Transition metal

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Transition metals in the periodic table

In chemistry, the term transition metal (or transition element) has two 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".

• Most 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".

Jensen^[4] reviews the history of the terms "transition element" (or "metal") and "d-block". The word transition was first used to describe the elements now known as the d-block by the English chemist Charles Bury in 1921, who 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]

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
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	57–71	Hf 72	Ta 73	W 74	Re 75	Os 76	Ir 77	Pt 78	Au 79	Hg 80
Period 7	89–103	Rf 104	Db 105	Sg 106	Bh 107	Hs 108	Mt 109	Ds 110	Rg 111	Cn 112

The typical electronic structure of transition metal atoms can be written as [ ]ns²(n-1)d^m, following the Madelung rule where the inner d orbital is predicted to be filled after the valence-shell s orbital. This is actually not the case; the 4s electrons are higher in energy than the 3d as shown spectroscopically. An ion such as Fe2+ has no 4s electrons: it has the electronic configuration [Ar]3d⁶ as compared with the configuration of the atom, [Ar]4s²3d⁶.

The elements of groups 3–12 are now generally recognized as transition metals, although the elements La-Lu and Ac-Lr and Group 12 attract different definitions from different authors.

1. 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.^[6][7] The two series together are classified as f-block elements, or (in older sources) as “inner transition elements”.
2. Some inorganic chemistry textbooks include La with the lanthanides and Ac with the actinides.^[8][9][10] 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.
3. A third classification defines the f-block elements as La-Yb and Ac-No, while placing Lu and Lr in Group 3.^[4] 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. Eric Scerri has proposed placing Lu and Lr in group 3 on the grounds of continuous sequences of atomic numbers in an expanded or long-form periodic table.^[11]

Zinc, cadmium, and mercury are sometimes excluded from the transition metals^[4] as they have the electronic configuration [ ]d¹⁰s², with no incomplete d shell.^[12] 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 be classed as post-transition metals in this case, because of the formation of a covalent bond between the two atoms of the dimer. 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 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 metal,^[13] but some authors still consider this compound to be exceptional.^[14]

Position in the Periodic Table

The d-block as stated earlier, is present in the centre of the long form of periodic table. These are flanked or surrounded by elements belonging to s and p-blocks on both sides. These are called transition elements since they represent a transition i.e., there is a change from metallic character of s-block elements to non-metallic character of p-block elements through d-block elements which are also metals. As pointed above there are four transition series in this block. Since the filling of electrons takes place in (n-1)d orbitals, the periods to which these series belong, is actually one more than the actual series. For example, the elements included in 3d series belong to fourth period ; the elements included in 4d series belong to the fifth period and so on.

Electronic configuration

The general electronic configuration of the d-block elements is [Inert gas] (n-1)d^1-10n s^1-2 The d-sub-shell is the penultimate (last but one) sub-shell and is denoted as (n-1) d-sub-shell. The number of s electrons may vary from one to two. The s-sub-shell in the valence shell is represented as the ns sub-shell. However, palladium (Pd) is an exception with no electron in the s-sub shell. In the periodic table, the transition metals are present in ten groups (3 to 12). Group-2 belongs to the s- block with an ns² configuration.

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 which has configuration [Ar]4s². The electronic configuration of scandium (Sc), the first element of group-3 with atomic number Z=21 is[Ar]4s²3d¹. As we move from left to right, electrons are added to the same d-sub-shell till it is complete. The element of group-12 in the first transition series is zinc (Zn) with configuration [Ar]4s²3d¹⁰. 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 transition series are given below:

First (3d) Transition 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) Transition 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	4d105s0	4d105s1	4d105s2

Third (5d) Transition Series (Lu-Hg)^[15]

At.No	71	72	73	74	75	76	77	78	79	80
Element	Lu	Hf	Ta	W	Re	Os	Ir	Pt	Au	Hg
Config.	5d16s2	5d26s2	5d36s2	5d46s2	5d56s2	5d66s2	5d76s2	5d96s1	5d106s1	5d106s2

Fourth (6d) Transition Series (Lr-Cn)

At. No.	103	104	105	106	107	108	109	110	111	112
Element	Lr	Rf	Db	Sg	Bh	Hs	Mt	Ds	Rg	Cn
Config.	7s27p1	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 shown by Pt, Au and Hg.. 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 reactivity of unpaired d electrons.^[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)

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.

• 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).^[17] 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 17 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.^[18] 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 and third rows the maximum occurs with ruthenium and osmium (+8). 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.

Magnetism

Transition metal compounds are paramagnetic when they have one or more unpaired d electrons.^[19] 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 than 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. Anti-ferromagnetism 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 catalytic hydrogenation) are some of the examples. Catalysts at a solid surface (nanomaterial-based catalysts) 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 catalysts.

Other properties

As implied by the name, all transition metals are metals and conductors of electricity.

In general, transition metals possess a high density and high melting points and boiling points. These properties are due to metallic bonding by delocalized d electrons, leading to 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. Mercury has a melting point of −38.83 °C (−37.89 °F) and is a liquid at room temperature.

Many transition metals can be bound to a variety of ligands.^[20]