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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
    • Leads to similar ionisation energies, atomic radii and ionic radii

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

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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