This section explains what a transition metal is in terms of its electronic structure, and then goes on to look at the general features of transition metal chemistry that include the following:

· Variable oxidation state (oxidation number)
· Complex ion formation
· Coloured ions
· Catalytic activity

The electronic structures of transition metals

What is a transition metal?
The terms transition metal (or element) and d block element are sometimes used as if they mean the same thing. They don't - there's a subtle difference between the two terms. We'll explore d block elements first:

d block elements

You will remember that when you are building the Periodic Table and working out where to put the electrons, something odd happens after argon.

At argon, the 3s and 3p levels are full, but rather than fill up the 3d levels next, the 4s level fills instead to give potassium and then calcium. Only after that do the 3d levels fill.

The elements in the Periodic Table which correspond to the d levels filling are called d block elements. The first row of these is shown in the shortened form of the Periodic Table below.




The electronic structures of the d block elements shown are:



You will notice that the pattern of filling isn't entirely tidy! It is broken at both chromium and copper.

Note: Why is this so? People sometimes say that a half-filled d level as in chromium (with one electron in each orbital) is stable, and so it is - sometimes! But you then have to look at why it is stable. The obvious explanation is that chromium takes up this structure because separating the electrons minimises the repulsions between them - otherwise it would take up some quite different structure.

Transition metals
Not all d block elements count as transition metals! There are discrepancies between various syllabuses, but the majority use the definition:

A transition metal is one which forms one or more stable ions which have incompletely filled d orbitals.

On the basis of this definition, scandium and zinc don't count as transition metals - even though they are members of the d block.

Scandium has the electronic structure [Ar] 3d14s2. When it forms ions, it always loses the 3 outer electrons and ends up with an argon structure. The Sc3+ ion has no d electrons and so doesn't meet the definition.

Zinc has the electronic structure [Ar] 3d104s2. When it forms ions, it always loses the two 4s electrons to give a 2+ ion with the electronic structure [Ar] 3d10. The zinc ion has full d levels and doesn't meet the definition either.

By contrast, copper, [Ar] 3d104s1, forms two ions. In the Cu+ ion the electronic structure is [Ar] 3d10. However, the more common Cu2+ ion has the structure [Ar] 3d9.

Copper is definitely a transition metal because the Cu2+ ion has an incomplete d level.

Transition metal ions
You have already come across the fact that when the Periodic Table is being built, the 4s orbital is filled before the 3d orbitals. This is because in the empty atom, 4s orbitals have a lower energy than 3d orbitals.

However, once the electrons are actually in their orbitals, the energy order changes - and in all the chemistry of the transition elements, the 4s orbital behaves as the outermost, highest energy orbital.

The reversed order of the 3d and 4s orbitals only applies to building the atom up in the first place. In all other respects, you treat the 4s electrons as being the outer electrons.

Remember this:

When d-block elements form ions, the 4s electrons are lost first.

To write the electronic structure for Co2+:

Co [Ar] 3d74s2
Co2+ [Ar] 3d7

The 2+ ion is formed by the loss of the two 4s electrons.


To write the electronic structure for V3+:

V [Ar] 3d34s2
V3+ [Ar] 3d2

The 4s electrons are lost first followed by one of the 3d electrons.