Editing Coordination complex (section)

==Electronic properties==
Many of the properties of transition metal complexes are dictated by their electronic structures. The electronic structure can be described by a relatively ionic model that ascribes formal charges to the metals and ligands. This approach is the essence of [[crystal field theory]] (CFT).  Crystal field theory, introduced by [[Hans Bethe]] in 1929, gives a [[quantum mechanics|quantum mechanically]] based attempt at understanding complexes.  But crystal field theory treats all interactions in a complex as ionic and assumes that the ligands can be approximated by negative point charges.

More sophisticated models embrace covalency, and this approach is described by [[ligand field theory]] (LFT) and [[Molecular orbital theory]] (MO). Ligand field theory, introduced in 1935 and built from molecular orbital theory, can handle a broader range of complexes and can explain complexes in which the interactions are [[covalent]]. The chemical applications of [[group theory]] can aid in the understanding of crystal or ligand field theory, by allowing simple, symmetry based solutions to the formal equations.

Chemists tend to employ the simplest model required to predict the properties of interest; for this reason, CFT has been a favorite for the discussions when possible. MO and LF theories are more complicated, but provide a more realistic perspective.

The electronic configuration of the complexes gives them some important properties:

[[File:Copper complex.jpg|thumb|Synthesis of copper(II)-tetraphenylporphyrin, a metal complex, from [[tetraphenylporphyrin]] and [[copper(II) acetate monohydrate]].]]

===Color of transition metal complexes===
Transition metal complexes often have spectacular colors caused by electronic transitions by the absorption of light. For this reason they are often applied as [[Pigment#Physical basis|pigments]]. Most transitions that are related to colored metal complexes are either d–d transitions or [[charge transfer band]]s.  In a d–d transition, an electron in a d&nbsp;orbital on the metal is excited by a photon to another d orbital of higher energy, therefore d–d transitions occur only for partially-filled d-orbital complexes (d<sup>1–9</sup>). For complexes having d<sup>0</sup> or d<sup>10</sup> configuration, charge transfer is still possible even though d–d transitions are not. A charge transfer band entails promotion of an electron from a metal-based orbital into an empty ligand-based orbital ([[Charge transfer complex|metal-to-ligand charge transfer]] or MLCT).  The converse also occurs: excitation of an electron in a ligand-based orbital into an empty metal-based orbital ([[Charge transfer complex|ligand-to-metal charge transfer]] or LMCT). These phenomena can be observed with the aid of electronic spectroscopy; also known as [[UV-Vis]].<ref>{{cite book |last1= Harris |first1= D. |last2= Bertolucci |first2= M. |title= Symmetry and Spectroscopy |year= 1989 |publisher= Dover Publications | isbn= 9780486661445 }}</ref> For simple compounds with high symmetry, the d–d transitions can be assigned using [[Tanabe–Sugano diagram]]s.  These assignments are gaining increased support with [[computational chemistry]].

{| class="wikitable"
|+ Colours of Various Example Coordination Complexes
|-
! &nbsp;
! Fe<sup>2+</sup>
! Fe<sup>3+</sup>
! Co<sup>2+</sup>
! Cu<sup>2+</sup>
! Al<sup>3+</sup>
! Cr<sup>3+</sup>
|-
! [[Metal ions in aqueous solution|Hydrated Ion]]
| style="background: #CBE9AD;" | {{tooltip|2=Hexaaquairon(2+) cation|[Fe(H<sub>2</sub>O)<sub>6</sub>]<sup>2+</sup>|dotted=no}}     <br/>Pale green<br/>{{tooltip|2=Suspended in solution|Solution|dotted=no}}
| style="background: #EAD558;" | {{tooltip|2=Hexaaquairon(3+) cation|[Fe(H<sub>2</sub>O)<sub>6</sub>]<sup>3+</sup>|dotted=no}}    <br/>Yellow/brown<br/>{{tooltip|2=Suspended in solution|Solution|dotted=no}}
| style="background: #FF99CB;" | {{tooltip|2=Hexaaquacobalt(2+) cation|[Co(H<sub>2</sub>O)<sub>6</sub>]<sup>2+</sup>|dotted=no}}<br/>Pink<br/>{{tooltip|2=Suspended in solution|Solution|dotted=no}}
| style="background: #C7D9F1;" | {{tooltip|2=Hexaaquacopper(2+) cation|[Cu(H<sub>2</sub>O)<sub>6</sub>]<sup>2+</sup>|dotted=no}}<br/>Blue<br/>{{tooltip|2=Suspended in solution|Solution|dotted=no}}
| style="background: #F2F2F2;" | {{tooltip|2=Hexaaquaaluminium(3+) cation|[Al(H<sub>2</sub>O)<sub>6</sub>]<sup>3+</sup>|dotted=no}}<br/>Colourless<br/>{{tooltip|2=Suspended in solution|Solution|dotted=no}}
| style="background: #D7E3BD;" | {{tooltip|2=Hexaaquachromium(3+) cation|[Cr(H<sub>2</sub>O)<sub>6</sub>]<sup>3+</sup>|dotted=no}}<br/>Green<br/>{{tooltip|2=Suspended in solution|Solution|dotted=no}}
|-
! (OH)<sup>−</sup>, dilute
| style="background: #92D14F;" | {{tooltip|2=Tetraaquadihydroxidoiron|[Fe(H<sub>2</sub>O)<sub>4</sub>(OH)<sub>2</sub>]|dotted=no}} <br/>Dark green<br/>{{tooltip|2=Precipitates from solution|Precipitate|dotted=no}}
| style="background: #9B752A;" | {{tooltip|2=Triaquatrihydroxidoiron|[Fe(H<sub>2</sub>O)<sub>3</sub>(OH)<sub>3</sub>]|dotted=no}}<br/>Brown<br/>{{tooltip|2=Precipitates from solution|Precipitate|dotted=no}}
| style="background: #8AE5D6;" | {{tooltip|2=Tetraaquadihydroxidocobalt|[Co(H<sub>2</sub>O)<sub>4</sub>(OH)<sub>2</sub>]|dotted=no}}<br/>Blue/green<br/>{{tooltip|2=Precipitates from solution|Precipitate|dotted=no}}
| style="background: #8EB2E2;" | {{tooltip|2=Tetraaquadihydroxidocopper|[Cu(H<sub>2</sub>O)<sub>4</sub>(OH)<sub>2</sub>]|dotted=no}}<br/>Blue<br/>{{tooltip|2=Precipitates from solution|Precipitate|dotted=no}}
| style="background: #FFFFFF;" | {{tooltip|2=Triaquatrihydroxidoaluminium|[Al(H<sub>2</sub>O)<sub>3</sub>(OH)<sub>3</sub>]|dotted=no}}<br/>White<br/>{{tooltip|2=Precipitates from solution|Precipitate|dotted=no}}
| style="background: #76923D;" | {{tooltip|2=Triaquatrihydroxidochromium|[Cr(H<sub>2</sub>O)<sub>3</sub>(OH)<sub>3</sub>]|dotted=no}}<br/>Green<br/>{{tooltip|2=Precipitates from solution|Precipitate|dotted=no}}
|-
! (OH)<sup>−</sup>, concentrated
| style="background: #92D14F;" | {{tooltip|2=Tetraaquadihydroxidoiron|[Fe(H<sub>2</sub>O)<sub>4</sub>(OH)<sub>2</sub>]|dotted=no}} <br/>Dark green<br/>{{tooltip|2=Precipitates from solution|Precipitate|dotted=no}}
| style="background: #9B752A;" | {{tooltip|2=Triaquatrihydroxidoiron|[Fe(H<sub>2</sub>O)<sub>3</sub>(OH)<sub>3</sub>]|dotted=no}}<br/>Brown<br/>{{tooltip|2=Precipitates from solution|Precipitate|dotted=no}}
| style="background: #8AE5D6;" | {{tooltip|2=Tetraaquadihydroxidocobalt|[Co(H<sub>2</sub>O)<sub>4</sub>(OH)<sub>2</sub>]|dotted=no}}<br/>Blue/green<br/>{{tooltip|2=Precipitates from solution|Precipitate|dotted=no}}
| style="background: #8EB2E2;" | {{tooltip|2=Tetraaquadihydroxidocopper|[Cu(H<sub>2</sub>O)<sub>4</sub>(OH)<sub>2</sub>]|dotted=no}}<br/>Blue<br/>{{tooltip|2=Precipitates from solution|Precipitate|dotted=no}}
| style="background: #FFFFFF;" | {{tooltip|2=Tetrahydroxidoaluminium|[Al(OH)<sub>4</sub>]<sup>−</sup>|dotted=no}}<br/>Colourless<br/>Solution
| style="background: #C3D59B;" | {{tooltip|2=Hexahydroxidochromate(3−) anion|[Cr(OH)<sub>6</sub>]<sup>3−</sup>|dotted=no}}<br/>Green<br/>{{tooltip|2=Suspended in solution|Solution|dotted=no}}
|-
! NH<sub>3</sub>, dilute
| style="background: #92D14F;" | {{tooltip|2=Hexaammineiron|[Fe(NH<sub>3</sub>)<sub>6</sub>]<sup>2+</sup>|dotted=no}}<br/>Dark green<br/>{{tooltip|2=Precipitates from solution|Precipitate|dotted=no}}
| style="background: #9B752A;" | {{tooltip|2=Hexaammineiron|[Fe(NH<sub>3</sub>)<sub>6</sub>]<sup>3+</sup>|dotted=no}}<br/>Brown<br/> {{tooltip|2=Precipitates from solution|Precipitate|dotted=no}}
| style="background: #D3D359;" | {{tooltip|2=Hexaamminecobalt(2+) cation|[Co(NH<sub>3</sub>)<sub>6</sub>]<sup>2+</sup>|dotted=no}}<br/>Straw coloured<br/>{{tooltip|2=Suspended in solution|Solution|dotted=no}}
| style="background: #548DD4;" | {{tooltip|2=Tetraamminediaquacopper(2+) cation|[Cu(NH<sub>3</sub>)<sub>4</sub>(H<sub>2</sub>O)<sub>2</sub>]<sup>2+</sup>|dotted=no}}<br/>Deep blue<br/>{{tooltip|2=Suspended in solution|Solution|dotted=no}}
| style="background: #FFFFFF;" | {{tooltip|2=Triamminealuminium|[Al(NH<sub>3</sub>)<sub>3</sub>]<sup>3+</sup>|dotted=no}}<br/>White<br/>{{tooltip|2=Precipitates from solution|Precipitate|dotted=no}}
| style="background: #CC0099;" | {{tooltip|2=hexaamminechromium(3+) cation|[Cr(NH<sub>3</sub>)<sub>6</sub>]<sup>3+</sup>|dotted=no}}<br/>Purple<br/>{{tooltip|2=Suspended in solution|Solution|dotted=no}}
|-
! NH<sub>3</sub>, concentrated
| style="background: #92D14F;" | {{tooltip|2=Hexaammineiron|[Fe(NH<sub>3</sub>)<sub>6</sub>]<sup>2+</sup>|dotted=no}}<br/>Dark green<br/>{{tooltip|2=Precipitates from solution|Precipitate|dotted=no}}
| style="background: #9B752A;" | {{tooltip|2=Hexaammineiron|[Fe(NH<sub>3</sub>)<sub>6</sub>]<sup>3+</sup>|dotted=no}}<br/>Brown<br/> {{tooltip|2=Precipitates from solution|Precipitate|dotted=no}}
| style="background: #D3D359;" | {{tooltip|2=Hexaamminecobalt(2+) cation|[Co(NH<sub>3</sub>)<sub>6</sub>]<sup>2+</sup>|dotted=no}}<br/>Straw coloured<br/>{{tooltip|2=Suspended in solution|Solution|dotted=no}}
| style="background: #548DD4;" | {{tooltip|2=Tetraamminediaquacopper(2+) cation|[Cu(NH<sub>3</sub>)<sub>4</sub>(H<sub>2</sub>O)<sub>2</sub>]<sup>2+</sup>|dotted=no}}<br/>Deep blue<br/>{{tooltip|2=Suspended in solution|Solution|dotted=no}}
| style="background: #FFFFFF;" | {{tooltip|2=Triamminealuminium|[Al(NH<sub>3</sub>)<sub>3</sub>]<sup>3+</sup>|dotted=no}}<br/>White<br/>{{tooltip|2=Precipitates from solution|Precipitate|dotted=no}}
| style="background: #CC0099;" | {{tooltip|2=hexaamminechromium(3+) cation|[Cr(NH<sub>3</sub>)<sub>6</sub>]<sup>3+</sup>|dotted=no}}<br/>Purple<br/>{{tooltip|2=Suspended in solution|Solution|dotted=no}}
|-
! (CO<sub>3</sub>)<sup>2-</sup>
| style="background: #92D14F;" | {{tooltip|2=Iron(II) carbonate|FeCO<sub>3</sub>|dotted=no}}<br/>Dark green<br                            />{{tooltip|2=Precipitates from solution|Precipitate|dotted=no}}
| style="background: #9B752A;" | {{tooltip|2=Iron(III) carbonate|Fe<sub>2</sub>(CO<sub>3</sub>)<sub>3</sub>|dotted=no}}<br/>Brown<br                            />{{tooltip|2=Precipitates from solution|Precipitate+bubbles|dotted=no}}
| style="background: #FF99CB;"| {{tooltip|2=Cobalt(II) carbonate|CoCO<sub>3</sub>|dotted=no}}<br/>Pink<br/>{{tooltip|2=Precipitates from solution with an effervescence of Carbon Dioxide|Precipitate|dotted=no}}
| style="background: #8AE5D6;" | {{tooltip|2=Copper(II) carbonate|CuCO<sub>3</sub>|dotted=no}}<br/>Blue/green<br/>{{tooltip|2=Precipitates from solution|Precipitate|dotted=no}}
|}

===Colors of lanthanide complexes===
Superficially [[lanthanide]] complexes are similar to those of the transition metals in that some are colored. However, for the common Ln<sup>3+</sup> ions (Ln = lanthanide) the colors are all pale, and hardly influenced by the nature of the ligand. The colors are due to 4f electron transitions. As the 4f orbitals in lanthanides are "buried" in the xenon core and shielded from the ligand by the 5s and 5p orbitals they are therefore not influenced by the ligands to any great extent leading to a much smaller [[Crystal field theory|crystal field]] splitting than in the transition metals. The absorption spectra of an Ln<sup>3+</sup> ion approximates to that of the free ion where the electronic states are described by [[Angular momentum coupling#Spin-orbit coupling|spin-orbit coupling]]. This contrasts to the transition metals where the ground state is split by the crystal field. Absorptions for Ln<sup>3+</sup> are weak as electric dipole transitions are parity forbidden ([[Laporte rule|Laporte forbidden]]) but can gain intensity due to the effect of a low-symmetry ligand field or mixing with higher electronic states (''e.g.'' d orbitals). f-f absorption bands are extremely sharp which contrasts with those observed for transition metals which generally have broad bands.<ref name = "C&W6th">{{Cotton&Wilkinson6th}}</ref><ref name=CottonSA2006>{{cite book |last=Cotton |first=Simon |year=2006 |title=Lanthanide and Actinide Chemistry|publisher= John Wiley & Sons Ltd}}</ref> This can lead to extremely unusual effects, such as significant color changes under different forms of lighting.

===Magnetism===
{{main|magnetochemistry}}
Metal complexes that have unpaired electrons are [[paramagnetic]]. This can be due to an odd number of electrons overall, or to incomplete electron-pairing.  Thus, monomeric Ti(III) species have one "d-electron" and must be [[paramagnetism|(para)magnetic]], regardless of the geometry or the nature of the ligands.  Ti(II), with two d-electrons, forms some complexes that have two unpaired electrons and others with none. This effect is illustrated by the compounds TiX<sub>2</sub>[(CH<sub>3</sub>)<sub>2</sub>PCH<sub>2</sub>CH<sub>2</sub>P(CH<sub>3</sub>)<sub>2</sub>]<sub>2</sub>: when X&nbsp;=&nbsp;[[Chlorine|Cl]], the complex is paramagnetic ([[high spin|high-spin]] configuration), whereas when X&nbsp;=&nbsp;[[methyl group|CH<sub>3</sub>]], it is diamagnetic ([[low spin|low-spin]] configuration). Ligands provide an important means of adjusting the [[ground state]] properties.

In bi- and polymetallic complexes, in which the individual centres have an odd number of electrons or that are high-spin, the situation is more complicated.  If there is interaction (either direct or through ligand) between the two (or more) metal centres, the electrons may couple ([[Antiferromagnetism|antiferromagnetic coupling]], resulting in a diamagnetic compound), or they may enhance each other ([[Ferromagnetism|ferromagnetic coupling]]).  When there is no interaction, the two (or more) individual metal centers behave as if in two separate molecules.

===Reactivity===
Complexes show a variety of possible reactivities:<ref>R. G. Wilkins Kinetics and Mechanism of Reactions of Transition Metal Complexes, 2nd Edition, VCH, Weinheim, 1991. {{ISBN|1-56081-125-0}}</ref>

* Electron transfers
*: [[Electron transfer]] (ET) between metal ions can occur via two distinct mechanisms, [[Inner sphere electron transfer|inner]] and [[outer sphere electron transfer]]s. In an inner sphere reaction, a [[bridging ligand]] serves as a conduit for ET.
* (Degenerate) [[ligand exchange]]
*: One important indicator of reactivity is the rate of degenerate exchange of ligands.  For example, the rate of interchange of coordinate water in [M(H<sub>2</sub>O)<sub>6</sub>]<sup>''n''+</sup> complexes varies over 20 orders of magnitude.  Complexes where the ligands are released and rebound rapidly are classified as labile.  Such labile complexes can be quite stable thermodynamically.  Typical labile metal complexes either have low-charge (Na<sup>+</sup>), electrons in d-orbitals that are [[antibonding]] with respect to the ligands (Zn<sup>2+</sup>), or lack covalency (Ln<sup>3+</sup>, where Ln is any lanthanide).  The lability of a metal complex also depends on the high-spin vs. low-spin configurations when such is possible.  Thus, high-spin Fe(II) and Co(III) form labile complexes, whereas low-spin analogues are inert.  Cr(III) can exist only in the low-spin state (quartet), which is inert because of its high formal oxidation state, absence of electrons in orbitals that are M–L antibonding, plus some "ligand field stabilization" associated with the d<sup>3</sup> configuration.
* Associative processes
*: Complexes that have unfilled or half-filled orbitals are often capable of reacting with substrates.  Most substrates have a singlet ground-state; that is, they have lone electron pairs (e.g., water, amines, ethers), so these substrates need an empty orbital to be able to react with a metal centre.  Some substrates (e.g., molecular oxygen) [[triplet oxygen|have a triplet ground state]], which results that metals with half-filled orbitals have a tendency to react with such substrates (it must be said that the [[dioxygen]] molecule also has lone pairs, so it is also capable to react as a 'normal' Lewis base).

If the ligands around the metal are carefully chosen, the metal can aid in ([[stoichiometric]] or [[catalytic]]) transformations of molecules or be used as a sensor.