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Transition metal

The thirty chemical elements 21 through 30, 39 through 48, and 71 through 80, are commonly referred to as the transition metals. This name comes from their position in the periodic table of elements, which represent the successive addition of electrons to the d atomic orbitals of the atoms as one progresses through each of the three periods. Transition elements are chemically defined as elements which form at least one ion with a partially filled subshell of d electrons.

Group   Period 4 Period 5 Period 6 Period 7
3 (III B)   Sc 21 Y 39 Lu 71 Lr 103
4 (IV B)   Ti 22 Zr 40 Hf 72 Unq 104
5 (V B)   V 23 Ta 73 Nb 41 Unp 105
6 (VI B)   Cr 24 Mo 42 W 74 Unh 106
7 (VII B)   Mn 25 Tc 43 Re 75 Uns 107
8 (VIII B)   Fe 26 Ru 44 Os 76 Uno 108
9 (VIII B)   Co 27 Rh 45 Ir 77 Une 109
10 (VIII B)   Ni 28 Pd 46 Pt 78  
11 (I B)   Cu 29 Ag 47 Au 79  
12 (II B)   Zn 30 Cd 48 Hg 80

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Electronic configuration Main group elements prior to the appearance of the transition group elements in the periodic chart (ie, elements number 1 through 20) have no electrons in d orbitals, but only in the s and p orbitals. (Though the low-lying, but empty d orbitals are thought to play a role in their d period elements such as silicon, phosphorus and sulfur)

From Scandium to Zinc, d block elements fill up their d orbitals across the period. With the exception of copper and chromium, all d block elements have two electrons in their outer s orbital, even elements with incomplete 3d orbitals. This is unusual: lower orbitals are usually filled up before outer shells. It happens that the s orbitals in d block elements are at lower energy states than the d subshells. As atoms always strive to be in states of lowest energy, s shells are filled up first. The copper and chromium exceptions - which have one electron in their outer orbital - do so because of electron repulsion. Sharing the electrons throughout the s and d orbitals gives lower energy states to the atoms than putting two electrons in the outer s orbital.

Not all d block elements are transition metals. Scandium and zinc don't qualify, due to the chemical definition given above. Scandium has one electron in its d subshell, and 2 electrons in its outer s orbital. As scandium's only ion (Sc3+) has no electrons in its d orbital it is clear that it doesn't have a 'partially filled d orbital'. Similarly, zinc is not applicable because its only ion, Zn2+, has a full d orbital.

Chemical properties Transition elements tend to have high tensile strength, density and melting and boiling points. As with many properties of transition metals, this is due to d orbital electrons' ability to delocalise within the metal lattice. In metallic substances, the more electrons shared between nuclei, the stronger the metal.

There are four common characteristic properties of transition elements:

Variable oxidation states Compared to Group II elements such as calcium, transition elements form ions with a wide variety of oxidation states. Calcium ions typically don't lose more than two electrons, whereas transition metals can lose up to nine. The reason for this can be obtained by studying the ionisation enthalpies[?] of both groups. The energies required to remove electrons from calcium are low until you try to remove electrons from below its outer two s orbitals. In fact Ca3+ has an ionisation enthalpy so high that it rarely occurs naturally. However a transition element like vanadium has roughly linear increasing ionisation enthalpies throughout its s and d orbitals, due to the close energy difference between the 3d and 4s orbitals. Transition metal ions are therefore commonly found in very high states.

Certain patterns can be seen to emerge across the period of transition elements:

  • The number of oxidation states of each ion increases up to Mn, after which they start to drop. This drop is due to the stronger pull from the protons in the nucleus towards the electrons, making them harder to remove.
  • When the elements are in lower oxidation states, they can be found as simple ions. However elements in higher oxidation states are usually bonded covalently to electronegative compounds such as O or F, often in an anion.

Properties with respect to the stability of oxidation states:

  • Higher oxidation state ions become less stable across the period.
  • Ions in higher oxidation states tend to make good oxidising agents, whereas elements in low oxidation states become reducing agents.
  • The 2+ ions across the period start as strong reducing agents, and become more stable.
  • The 3+ ions start stable and become more oxidising[?] across the period.

Catalytic activity Transition metals form good homogeneous or heterogeneous[?] catalysts, for example iron is the catalyst for the Haber process. Nickel or platinum is used in the hydrogenation of alkenes.

Coloured compounds We observe colour as varying frequencies of electromagnetic radiation in the visible region of the electromagnetic spectrum. Different colours result from the changed composition of light after it has been reflected, transmitted or absorbed after hitting a substance. Because of their structure, transition metals form many different coloured ions and complexes. Colour even varies between the different ions of a single element - MnO4- (Mn in oxidation state 7+) is a purple compound, whereas Mn2+ is pale-pink.

Complex formation can play a part in determining colour in a transition compound. This is because of the effect that ligands have on the 3d subshell. Ligands pull on some of the 3d electrons and split them in to higher and lower (in terms of energy) groups. Electromagnetic radiation is only absorbed if its frequency is proportional to the difference in energies between two energy states present in an atom (through the formula e=hf.) When light hits an atom which has had its 3d orbitals split, some of the electrons become promoted to the higher group. Compared to an un-complexed ion, different frequencies can be absorbed, hence different colours are observed.

The colour of a complex depends on:

  • The nature of the metal ion, specifically the number of electrons in the d orbitals
  • The arrangement of the ligands around the metal ion (for example geometric isomers[?] can display different colours)
  • The nature of the ligands surrounding the metal ion. The stronger the ligands then the greater the energy difference between the split high and low 3d groups.

The complex formed by the d block element zinc (though not strictly a transition element) is colourless, because the 3d orbitals are full - no electrons are able to move up to the higher group.

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