Transition metal


In chemistry, a transition metal is a chemical element in the d-block of the periodic table, though the elements of group 12 are sometimes excluded. The lanthanide and actinide elements are called inner transition metals and are sometimes considered to be transition metals as well.
They are lustrous metals with good electrical and thermal conductivity. Most are hard and strong, and have high melting and boiling temperatures. They form compounds in any of two or more different oxidation states and bind to a variety of ligands to form coordination complexes that are often coloured. They form many useful alloys and are often employed as 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 or inner transition metals.
English chemist Charles Rugeley Bury 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 from a stable group of 8 to one of 18, or from 18 to 32. These elements are now known as the d-block.

Definition and classification

The 2011 IUPAC Principles of Chemical Nomenclature describe a "transition metal" as any element in groups 3 to 12 on the periodic table. This corresponds exactly to the d-block elements, and many scientists use this definition. In actual practice, the f-block lanthanide and actinide series are called "inner transition metals". The 2005 Red Book allows for the group 12 elements to be excluded, but not the 2011 Principles.
The IUPAC Gold Book 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", but this definition is taken from an old edition of the Red Book and is no longer present in the current edition.
In the d-block, the atoms of the elements have between zero and ten d electrons.
Group3456789101112
Period 421Sc22Ti23V24Cr25Mn26Fe27Co28Ni29Cu30Zn
539Y40Zr41Nb42Mo43Tc44Ru45Rh46Pd47Ag48Cd
671Lu72Hf73Ta74W75Re76Os77Ir78Pt79Au80Hg
7103Lr104Rf105Db106Sg107Bh108Hs109Mt110Ds111Rg112Cn

Published texts and periodic tables show variation regarding the heavier members of group 3. The common placement of lanthanum and actinium in these positions is not supported by physical, chemical, and electronic evidence, which overwhelmingly favour putting lutetium and lawrencium in those places. Some authors prefer to leave the spaces below yttrium blank as a third option, but there is confusion on whether this format implies that group 3 contains only scandium and yttrium, or if it also contains all the lanthanides and actinides; additionally, it creates a 15-element-wide f-block, when quantum mechanics dictates that the f-block should only be 14 elements wide. The form with lutetium and lawrencium in group 3 is supported by a 1988 IUPAC report on physical, chemical, and electronic grounds, and again by a 2021 IUPAC preliminary report as it is the only form that allows simultaneous preservation of the sequence of increasing atomic numbers, a 14-element-wide f-block, and avoidance of the split in the d-block. 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 ; the majority of investigators considering the problem agree with the updated form with lutetium and lawrencium.
The group 12 elements zinc, cadmium, and mercury are sometimes excluded from the transition metals. This is because they have the electronic configuration d10s2, where the d shell is complete, and they still have a complete d shell in all their known oxidation states. 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 and 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. Moreover, Zn, Cd, and Hg can use their d orbitals for 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.
The 2007 synthesis of mercury fluoride has been taken by some to reinforce the view that the group 12 elements should be considered transition metals, but some authors still consider this compound to be exceptional. 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. 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. 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 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. The same is likely true of radium.
The f-block elements La–Yb and Ac–No have chemical activity of the d shell, but importantly also have chemical activity of the f shell that is absent in d-block elements. Hence they are often treated separately as inner transition elements.

Electronic configuration

The general electronic configuration of the d-block atoms is d0–10ns0–2np0–1. Here "" 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 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 3d24s2. The period 6 and 7 transition metals also add core f14 electrons, which are omitted from the tables below. The p orbitals are almost never filled in free atoms, 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 ns2dm. 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 d subshell. The number of s electrons in the outermost s subshell is generally one or two except palladium, 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.
The elements in group 3 have an ns2d1 configuration, except for lawrencium : 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 of group 2 with the configuration 4s2, or scandium, the first element of group 3 with atomic number Z = 21 and configuration 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 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:
Group3456789101112
Atomic number21222324252627282930
ElementScTiVCrMnFeCoNiCuZn
Electron
configuration		3d14s23d24s23d34s23d54s13d54s23d64s23d74s23d84s23d104s13d104s2

Atomic number39404142434445464748
ElementYZrNbMoTcRuRhPdAgCd
Electron
configuration		4d15s24d25s24d45s14d55s14d55s24d75s14d85s14d105s04d105s14d105s2

Atomic number71727374757677787980
ElementLuHfTaWReOsIrPtAuHg
Electron
configuration		5d16s25d26s25d36s25d46s25d56s25d66s25d76s25d96s15d106s15d106s2

Atomic number103104105106107108109110111112
ElementLrRfDbSgBhHsMtDsRgCn
Electron
configuration		7s27p16d27s26d37s26d47s26d57s26d67s26d77s26d87s26d97s26d107s2

A careful look at the electronic configuration of the elements reveals that there are certain exceptions to the Madelung rule. For Cr as an example the rule predicts the configuration 3d44s2, but the observed atomic spectra show that the real ground state is 3d54s1. To explain such exceptions, it is necessary to consider the effects of increasing nuclear charge on the orbital energies, as well as the electron–electron interactions including both Coulomb repulsion and exchange energy. 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.
The 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 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.