# Transition Metals and Complexes¶

• Transition metals are defined by their many d orbitals that can be filled and configured in many different ways.
• This gives rise to many different properties, oxidations states, and stable ions
• Through the formation of complexes, they also tend to create highly coloured species

## Physical properties¶

• Transition metals tend to have similar properties, due to their relatively small difference in Zeff

## Oxidation states¶

• Possible oxidation states of a transition metal are based on their electronic configuration
• Namely, how many electrons they can stably lose to form a cation
• The s electrons will likely be removed before any of the d electrons (as per $$V^{2+}$$)
• An s electron will likely be promoted to complete the d orbitals (as per $$Cr$$)

E.g.

$$\hskip{1cm}Sc=[Ar]\:4s^2\:3d^1$$

$$\hskip{1cm}Ti=[Ar]\:4s^2\:3d^2$$

$$\hskip{1cm}V=[Ar]\:4s^2\:3d^3\hskip{2cm}V^{2+}=[Ar]\:4s^0\:3d^3$$

$$\hskip{1cm}Cr=[Ar]\:4s^1\:3d^5$$

E.g. Manganese can form a $$+7$$ ion when it loses all of its $$d$$ and $$s$$ electrons

• +6 refers to $$4s^0\:3d^1$$
• +4 refers to $$4s^0\:3d^3$$
• +2 refers to $$4s^0\:3d^5$$ which is stable due to it’s complete d orbitals and has only lost the $$4s$$ electrons

## Complexes¶

• Transition metals often form complexes with ligands that have lone pair electrons (Lewis bases)
• The lone pair electrons on the ligands occupy and hybridise higher angular momentum ($$l$$) orbital
• Though this theory is considered obsolete