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Transition metal

The term transition metal (sometimes also called a transition element) has two possible meanings:

* In the past it referred to any element in the d-block of the periodic table, which includes groups 3 to 12 on the periodic table. All elements in the d-block are metals.
* The modern, IUPAC definition[1] states that a transition metal is "an element whose atom has an incomplete d sub-shell, or which can give rise to cations with an incomplete d sub-shell." Group 12 elements are not transition metals in this definition.

The name transition comes from their position in the periodic table of elements. In each of the four periods in which they occur, these elements represent the successive addition of electrons to the d atomic orbitals of the atoms. In this way, the transition metals represent the transition between group 2 elements and group 13 elements.


In the d-block the atoms of the elements have between 1 and 10 d electrons.
Group 3 4 5 6 7 8 9 10 11 12
Period 4 Sc 21 Ti 22 V 23 Cr 24 Mn 25 Fe 26 Co 27 Ni 28 Cu 29 Zn 30
Period 5 Y 39 Zr 40 Nb 41 Mo 42 Tc 43 Ru 44 Rh 45 Pd 46 Ag 47 Cd 48
Period 6 La 57 Hf 72 Ta 73 W 74 Re 75 Os 76 Ir 77 Pt 78 Au 79 Hg 80
Period 7 Ac 89 Rf 104 Db 105 Sg 106 Bh 107 Hs 108 Mt 109 Ds 110 Rg 111 Cn 112

Although atoms of scandium and yttrium have a single d electron in the outermost shell, these elements are not usually considered transition metals as all their compounds contain the ions Sc3+ and Y3+ in which there are no d electrons. Lanthanum is usually considered a lanthanoid element and actinium an actinoid element.[2]

The electronic structure of transition metal atoms can be written, with a few minor exceptions, as [ ]ns2(n-1)dm, as the inner d orbital has more energy than the valence-shell s orbital. In divalent and trivalent ions of the transition metals the situation is reversed so that the s electrons have higher energy. Consequently an ion such as Fe2+ has no s electrons, it has the electronic configuration [Ar]3d6 as compared with the configuration of the atom, [Ar]4s23d6.

Zinc, cadmium, and mercury are not transition metals.[3] as they have the electronic configuration [ ]d10s2, with no incomplete d shell.[4] In the oxidation state +2 the ions have the electronic configuration [ ] d10. While compounds in the +1 oxidation state, such as Hg22+, are known there are no unpaired electrons because of the formation of a covalent bond between the two atoms of the dimer. Zn, Cd and Hg may be classed as post-transition metals. However, it is often convenient to include these elements in a discussion of the transition elements. For example, when discussing the crystal field stabilization energy of first-row transition elements, it is convenient to include the non-transition elements calcium and zinc, as both Ca2+ and Zn2+ have a value of zero against which the value for other transition metal ions may be compared. Another example occurs in the Irving-Williams series of stability constants of complexes.

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 reactivity of unpaired d electrons.[5]
* the formation of many paramagnetic compounds due to the presence of unpaired d electrons. A few compounds of main group elements are paramagnetic (e.g. nitric oxide, oxygen)

Coloured compounds
From left to right, aqueous solutions of: Co(NO3)2 (red); K2Cr2O7 (orange); K2CrO4 (yellow); NiCl2 (turquoise); CuSO4 (blue); KMnO4 (purple).

Colour in transition metal compounds may be due to electronic transitions of two principal types.

* charge transfer transitions. An electron may jump from a predominantly ligand 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. The colour of chromate, dichromate and permanganate ions is due to LMCT transitions. 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 oxidised. Mercuric iodide, HgI2, is red because of a MLCT transition. As this example shows, charge transfer transitions are not restricted to transition metals.[6]

* 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 diagrams.

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−1cm−1 (where M = mol dm−3).[7] Some d-d transitions are spin forbidden. An example occurs in octahedral, high-spin complexes of manganese(II), which has a d5 configuration in which all five electron has parallel spins; the colour of such complexes is much weaker than in complexes with spin-allowed transitions. In fact many compounds of manganese(II) appear almost colourless. The spectrum of [Mn(H2O)6]2+ shows a maximum molar absorptivity of about 0.04 M−1cm−1 in the visible spectrum.
Oxidation states

A characteristic of transition metals is that they exhibit two or more oxidation states, usually differing by one. For example, compounds of vanadium are known in all oxidation states between −1, such as [V(CO)6]−, and +5, such as VO3−4.

Main group elements in groups 13 to 17 also exhibit multiple oxidation states. The "common" oxidation states of those elements differ by two. For example, compounds of gallium in oxidation states +1 and +3 exist in which there is a single gallium atom. No such compound of Ga(II) is known: any such compound would have an unpaired electron and would behave as a free radical and be destroyed rapidly. However, under the right conditions dimeric compounds such as [Ga2Cl6]2− can be made in which a Ga-Ga bond is formed from the unpaired electron on each Ga atom.[8] Thus the main difference, regarding 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 electron from titanium (+4) up to manganese (+7), but decreases in the later elements. In the second and third rows the maximum occurs with ruthenium and osmium (+8). In compounds such as [MnO4]− and OsO4 the elements achieve a stable octet by forming four covalent bonds.

The lowest oxidation states are exhibited in such compounds as Cr(CO)6 (oxidation state zero) and [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.

Transition metal compounds are paramagnetic when they have one or more unpaired d electrons.[9] 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 [FeCl4]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 that the energy needed to pair up the spins. Some compounds are diamagnetic. These include octahedral, low-spin, d6 and square-planar d8 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.
See also

* Inner transition element, a name given to any member of the f-block
* Ligand field theory a development of crystal field theory taking covalency into account
* Post-transition metal


1. ^ International Union of Pure and Applied Chemistry. "transition element". Compendium of Chemical Terminology Internet edition.
2. ^ Greenwood, Norman N.; Earnshaw, A. (1997), Chemistry of the Elements (2nd ed.), Oxford: Butterworth-Heinemann, ISBN 0080379419
3. ^ Jensen, William B. (2003). "The Place of Zinc, Cadmium, and Mercury in the Periodic Table". Journal of Chemical Education 80 (8): 952–961. doi:10.1021/ed080p952. http://www.uv.es/~borrasj/ingenieria_web/temas/tema_1/lecturas_comp/p952.pdf.
4. ^ Cotton, F. Albert; Wilkinson, G.; Murillo, C. A. (1999). Advanced Inorganic Chemistry (6th ed.). New York: Wiley.
5. ^ Matsumoto, Paul S (2005). "Trends in Ionization Energy of Transition-Metal Elements". Journal of Chemical Education 82: 1660. doi:10.1021/ed082p1660. http://www.jce.divched.org/Journal/Issues/2005/Nov/abs1660.html.
6. ^ T.M. Dunn in Lewis, J.; Wilkins,R.G. (1960). Modern Coordination Chemistry. New York: Interscience. , Chapter 4, Section 4, "Charge Transfer Spectra", pp. 268-273.
7. ^ Orgel, L.E. (1966). An Introduction to Transition-Metal Chemistry, Ligand field theory (2nd. ed.). London: Methuen.
8. ^ Greenwood, Norman N.; Earnshaw, A. (1997), Chemistry of the Elements (2nd ed.), Oxford: Butterworth-Heinemann, ISBN 0080379419 p. 240
9. ^ Figgis, B.N.; Lewis, J. (1960). Lewis, J. and Wilkins, R.G.. ed. The Magnetochemistry of Complex Compounds. Modern Coordination Chemistry. New York: Interscience. pp. 400–454.

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