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In chemistry, the term transition metal (or transition element) has three possible meanings:

The IUPAC definition^[1] defines a transition metal as “an element whose atom has a partially filled d sub-shell, or which can give rise to cations with an incomplete d sub-shell”.
Many scientists describe a “transition metal” as any element in the d-block of the periodic table, which includes groups 3 to 12 on the periodic table.^[2]^[3] In actual practice, the f-block lanthanide and actinide series are also considered transition metals and are called “inner transition metals”.
Cotton and Wilkinson^[4] expand the brief IUPAC definition (see above) by specifying which elements are included. As well as the elements of groups 4 to 11, they add scandium and yttrium in group 3 which have a partially filled d subshell in the metallic state. These last two elements are included even though they do not (so far) seem to possess the catalytic properties which are so characteristic of the transition metals in general. Lanthanum and actinium in Group 3 are however classified as lanthanides and actinides respectively.

English chemist Charles Bury (1890-1968) first used the word transition in this context in 1921, when he referred to a transition series of elements during the change of an inner layer of electrons (for example n=3 in the 4th row of the periodic table) from a stable group of 8 to one of 18, or from 18 to 32.^[5]^[6]^[7] These elements are now known as the d-block.

Classification

In the d-block the atoms of the elements have between 1 and 10 d electrons.

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

The elements of groups 4-11 are generally recognized as transition metals, justified by their typical chemistry, i.e. a large range of complex ions in various oxidation states, coloured complexes, and catalytic properties either as the element or as ions (or both). Sc and Y in Group 3 are also generally recognized as transition metals. However, the elements La–Lu and Ac–Lr and Group 12 attract different definitions from different authors.

Many chemistry textbooks and printed periodic tables classify La and Ac as Group 3 elements and transition metals, since their atomic ground-state configurations are s²d¹ like Sc and Y. The elements Ce-Lu are considered as the “lanthanide” series (or “lanthanoid” according to IUPAC) and Th-Lr as the “actinide” series.^[8]^[9] The two series together are classified as f-block elements, or (in older sources) as “inner transition elements”.
Some inorganic chemistry textbooks include La with the lanthanides and Ac with the actinides.^[4]^[10]^[11] This classification is based on similarities in chemical behaviour, and defines 15 elements in each of the two series even though they correspond to the filling of an f subshell which can only contain 14 electrons.
A third classification defines the f-block elements as La-Yb and Ac-No, while placing Lu and Lr in Group 3.^[5] This is based on the Aufbau principle (or Madelung rule) for filling electron subshells, in which 4f is filled before 5d (and 5f before 6d), so that the f subshell is actually full at Yb (and No) while Lu (and Lr) has an [ ]s²f¹⁴d¹ configuration. However La and Ac are exceptions to the Aufbau principle with electron configuration [ ]s²d¹ (not [ ]s²f¹ as the Aufbau principle predicts) so it is not clear from atomic electron configurations whether La or Lu (Ac or Lr) should be considered as transition metals.^[12]

Zinc, cadmium, and mercury are generally excluded from the transition metals^[5] as they have the electronic configuration [ ]d¹⁰s², with no incomplete d shell.^[13] In the oxidation state +2 the ions have the electronic configuration [ ] d¹⁰. However, these elements can exist in other oxidation states, including the +1 oxidation state, as in the diatomic ion Hg2+
2. The group 12 elements Zn, Cd and Hg may therefore, under certain criteria, be classed as post-transition metals in this case. 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 also include the 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.

The recent (though disputed and so far not reproduced independently) synthesis of mercury(IV) fluoride (HgF
4) has been taken by some to reinforce the view that the group 12 elements should be considered transition metals,^[14] but some authors still consider this compound to be exceptional.^[15]

Although meitnerium, darmstadtium, and roentgenium are within the d-block and are expected to behave as transition metals analogous to their lighter congeners iridium, platinum, and gold, this has not yet been experimentally confirmed.

Electronic configuration

The general electronic configuration of the d-block elements is [Inert gas] (n − 1)d^1–10n s^0–2. The period 6 and 7 transition metals also add (n − 2)f^0–14 electrons, which are omitted from the tables below.

The Madelung rule predicts that the typical electronic structure of transition metal atoms can be written as [inert gas]ns²(n − 1)d^m where the inner d orbital is predicted to be filled after the valence-shell s orbital. This rule is however only approximate – it only holds for some of the transition elements, and only then in their neutral ground state.

The d-sub-shell is the next-to-last sub-shell and is denoted as $(n-1)d$ -sub-shell. The number of s electrons in the outermost s sub-shell is generally one or two except palladium (Pd), with no electron in that s-sub shell in its ground state. The s-sub-shell in the valence shell is represented as the ns sub-shell, e.g. 4s. In the periodic table, the transition metals are present in eight groups (4 to 11), with some authors including some elements in groups 3 or 12.

The elements in group 3 have an ns²(n − 1)d¹ configuration. The first transition series is present in the 4th period, and starts after Ca (Z = 20) of group-2 with the configuration [Ar]4s², or scandium (Sc), the first element of group 3 with atomic number Z = 21 and configuration [Ar]4s²3d¹, depending on the definition used. As we move from left to right, electrons are added to the same d-sub-shell till it is complete. The element of group 11 in the first transition series is copper (Cu) with an untypical configuration [Ar]4s¹3d¹⁰. Despite the filled d subshell in metallic copper it nevertheless forms a stable ion with an incomplete d subshell. Since the electrons added fill the $(n-1)d$ orbitals, the properties of the d-block elements are quite different from those of s and p block elements in which the filling occurs either in s or in p-orbitals of the valence shell. The electronic configuration of the individual elements present in all the d-block series are given below:

First (3d) d-block Series (Sc–Zn)
Group	3	4	5	6	7	8	9	10	11	12
At.no.	21	22	23	24	25	26	27	28	29	30
Element	Sc	Ti	V	Cr	Mn	Fe	Co	Ni	Cu	Zn
Config.	3d¹4s²	3d²4s²	3d³4s²	3d⁵4s¹	3d⁵4s²	3d⁶4s²	3d⁷4s²	3d⁸4s²	3d¹⁰4s¹	3d¹⁰4s²

Second (4d) d-block Series (Y–Cd)
At. No.	39	40	41	42	43	44	45	46	47	48
Element	Y	Zr	Nb	Mo	Tc	Ru	Rh	Pd	Ag	Cd
Config.	4d¹5s²	4d²5s²	4d⁴5s¹	4d⁵5s¹	4d⁵5s²	4d⁷5s¹	4d⁸5s¹	4d¹⁰	4d¹⁰5s¹	4d¹⁰5s²

Third (5d) d-block Series (La–Hg)
At.No	57	72	73	74	75	76	77	78	79	80
Element	La	Hf	Ta	W	Re	Os	Ir	Pt	Au	Hg
Config.	5d¹6s²	5d²6s²	5d³6s²	5d⁴6s²	5d⁵6s²	5d⁶6s²	5d⁷6s²	5d⁹6s¹	5d¹⁰6s¹	5d¹⁰6s²

Fourth (6d) d-block Series (Ac–Cn)
At. No.	89	104	105	106	107	108	109	110	111	112
Element	Ac	Rf	Db	Sg	Bh	Hs	Mt	Ds	Rg	Cn
Config.	6d¹7s²	6d²7s²	6d³7s²	6d⁴7s²	6d⁵7s²	6d⁶7s²	6d⁷7s²	6d⁸7s²	6d⁹7s²	6d¹⁰7s²

A careful look at the electronic configuration of the elements reveals that there are certain exceptions, for example Cr and Cu. These are either because of the symmetry or nuclear-electron and electron-electron force.

The $(n-1)d$ orbitals that are involved in the transition metals are very significant because they influence such properties as magnetic character, variable oxidation states, formation of colored compounds etc. The valence $s(ns)$ and $p(np)$ orbitals have very little contribution in this regard since they hardly change in the moving from left to the right in a transition series. In transition metals, there is a greater horizontal similarities in the properties of the elements in a period in comparison to the periods in which the d-orbitals are not involved. This is because in a transition series, the valence shell electronic configuration of the elements do not change. However, there are some group similarities as well.

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^[16]
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.^[17]

Coloured compounds

From left to right, aqueous solutions of: Co(NO
3)
2 (red); K
2Cr
2O
7 (orange); K
2CrO
4(yellow); NiCl
2 (turquoise); CuSO
4 (blue); KMnO
4 (purple).

Colour in transition-series metal compounds is generally 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. For example, the colour of chromate, dichromate and permanganate ions is due to LMCT transitions. Another example is that mercuric iodide, HgI₂, is red because of a LMCT transition.

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 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-dtransition. 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).^[18] 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. Many compounds of manganese(II) appear almost colourless. The spectrum of [Mn(H
2O)
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 18 also exhibit multiple oxidation states. The “common” oxidation states of these elements typically differ by two. For example, compounds of gallium in oxidation states +1 and +3 exist in which there is a single gallium atom. No 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. The only compounds in which gallium has a formal oxidation state of +2 are dimeric compounds, such as [Ga
2Cl
6]2−
, which contain a Ga-Ga bond formed from the unpaired electron on each Ga atom.^[19] 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 [MnO
4]−
and OsO
4 the elements achieve a stable octet by forming four covalent bonds.

The lowest oxidation states are exhibited in metal carbonyl complexes such 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.