihren einzigartigen Eigenschaften, wie z. B. Farben, beitragen. Doch wie wird die Oxidationszahl des Metalls bestimmt?
Koordinationsverbindungen sind elektrisch neutrale Spezies, die aus einem Koordinationskomplex und Gegenionen mit einer primären und sekundären Valenz bestehen.
Die primäre Valenz ist die Oxidationszahl des Metallions. Um die Oxidationszahl zu ermitteln, beginnen Sie mit der Identifizierung der Ladungen, die von den Liganden und Gegenionen beigesteuert werden. 
Als nächstes summieren Sie die Ladungen und bestimmen die Oxidationszahl des Metallions. Sind alle Liganden neutral, wird die komplexe Ionenladung zur Oxidationszahl des Metallions.
Die sekundäre Valenz bezieht sich auf die Anzahl der Liganden, die direkt an das zentrale Metallion gebunden sind, auch Koordinationszahl genannt. Hier beträgt die Koordinationszahl von Rhodium sechs.
Einige Metallionen besitzen nur eine Koordinationszahl. Kobalt(III) und Platin(II) haben eine Koordinationszahl von 6 und 4. Bei vielen Metallionen variiert die Koordinationszahl jedoch zwischen 2 und 6.
Die relative Größe von Liganden und Metallionen beeinflusst die Koordinationszahl. Zum Beispiel koordinieren sich kleinere Liganden wie Fluor sechsmal mit Eisen(III), verglichen mit dem größeren Chlor, das nur viermal koordiniert. 
Negative Ladungen, die durch Liganden an das Metallion abgegeben werden, beeinflussen ebenfalls die Koordinationszahl. Die Koordinationszahl von Nickel(II) mit neutralen Wassermolekülen beträgt 6, die mit anionischen Chloridionen auf 4 reduziert wird.
Die geometrische Form des komplexen Ions hängt teilweise von der Koordinationszahl eines Metallions ab. Ein Komplex mit einer Koordinationszahl von zwei hat eine lineare Geometrie, bei der zwei Liganden auf beiden Seiten des Metallions um 180° voneinander entfernt sind.
Ein Komplex mit einer Koordinationszahl von 4 weist zwei Arten von Geometrien auf, die auf dem Valenzelektron in der d-Unterschale basieren. Metallionen mit acht d-Elektronen, wie z. B. Palladium(II), sind quadratisch planar. Während Metallionen mit zehn d-Elektronen, wie z.B. Zink(II), eine tetraedrische Geometrie aufweisen. 
Ein Komplex mit einer Koordinationszahl von 6 ist oktaedrisch. Die sechs Liganden nehmen sechs Eckpunkte ein, vier Liganden bilden die Ecken eines Quadrats und die restlichen zwei die Ebenen darüber und unten in gleichem Abstand. So erscheint ein Oktaeder aus zwei Pyramiden mit einer gemeinsamen quadratischen Grundfläche und acht Flächen.

koordinationszahl und geometrie

Bei Übergangsmetallkomplexen bestimmt die Koordinationszahl die Geometrie um das zentrale Metallion. Tabelle 1 vergleicht die Koordinationszahlen mit der Molekülgeometrie. Die häufigsten Strukturen der Komplexe in Koordinationsverbindungen sind oktaedrisch, tetraedrisch und quadratisch-planar. 

<table style="margin-left: auto; margin-right: auto;"> 
 <tbody> 

 <tr> 
 <td style="color: black; border: 1px solid black; text-align: center;">Koordinationszahl</td>
 <td style="color: black; border: 1px solid black; text-align: center;">Molekülgeometrie</td>
 <td style="color: black; border: 1px solid black; text-align: center;">Beispiel</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;">lineares</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 planares</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;"> tetraedrisch (d0 oder d10), niedrige Oxidationsstufen für 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;">quadratisch planar (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;">trigonal-bipyramidales</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;">quadratisch-pyramidenförmiges</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;">oktaedrisches</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;">fünfeckige Bipyramide </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;">quadratisches Antiprisma </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;">Dodekaeder</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 und höher</td>
 <td style="color: black; border: 1px solid black; text-align: center;">kompliziertere Strukturen</td>
 <td style="color: black; border: 1px solid black; text-align: center;">[ReH9]2−
 </td>
 </tr> 
 </tbody> 
</table> 

Tabelle 1. Koordinationszahlen und Molekülgeometrie. 

Im Gegensatz zu Hauptgruppenatomen, bei denen sowohl die bindenden als auch die nichtbindenden Elektronen die Molekülform bestimmen, verändern die nichtbindenden d-Elektronen die Anordnung der Liganden nicht. Oktaederkomplexe haben eine Koordinationszahl von sechs und die sechs Donoratome sind an den Ecken eines Oktaeders um das zentrale Metallion angeordnet. Beispiele sind in Figure 1 dargestellt. Die Chlorid- und Nitratanionen in [Co(H2O)6]Cl2 und [Cr(en)3](NO3)3 sowie die Kaliumkationen in K2[PtCl6], stehen außerhalb der Klammern und sind nicht an das Metallion gebunden. 

<img alt="Image1" src="/files/ftp_upload/11457/11457_Image_1.svg" style="text-align: center; height: 200px;"> 

Figure 1. Viele Übergangsmetallkomplexe nehmen oktaedrische Geometrien an, wobei sechs Donoratome mit benachbarten Liganden Bindungswinkel von 90° um das Zentralatom bilden. Beachten Sie, dass nur Liganden innerhalb der Koordinationssphäre die Geometrie um das Metallzentrum beeinflussen. 

Für Übergangsmetalle mit einer Koordinationszahl von vier sind zwei verschiedene Geometrien möglich: tetraedrisch oder quadratisch-planar. In tetraedrischen Komplexen wie [Zn(CN)4]2− (Figure 3) bildet jedes der Ligandenpaare einen Winkel von 109,5°. In quadratisch-planaren Komplexen wie [Pt(NH3)2Cl2] hat jeder Ligand zwei weitere Liganden im 90°-Winkel (die sogenannten cis-Positionen) und einen zusätzlichen Liganden im 180°-Winkel in der trans-Position. 

<img alt="Image2" src="/files/ftp_upload/11457/11457_Image_2.svg" style="text-align: center; height: 200px;"> 

Figure 2. Übergangsmetalle mit einer Koordinationszahl von vier können eine tetraedrische Geometrie (a) wie in K2[Zn(CN)4] oder eine quadratisch-planare Geometrie (b) wie in [Pt(NH3)2Cl2] annehmen. 

Dieser Text wurde angepasst von 
 <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>

Koordinationszahl und Geometrie

Bei Übergangsmetallkomplexen bestimmt die Koordinationszahl die Geometrie um das zentrale Metallion. Tabelle 1 vergleicht die Koordinationszahlen mit ...

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Transition Metals and Coordination Complexes

Übergangsmetalle und Koordinationskomplexe

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KEY TERMS AND CONCEPTS

KERNBEGRIFFE UND KONZEPTE

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...

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

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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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Unless individual gases chemically react with each other, the individual gases in a mixture of gases do not affect each other&rsquo;s pressure. Each gas ...

Analyze data to support the claim that the combined gas law describes the relationships among volume, pressure, and temperature for a sample of an ideal gas.

HS-PS1-9.

Transition Metal Shapes from Coordination Number

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