Transition metals: electron configuration and oxidation number

In a nutshell 

Transition metal ions have partially filled $$d$$ sub-shells. Electrons pair up in shells only if necessary. Transition metal ions can have different oxidation numbers as the ionisation energies of removing successive electrons does not increase substantially. 

Key words 

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d-block elements

Transition metals make up the bulk of the d-block elements in the periodic table. All of the d-block elements in period four are transition metals apart from scandium and zinc.

Transition metals have similar physical properties. They all have a high density as well as high melting and boiling points. The ionic radii of transition metals are very similar.

d sub-shell

The $$d$$ sub-shell of an atom can hold up to ten electrons. An element can only be classed as a transition metal if it can form at least one stable ion which has one to nine electrons in the $$d$$ sub-shell. 

The incomplete $$d$$ sub-shell of transition metals give them special chemical properties. Transition metals can have variable oxidation numbers, and can form complex ions. They also form coloured ions and are good catalysts.

Electron configuration 

The diagram below shows the electron configuration of transition metals. The ​$$4s$$orbitals of period four transition metals usually fill up before the $$3d$$ orbitals. Electrons prefer to singly occupy orbitals and double up only if they have to.

Chromium is more stable having singly occupied ​$$4s$$ and $$3d$$​ orbitals rather than doubly occupied orbitals. Copper is stable with a full $$3d$$ sub-shell and one electron in the $$4s$$ sub-shell.

4_s

Scandium and zinc 

The electron configuration of scandium is $$[Ar]3d^14s^2$$. Scandium can only form a $$Sc^{3+}$$ ion which has an empty $$3d$$ sub-shell, therefore it is not a transition metal. The electron configuration of a $$Sc^{3+}$$ ion is $$[Ar]$$.

The electron configuration of zinc is $$[Ar]3d^{10}4s^2$$. Zinc can only form a $$Zn^{2+}$$ ion. When a ​Zn^{2+}$$Zn^{2+}$$ ion forms, the $$4s$$ sub-shell loses two electrons and the $$3d$$ sub-shell remains full. 

Oxidation numbers

Period four transition metal ions form when electrons are lost from the $$4s$$​ and $$3d$$​sub-shells. The $$4s$$​ and $$3d$$​ sub-shells have similar energies, therefore similar amounts of energy is required to remove an electron from $$4s$$​ and $$3d$$​ sub-shells.

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Example

​The electron configuration of vanadium is $$[Ar]3d^{3}4s^2$$. The ionisation energy does not increase significantly when five successive electrons are removed since the five electrons are being removed from the $$3d$$​ and $$4s$$​ sub-shells.  

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There is a large increase between the fifth and sixth ionisation energy as the sixth electron is being removed from the $$3p$$​ sub-shell which is closer to the nucleus. ​

Ionisation energies of removing successive electrons from $$s$$​ and $$d$$ sub-shells do not increase substantially, therefore, multiple electrons can be removed and ions with different oxidation numbers can form.

Example

Manganese has various oxidation numbers. 

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Higher oxidation numbers  

​Compounds and complexes will contain an ion with a given oxidation number if the energy released upon the formation of the compound/complex is greater than the ionisation energy.

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Forming transition metal ions with higher oxidation numbers involves higher ionisation energies. However, as ionic charge increases, the energy released upon the formation of a compound/complex also increases; this counteracts the higher ionisation energies involved.

[Ar]3d^{3}4s^[Ar]3d^{3}4

Sc^{3+}