5.3.1 d & e Ligands and Complex Ions

(d) explanation and use of the term ligand in terms of coordinate (dative covalent) bonding to a metal ion or metal, including bidentate ligands

{Examples should include: monodentate: H₂O, Cl^– and NH₃ ; bidentate: NH₂CH₂CH₂NH₂ (‘en’).}

{In exams, other ligands could be introduced.}

(e) use of the terms complex ion and coordination number and examples of complexes with:

(i) six-fold coordination with an octahedral shape

(ii) four-fold coordination with either a planar or tetrahedral shape

{Examples: Octahedral: many hexaaquo complexes, e.g. [Cu(H₂O)₆] ²⁺, [Fe(H₂O)₆] ³⁺

Tetrahedral: many tetrachloro complexes, e.g. CuCl₄^2– and CoCl₄^2–

Square planar: complexes of Pt, e.g. platin: Pt(NH₃)₂Cl₂ }

What does this mean?

Transition metals ions are able to accept electron pairs from other ions, or compounds.

This forms a dative covalent bond (a shared pair of electrons where both electrons were provided by one of the atoms).

The ion/compounds providing the electron pair is called a ligand.

The metal ion will surround itself with 2, 4 or 6 ligands to form a complex ion.

The number of ligands depends on several factors, not least of which is the size of the ligand.

The simplest ligands only donate one pair of electrons - so they form one dative bond and are said to be monodentate (one tooth).

They can be ions like the Chloride ion (Cl^-) or molecules like Water or Ammonia.

In the latter two cases the examiner would expect you to draw the bond from the atom that hold the lone pair - ie from the O atom in water and the N atom in Ammonia rather than the H atoms.

The co-ordination number of the complex ion is the number of dative (coordinate) bonds - this is not always the same as the number of ligands as some ligands have two or more pairs of electrons they can donate.

If the ligand can form two dative bonds it is said to be bidentate (two teeth)

Ethane- 1,2-diamine is the only example that you have to learn but the syllabus states that others may be introduced in exams.

So, you should notice that the pairs of electrons in bidentate ligands are (almost) always two carbon atoms apart - any closer and the ligand cannot "wrap itself" around the metal ion well enough.

If you were asked to label the binding sites on these ligands then you should (obviously) pick the atoms with lone pairs which (with the exception of carbonate ion) are two carbon atoms apart.

On the "Oxalate ion" (ethanedioate ion according to IUPAC) it wouldn't matter which 2 Oxygen atoms you picked so long as they weren't both on the same side of the molecule.

The syllabus no longer mentions hexadentate ligands but EDTA was once a staple of A level exams and it isn't unreasonable to suppose it might make an appearance

Here you can see lone pairs on two N atoms separated by two C atoms. There are also lone pairs of the O atoms in the acid groups. So six dative bonds can form.

The coordination number of this complex ion is still 6 even though there is only one ligand.

a - Octahedral

To form an octahedral complex 6 dative bonds are neede.

These can all come from a hexadentate ligand (as above),

From 6 monodentate ligands which can either be mixed or all identical as below.

Or, there could be 3 bidentate ligands

Or a mixture of bidentate and monodentate ligands

Or two tridentate ligands, or even a mixture of tridentate and monodentate ligands

It's safe to assume for A level that any metal ion can form a hexaqua complex - that is, a complex with six water ligands.

Here we see hexaqua copper (II).

The water ligands are neutral so the overall charge is still 2+.

The bond angles are all 90^o, and the examiner would expect to see an attempt at wedges and dots to show the spatial arrangement of the datvie bonds.

b- Tetrahedral

If the coordination number is four then the usual geometry of the ion will be tetrahedral.

These are the two examples mentioned on the syllabus. In both cases the central metal atom must have a 2⁺ charge because there are 4 Cl^- ligands and yet the overall charge is 2^- .

It's also perfectly possible to form a tetrahedral complex with bidentate ligands but not this isn't usual with the ones that are mentioned on the syllabus which have a better geometry for forming square planar complexes.

c - Square planar

If the coordination number is four then the geometry of the ion can be square planar.

We don't have to worry about Crystal Field theory.

So, for the sake of A level, we can say that Copper and Cobalt (both fairly cheap) tend towards tetrahedral whereas Platinum and Rhodium (expensive) tend to form square planar complexes.

In reality, it's the Chloride ions which tend to make tetrachloro-cobalt and tetrachloro-copper ions tetrahedral.

The only square planar complex we're expected to know about in any detail is Cis-Platin (above) - about which more in part (f,g).