|Transition metals in the periodic table|
In chemistry, the term transition metal (or transition element) has two possible meanings:
Jensen 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.
In the d-block the atoms of the elements have between 1 and 10 d electrons.
|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|
With a few minor exceptions, the electronic structure of transition metal atoms can be written as [ ]ns2(n-1)dm, where the inner d orbital has more energy than the valence-shell s orbital. In divalent and trivalent ions of the transition metals, the situation is reversed such that the s electrons have higher energy. Consequently, an ion such as Fe2+ has no s 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.
Zinc, cadmium, and mercury are sometimes incorrectly classified as non-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. However, these elements can exist in many 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.
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
Colour in transition-series metal compounds is generally due to electronic transitions of two principal types.
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 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). 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. In fact many compounds of manganese(II) appear almost colourless. The spectrum of [Mn(H2O)6]2+ shows a maximum molar absorptivity of about 0.04 M−1cm−1 in the visible spectrum.
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−
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 [Ga2Cl6]2−, 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 (+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 [MnO4]− and OsO4 the elements achieve a stable octet by forming four covalent bonds.
The lowest oxidation states are exhibited in such compounds 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.
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 [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.
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.
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. In fact mercury has a melting point of −38.83 °C (−37.89 °F) and is a liquid at room temperature.
Here you can share your comments or contribute with more information, content, resources or links about this topic.