Transition metal

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Transition metals in the periodic table
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson
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In chemistry, a transition metal (or transition element) is a

inner transition metals
and are sometimes considered to be transition metals as well.

Since they are metals, they are lustrous and have good electrical and thermal conductivity. Most (with the exception of

catalysts in elemental form or in compounds such as coordination complexes and oxides. Most are strongly paramagnetic because of their unpaired d electrons, as are many of their compounds. All of the elements that are ferromagnetic near room temperature are transition metals (iron, cobalt and nickel) or inner transition metals (gadolinium
).

English chemist Charles Rugeley 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.[1][2][3] These elements are now known as the d-block.

The first row of transition metals, in order

Definition and classification

The 2011

Red Book allows for the group 12 elements to be excluded, but not the 2011 Principles.[7]

The IUPAC

cations with an incomplete d sub-shell", but this definition is taken from an old edition of the Red Book and is no longer present in the current edition.[7]

In the d-block, the atoms of the elements have between zero and ten d electrons.

Transition metals in the d-block
Group 3 4 5 6 7 8 9 10 11 12
Period 4 21Sc 22Ti 23V 24Cr 25Mn 26Fe 27Co 28Ni 29Cu 30Zn
5 39Y 40Zr 41Nb 42Mo 43Tc 44Ru 45Rh 46Pd 47Ag 48Cd
6 71Lu 72Hf 73Ta 74W 75Re 76Os 77Ir 78Pt 79Au 80Hg
7 103Lr 104Rf 105Db 106Sg 107Bh 108Hs 109Mt 110Ds 111Rg 112Cn

Published texts and periodic tables show

IUPAC report on physical, chemical, and electronic grounds,[20] and again by a 2021 IUPAC preliminary report as it is the only form that allows simultaneous (1) preservation of the sequence of increasing atomic numbers, (2) a 14-element-wide f-block, and (3) avoidance of the split in the d-block.[15] Argumentation can still be found in the contemporary literature purporting to defend the form with lanthanum and actinium in group 3, but many authors consider it to be logically inconsistent (a particular point of contention being the differing treatment of actinium and thorium, which both can use 5f as a valence orbital but have no 5f occupancy as single atoms);[14][21][22] the majority of investigators considering the problem agree with the updated form with lutetium and lawrencium.[14]

The group 12 elements

bonding even though they are not known in oxidation states that would formally require breaking open the d-subshell, which sets them apart from the p-block elements.[24][25][26]

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,[27] but some authors still consider this compound to be exceptional.[28] Copernicium is expected to be able to use its d electrons for chemistry as its 6d subshell is destabilised by strong relativistic effects due to its very high atomic number, and as such is expected to have transition-metal-like behaviour and show higher oxidation states than +2 (which are not definitely known for the lighter group 12 elements). Even in bare dications, Cn2+ is predicted to be 6d87s2, unlike Hg2+ which is 5d106s0.

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. Whether copernicium behaves more like mercury or has properties more similar to those of the noble gas radon is not clear. Relative inertness of Cn would come from the relativistically expanded 7s–7p1/2 energy gap, which is already adumbrated in the 6s–6p1/2 gap for Hg, weakening metallic bonding and causing its well-known low melting and boiling points.

Transition metals with lower or higher group numbers are described as 'earlier' or 'later', respectively. When described in a two-way classification scheme, early transition metals are on the left side of the d-block from group 3 to group 7. Late transition metals are on the right side of the d-block, from group 8 to 11 (or 12, if they are counted as transition metals). In an alternative three-way scheme, groups 3, 4, and 5 are classified as early transition metals, 6, 7, and 8 are classified as middle transition metals, and 9, 10, and 11 (and sometimes group 12) are classified as late transition metals.

The heavy group 2 elements calcium, strontium, and barium do not have filled d-orbitals as single atoms, but are known to have d-orbital bonding participation in some compounds, and for that reason have been called "honorary" transition metals.[29] Probably the same is true of radium.[30]

The f-block elements La–Yb and Ac–No have chemical activity of the (n−1)d shell, but importantly also have chemical activity of the (n−2)f shell that is absent in d-block elements. Hence they are often treated separately as inner transition elements.

Electronic configuration

Main article: Electron configuration

The general electronic configuration of the d-block atoms is [noble gas](n − 1)d0–10ns0–2np0–1. Here "[noble gas]" is the electronic configuration of the last noble gas preceding the atom in question, and n is the highest principal quantum number of an occupied orbital in that atom. For example, Ti (Z = 22) is in period 4 so that n = 4, the first 18 electrons have the same configuration of Ar at the end of period 3, and the overall configuration is [Ar]3d24s2. The period 6 and 7 transition metals also add core (n − 2)f14 electrons, which are omitted from the tables below. The p orbitals are almost never filled in free atoms (the one exception being lawrencium due to relativistic effects that become important at such high Z), but they can contribute to the chemical bonding in transition metal compounds.

The Madelung rule predicts that the inner d orbital is filled after the valence-shell s orbital. The typical electronic structure of transition metal atoms is then written as [noble gas]ns2(n − 1)dm. This rule is approximate, but holds for most of the transition metals. Even when it fails for the neutral ground state, it accurately describes a low-lying excited state.

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

The elements in group 3 have an ns2(n − 1)d1 configuration, except for lawrencium (Lr): its 7s27p1 configuration exceptionally does not fill the 6d orbitals at all. 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 subshell till it is complete. 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:[31]

First (3d) d-block Series (Sc–Zn)
Group 3 4 5 6 7 8 9 10 11 12
Atomic number 21 22 23 24 25 26 27 28 29 30
Element Sc Ti V Cr Mn Fe Co Ni Cu Zn
Electron
configuration 		3d14s2 3d24s2 3d34s2 3d54s1 3d54s2 3d64s2 3d74s2 3d84s2 3d104s1 3d104s2
Second (4d) d-block Series (Y–Cd)
Atomic number 39 40 41 42 43 44 45 46 47 48
Element Y Zr Nb Mo Tc Ru Rh Pd Ag Cd
Electron
configuration 		4d15s2 4d25s2 4d45s1 4d55s1 4d55s2 4d75s1 4d85s1 4d105s0 4d105s1 4d105s2
Third (5d) d-block Series (Lu–Hg)
Atomic number 71 72 73 74 75 76 77 78 79 80
Element Lu Hf Ta W Re Os Ir Pt Au Hg
Electron
configuration 		5d16s2 5d26s2 5d36s2 5d46s2 5d56s2 5d66s2 5d76s2 5d96s1 5d106s1 5d106s2
Fourth (6d) d-block Series (Lr–Cn)
(Configurations predicted for Mt–Cn)
Atomic number 103 104 105 106 107 108 109 110 111 112
Element Lr Rf Db Sg Bh Hs Mt Ds Rg Cn
Electron
configuration 		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 to the

exchange energy.[31] The exceptions are in any case not very relevant for chemistry because the energy difference between them and the expected configuration is always quite low.[32]

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 coloured compounds etc. The valence s and p orbitals (ns and np) 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 are 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

Most transition metals can be bound to a variety of

ligands, allowing for a wide variety of transition metal complexes.[34]

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.

  • 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
    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

spin forbidden. An example occurs in octahedral, high-spin complexes of manganese
(II), which has a d5 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 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.

Oxidation states of the transition metals. The solid dots show common oxidation states, and the hollow dots show possible but unlikely states.

free radical and generally be destroyed rapidly, but some stable radicals of Ga(II) are known.[36] Gallium also has a formal oxidation state of +2 in dimeric compounds, such as [Ga
2Cl
6]2−
, which contain a Ga-Ga bond formed from the unpaired electron on each Ga atom.[37]
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

covalent bonding
.

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

Main article: Magnetochemistry

Transition metal compounds are

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 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
.

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)).[39] Once a little Mn2+ has been produced, it can react with MnO4− forming Mn3+. This then reacts with C2O4− ions forming Mn2+ again.

Physical properties

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

In general, transition metals possess a high

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, which again tends to differentiate them from the accepted transition metals. Mercury has a melting point of −38.83 °C (−37.89 °F) and is a liquid at room temperature.

See also

Scholia has a profile for transition metal (Q19588).

References

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