Editing Transition metal (section)

==Characteristic properties==
There are a number of properties shared by the transition elements that are not found in other elements, which results from the partially filled d shell. These include
* the formation of compounds whose colour is due to d–d electronic transitions
* the formation of compounds in many oxidation states, due to the relatively low energy gap between different possible oxidation states<ref>{{cite journal|author=Matsumoto, Paul S|journal=Journal of Chemical Education|year=2005|volume=82|page=1660|title=Trends in Ionization Energy of Transition-Metal Elements|doi=10.1021/ed082p1660 |bibcode = 2005JChEd..82.1660M|issue=11 }}</ref>
* the formation of many [[paramagnetic]] compounds due to the presence of unpaired d electrons. A few compounds of main-group elements are also paramagnetic (e.g. [[nitric oxide]], [[oxygen]])
Most transition metals can be bound to a variety of [[ligands]], allowing for a wide variety of transition metal complexes.<ref>Hogan, C. Michael (2010). [http://www.eoearth.org/article/Heavy_metal?topic=49498 "Heavy metal"] in ''Encyclopedia of Earth''. National Council for Science and the Environment. E. Monosson and C. Cleveland (eds.) Washington DC.</ref>

===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]].

===Oxidation states===
A characteristic of transition metals is that they exhibit two or more [[oxidation state]]s, usually differing by one. For example, compounds of [[vanadium]] are known in all oxidation states between −1, such as {{chem|[V(CO)|6|]|-}}, and +5, such as {{chem|VO|4|3-}}.

[[File:Transition metal oxidation states.svg|frame|center|Oxidation states of the transition metals. The solid dots show common oxidation states, and the hollow dots show possible but unlikely states.]]

[[Main-group element]]s in groups 13 to 18 also exhibit multiple oxidation states. The "common" oxidation states of these elements typically differ by two instead of one. For example, compounds of [[gallium]] in oxidation states +1 and +3 exist in which there is a single gallium atom. Compounds of Ga(II) would have an unpaired electron and would behave as a [[free radical]] and generally be destroyed rapidly, but some stable radicals of Ga(II) are known.<ref>{{cite journal|last1=Protchenko|first1=Andrey V.|last2=Dange|first2=Deepak|last3=Harmer|first3=Jeffrey R.|last4=Tang|first4=Christina Y.|last5=Schwarz|first5=Andrew D.|last6=Kelly|first6=Michael J.|last7=Phillips|first7=Nicholas|last8=Tirfoin|first8=Remi|last9=Birjkumar|first9=Krishna Hassomal|last10=Jones|first10=Cameron|last11=Kaltsoyannis|first11=Nikolas|last12=Mountford|first12=Philip|last13=Aldridge|first13=Simon|title=Stable GaX<sub>2</sub>, InX<sub>2</sub> and TlX<sub>2</sub> radicals|journal=Nature Chemistry|date=16 February 2014|volume=6|issue=4|pages=315–319|doi=10.1038/nchem.1870|pmid=24651198|bibcode = 2014NatCh...6..315P }}</ref> Gallium also has a formal oxidation state of +2 in dimeric compounds, such as {{chem|[Ga|2|Cl|6|]|2-}}, which contain a Ga-Ga bond formed from the unpaired electron on each Ga atom.<ref>{{Greenwood&Earnshaw}} p. 240</ref> Thus the main difference in oxidation states, between transition elements and other elements is that oxidation states are known in which there is a single atom of the element and one or more unpaired electrons.

The maximum oxidation state in the first row transition metals is equal to the number of valence electrons from [[titanium]] (+4) up to [[manganese]] (+7), but decreases in the later elements. In the second row, the maximum occurs with [[ruthenium]] (+8), and in the third row, the maximum occurs with [[iridium]] (+9). In compounds such as {{chem|[MnO|4|]|-}} and {{chem|OsO|4}}, the elements achieve a stable configuration by [[covalent bonding]].

The lowest oxidation states are exhibited in [[metal carbonyl]] complexes such as {{chem|Cr(CO)|6}} (oxidation state zero) and {{chem|[Fe(CO)|4|]|2-}} (oxidation state −2) in which the [[18-electron rule]] is obeyed. These complexes are also covalent.

Ionic compounds are mostly formed with oxidation states +2 and +3. In aqueous solution, the ions are hydrated by (usually) six water molecules arranged octahedrally.
<!-- [[Image:Transition metal oxidation states 3.png|center|frame|This table shows some of the oxidation states found in compounds of the transition-metal elements.<br> A solid circle represents a common oxidation state, and a ring represents a less common oxidation state.]] -->

===Magnetism===
{{main|Magnetochemistry}}
Transition metal compounds are [[paramagnetic]] when they have one or more unpaired d electrons.<ref>{{cite book|last1=Figgis|first1=B.N.|last2=Lewis|first2=J.|title=The Magnetochemistry of Complex Compounds|editor=Lewis, J.|editor2=Wilkins, R.G.|publisher=Wiley Interscience|location=New York|year=1960|series=Modern Coordination Chemistry|pages=400–454}}</ref> In octahedral complexes with between four and seven d electrons both [[high spin]] and [[low spin]] states are possible. Tetrahedral transition metal complexes such as {{chem|[FeCl|4|]|2-}} are [[high spin]] because the crystal field splitting is small so that the energy to be gained by virtue of the electrons being in lower energy orbitals is always less than the energy needed to pair up the spins. Some compounds are [[diamagnetic]]. These include octahedral, low-spin, d<sup>6</sup> and square-planar d<sup>8</sup> complexes. In these cases, [[crystal field]] splitting is such that all the electrons are paired up.

[[Ferromagnetism]] occurs when individual atoms are paramagnetic and the spin vectors are aligned parallel to each other in a crystalline material. Metallic iron and the alloy [[alnico]] are examples of ferromagnetic materials involving transition metals. [[Antiferromagnetism]] is another example of a magnetic property arising from a particular alignment of individual spins in the solid state.

===Catalytic properties===
The transition metals and their compounds are known for their homogeneous and heterogeneous [[catalytic]] activity. This activity is ascribed to their ability to adopt multiple oxidation states and to form complexes. [[Vanadium]](V) oxide (in the [[contact process]]), finely divided [[iron]] (in the [[Haber process]]), and [[nickel]] (in [[Hydrogenation|catalytic hydrogenation]]) are some of the examples. Catalysts at a solid surface ([[nanomaterial-based catalyst]]s) involve the formation of bonds between reactant molecules and atoms of the surface of the catalyst (first row transition metals utilize 3d and 4s electrons for bonding). This has the effect of increasing the concentration of the reactants at the catalyst surface and also weakening of the bonds in the reacting molecules (the activation energy is lowered). Also because the transition metal ions can change their oxidation states, they become more effective as [[Catalysis|catalysts]].

An interesting type of catalysis occurs when the products of a reaction catalyse the reaction producing more catalyst ([[autocatalysis]]). One example is the reaction of [[oxalic acid]] with acidified [[potassium permanganate]] (or manganate (VII)).<ref>{{cite journal | title = Revising the Mechanism of the Permanganate/Oxalate Reaction |vauthors=Kovacs KA, Grof P, Burai L, Riedel M | journal = J. Phys. Chem. A | doi = 10.1021/jp047061u | year = 2004 | volume = 108 | pages = 11026–11031 | issue = 50| bibcode = 2004JPCA..10811026K }}</ref> Once a little Mn<sup>2+</sup> has been produced, it can react with MnO<sub>4</sub><sup>−</sup> forming Mn<sup>3+</sup>. This then reacts with C<sub>2</sub>O<sub>4</sub><sup>−</sup> ions forming Mn<sup>2+</sup> again.

===Physical properties===
As implied by the name, all transition metals are [[metal]]s and thus conductors of electricity.

In general, transition metals possess a high [[density]] and high [[melting point]]s and [[boiling point]]s. These properties are due to [[metallic bond]]ing by delocalized d electrons, leading to [[Cohesion (chemistry)|cohesion]] which increases with the number of shared electrons. However the group 12 metals have much lower melting and boiling points since their full d subshells prevent d–d bonding, which again tends to differentiate them from the accepted transition metals. Mercury has a melting point of {{convert|−38.83|°C|F}} and is a liquid at room temperature.