Transition metal

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

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 (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.

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.[8][9] 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.[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 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.[12] This is based on the Aufbau principle (or Madelung rule) for filling electron sub-shells, in which 4f is filled before 5d (and 5f before 6d), so that the f sub-shell 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.[13] 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.[14]

Zinc, cadmium, and mercury are sometimes excluded from the transition metals,[5] as they have the electronic configuration [ ]d¹⁰s², with no incomplete d shell.[15] 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 Hg²⁺
₂, 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 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,[16] but some authors still consider this compound to be exceptional.[17] 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 (which are not definitely known for the lighter group 12 elements).

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.

The general electronic configuration of the d-block elements is (noble gas) (n − 1)d^1–10ns^0–2. Here "(noble gas)" is the 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 (n − 2)f^0–14 electrons, which are omitted from the tables below.

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)d^m. This rule is however only approximate – it only holds for some of the transition elements, and only then in the neutral ground states.

The d sub-shell is the next-to-last sub-shell and is denoted as ${\displaystyle (n-1)d}$ 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 atypical configuration [Ar]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 ${\displaystyle (n-1)d}$ 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:[18]

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.[18]

The ${\displaystyle (n-1)d}$ 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 ${\displaystyle s(ns)}$ and ${\displaystyle p(np)}$ 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.

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 color 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[19]
• 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.[20]

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-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).[21] 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.

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

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(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.[22] 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 configuration by 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.

• Inner transition element, a name given to any member of the f-block
• Main group element, an element other than a transition metal
• Ligand field theory a development of crystal field theory taking covalency into account
• Crystal field theory a model that describes the breaking of degeneracies of electronic orbital states
• Post-transition metal, a metallic element to the right of the transition metals in the periodic table