L2 - Complex Ions

objectives

Understandings:

• The d sub-level splits into two sets of orbitals of different energy in a complex ion.
• Complexes of d-block elements are coloured, as light is absorbed when an electron is excited between the d-orbitals.
• The colour absorbed is complementary to the colour observed.

Applications and skills:

• Explanation of the effect of the identity of the metal ion, the oxidation number of the metal and the identity of the ligand on the colour of transition metal ion complexes.
• Explanation of the effect of different ligands on the splitting of the d-orbitals in transition metal complexes and colour observed using the spectrochemical series.

Guidance:

• The spectrochemical series is given in the data booklet in section 15. A list of polydentate ligands is given in the data booklet in section 16.
• Students are not expected to recall the colour of specific complex ions.
• The relation between the colour observed and absorbed is illustrated by the colour wheel in the data booklet in section 17.
• Students are not expected to know the different splitting patterns and their relation to the coordination number. Only the splitting of the 3-d orbitals in an octahedral crystal field is required.

Despite it's title this is not a particularly difficult area to learn!

Complexes

• A Complex Ion is an ion comprising of one (or more) Ligands attached to a central metal cation by a Dative Covalent Bond.
• A Ligand is a species which can use a lone pair of electrons to form a dative covalent bond with a transition metal.
• An examples of a Ligand is Water - the below shows some other simple ligands:

Why do some Cations form Complexes?

• For a Cation to form a complex ion it must have 2 features:
  • A High Charge Density - this attracts electrons from ligands
  • Empty Orbitals of Low Energy - To accept Lone Pairs from ligands

• As cations of d-block elements are small, have a high charge and available 4s and 4d orbitals of low energy they can easily form complex ions.

Coordination Number

• This is a brief note as the IB does not really venture far into complex ion chemistry ].
• The number of Dative Coordinate Bonds (or Lone Pairs a central transition metal ion can accept) is called the Coordinate Number.
• It is often the same as the number of ligands But not always as some ligands can form more than one dative covalent bond per molecule.
• EDTA and Ethanedioic acid are good examples of ligands that form more than one dative covalent bond per molecule.
• We can get coordination numbers of 2, 4 or 6.

What do Complex Ions look like and how do we draw them?

Examples of Complex ions can appear as:

• [Fe(H₂O)₆]⁺²
• [CoCl₄]^2-
• [Cu(NH₃)₄(H₂O)₂]⁺²

• Square Brackets - Used to show that the charge pf the ion is delocalised over the entire complex
• Complex Ion Charge - This is shown in the top right corner and is derived by subtracting the charge on the ligands from the charge on the central metal ion.

• As examples you should be able to work out the central ion given the complex ion formula:

• You should also note that complex ions can appear as Salts with other anions appearing outside the square brackets.
• EG - [Fe(H₂O)₅(Cl)]SO₄
• In these examples you will need to derive the:
  • Charge on the metal ion
  • Charge on the complex ion

For the above example we know that SO₄ has a charge of -2 so the charge on the complex ion must equal +2.

If the complex ion has a charge of +2 then the charge on the Fe must be +3

This is because it contains 5 water ligands (charge 0) and 1 Cl- ligand (charge -1)