Editing Transition metal (section)

===Coloured compounds<span class="anchor" id="Colored compounds"></span>===
[[Image:Coloured-transition-metal-solutions.jpg|thumb|right|250px|From left to right, aqueous solutions of: {{chem|link=cobalt(II) nitrate|Co(NO|3|)|2}} (red); {{chem|link=potassium dichromate|K|2|Cr|2|O|7}} (orange); {{chem|link=potassium chromate|K|2|CrO|4}} (yellow); {{chem|link=nickel(II) chloride|NiCl|2}} (turquoise); {{chem|link=copper(II) sulfate|CuSO|4}} (blue); {{chem|link=potassium permanganate|KMnO|4}} (purple).]]

Colour in transition-series metal compounds is generally due to electronic transitions of two principal types.
*[[Charge transfer complex|charge transfer]] transitions. An electron may jump from a predominantly [[ligand]] [[Atomic orbital|orbital]] to a predominantly metal orbital, giving rise to a ligand-to-metal charge-transfer (LMCT) transition. These can most easily occur when the metal is in a high oxidation state. For example, the colour of [[Chromate ion|chromate]], [[dichromate]] and [[permanganate]] ions is due to LMCT transitions. Another example is that [[mercuric iodide]], HgI<sub>2</sub>, is red because of a LMCT transition.<!--As this example shows, charge transfer transitions are not restricted to transition metals.<ref>{{cite book|author=Dunn, T.M.|editor=Lewis, J. and Wilkins, R.G.|title=Modern Coordination Chemistry|publisher=Wiley Interscience|location=New York|year=1960|pages= Chapter 4, Section 4, "Charge Transfer Spectra", pp. 268–273}}</ref>-->

A metal-to-ligand charge transfer (MLCT) transition will be most likely when the metal is in a low oxidation state and the ligand is easily reduced.

In general charge transfer transitions result in more intense colours than d–d transitions.
*d–d transitions. An electron jumps from one [[d orbital]] to another. In complexes of the transition metals the d orbitals do not all have the same energy. The pattern of splitting of the d orbitals can be calculated using [[crystal field]] theory. The extent of the splitting depends on the particular metal, its oxidation state and the nature of the ligands. The actual energy levels are shown on [[Tanabe–Sugano diagram]]s.

In [[centrosymmetric]] complexes, such as octahedral complexes, d–d transitions are forbidden by the [[Laporte rule]] and only occur because of [[vibronic coupling]] in which a [[molecular vibration]] occurs together with a d–d transition. Tetrahedral complexes have somewhat more intense colour because mixing d and p orbitals is possible when there is no centre of symmetry, so transitions are not pure d–d transitions. The [[molar absorptivity]] (ε) of bands caused by d–d transitions are relatively low, roughly in the range 5-500 M<sup>−1</sup>cm<sup>−1</sup> (where [[Molar concentration|M]] = mol dm<sup>−3</sup>).<ref>{{cite book|last=Orgel|first=L.E.|title=An Introduction to Transition-Metal Chemistry, Ligand field theory|publisher=Methuen|location=London|year=1966|edition=2nd.}}</ref> Some d–d transitions are [[spin forbidden]]. An example occurs in octahedral, high-spin complexes of [[manganese]](II),
which has a d<sup>5</sup> configuration in which all five electrons have parallel spins; the colour of such complexes is much weaker than in complexes with spin-allowed transitions. Many compounds of manganese(II) appear almost colourless. The [[Tanabe–Sugano diagram#Manganese(II) Hexahydrate|spectrum of {{chem|[Mn(H|2|O)|6|]|2+}}]] shows a maximum molar absorptivity of about 0.04 M<sup>−1</sup>cm<sup>−1</sup> in the [[visible spectrum]].