Transition metal


In chemistry, the term transition metal has three possible definitions:
English chemist Charles 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.

Classification

In the d-block, the atoms of the elements have between one and ten d electrons.
Group3456789101112
Period 421Sc22Ti23V24Cr25Mn26Fe27Co28Ni29Cu30Zn
539Y40Zr41Nb42Mo43Tc44Ru45Rh46Pd47Ag48Cd
657La72Hf73Ta74W75Re76Os77Ir78Pt79Au80Hg
789Ac104Rf105Db106Sg107Bh108Hs109Mt110Ds111Rg112Cn

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, colored complexes, and catalytic properties either as the element or as ions. 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.
  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 and Th–Lr as the "actinide" series. The two series together are classified as f-block elements, or as "inner transition elements".
  2. Some inorganic chemistry textbooks include La with the lanthanides and Ac with the actinides. 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 sub-shell, 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. This is based on the Aufbau principle for filling electron sub-shells, in which 4f is filled before 5d, so that the f sub-shell is actually full at Yb, while Lu has an s2f14d1 configuration. However La and Ac are exceptions to the Aufbau principle with electron configuration s2d1, so it is not clear from atomic electron configurations whether La or Lu should be considered as transition metals. This can be resolved by considering that what are excited states for the free atom and ion can become the ground state in chemical environments; La and Ac have vacant low-lying f sub-shells which are filled in Lu and Lr, so excitation to f orbitals is possible in La and Ac but not in Lu or Lr.
Zinc, cadmium, and mercury are generally excluded from the transition metals, as they have the electronic configuration d10s2, with no incomplete d shell. In the oxidation state +2, the ions have the electronic configuration …d10. Although these elements can exist in other oxidation states, including the +1 oxidation state, as in the diatomic ion, they still have a complete d shell in these 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.
The recent 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 sub-shell is destabilised by strong relativistic effects due to its very high atomic number, and as such is expected to have transition-metal-like behaviour when it shows higher oxidation states than +2.
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.

Subclasses

Early transition metals are on the left side of the periodic table from group 3 to group 7. Late transition metals are on the right side of the d-block, from group 8 to 11.

Electronic configuration

The general electronic configuration of the d-block elements is d1–10n s0–2. The period 6 and 7 transition metals also add f0–14 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 ns2dm where the inner d orbital is predicted to be filled after the valence-shell's s orbital is filled. 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 sub-shell. The number of s electrons in the outermost s sub-shell is generally one or two except palladium, 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, with some authors including some elements in groups 3 or 12.
The elements in group 3 have an ns2d1 configuration. 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 sub-shell till it is complete. The element of group 11 in the first transition series is copper with an atypical configuration 4s13d10. Despite the filled d sub-shell in metallic copper it nevertheless forms a stable ion with an incomplete d sub-shell. Since the electrons added fill the 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 nr21222324252627282930
ElementScTiVCrMnFeCoNiCuZn
Electron
configuration		3d14s23d24s23d34s23d54s13d54s23d64s23d74s23d84s23d104s13d104s2

Atomic nr39404142434445464748
ElementYZrNbMoTcRuRhPdAgCd
Electron
configuration		4d15s24d25s24d45s14d55s14d55s24d75s14d85s14d104d105s14d105s2

Atomic nr57727374757677787980
ElementLaHfTaWReOsIrPtAuHg
Electron
configuration		5d16s25d26s25d36s25d46s25d56s25d66s25d76s25d96s15d106s15d106s2

Atomic nr89104105106107108109110111112
ElementAcRfDbSgBhHsMtDsRgCn
Electron
configuration		6d17s26d27s26d37s26d47s26d57s26d67s26d77s26d87s26d97s26d107s2

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 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 and 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
Most transition metals can be bound to a variety of ligands, allowing for a wide variety of transition metal complexes.

Coloured compounds

Colour in transition-series metal compounds is generally due to electronic transitions of two principal types.
A metal-to-ligand charge transfer 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.
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 of bands caused by d-d transitions are relatively low, roughly in the range 5-500 M−1cm−1. Some d-d transitions are spin forbidden. An example occurs in octahedral, high-spin complexes of manganese,
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 appear almost colourless. The spectrum of 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, and +5, such as.
Main group elements in groups 13 to 18 also exhibit multiple oxidation states. The "common" oxidation states of these elements typically differ by two instead of one. For example, compounds of gallium in oxidation states +1 and +3 exist in which there is a single gallium atom. No compound of Ga 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, which contain a Ga-Ga bond formed from the unpaired electron on each Ga atom. 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 up to manganese, but decreases in the later elements. In the second row, the maximum occurs with ruthenium, and in the third row, the maximum occurs with iridium. In compounds such as and, the elements achieve a stable configuration by covalent bonding.
The lowest oxidation states are exhibited in metal carbonyl complexes such as and 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 six water molecules arranged octahedrally.

Magnetism

Transition metal compounds are paramagnetic when they have one or more unpaired d electrons. 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 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 oxide, finely divided iron, and nickel are some of the examples. Catalysts at a solid surface involve the formation of bonds between reactant molecules and atoms of the surface of the catalyst. 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. 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. One example is the reaction of oxalic acid with acidified potassium permanganate. 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 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 sub-shells prevent d–d bonding, which again tends to differentiate them from the accepted transition metals. Mercury has a melting point of and is a liquid at room temperature.