Shapes of Complex Ions

Complex ions are molecules with a transition metal or its ion at the centre, surrounded by other molecules called ligands. Ligands determine the shapes of complex ions by forming coordinate bonds.

Sound complicated? Not to worry, complex ions will not seem complicated anymore once you have read this!

• We will learn how coordinate bonds determine the shapes of complex ions.
• We will explore the different types of ligands in complex ions, and consider the four shapes transition metal complexes can take.
• We will also discover how octahedral and square planar complex ions show stereoisomerism.
• Finally, we will learn how to draw structures of complex ions.

Complex metal ions

This means ligands donate a pair of nonbonding electrons. Ligands can be very simple like the water molecule, or they can be much more complex, like ethylenediamine (en) or the ethanedioate ion (ox).

Fig. 1 - Simple and complex ligands

Complexes consist of a transition metal or its ion bonded coordinately to ligands. Let us take a closer look at what we mean by coordinate bonding.

Bonding in transition metal complexes

Transition metal ions have a partially filled 3d sublevel. Ligands form coordinate bonds when an orbital that contains a lone pair of electrons overlaps with a vacant orbital on the metal ion. In other words, the metal acts as a Lewis acid by accepting a pair of electrons, and the ligand acts as a Lewis base by donating an electron pair.

For example, aluminium has the electronic configuration: 1s² 2s² 2p⁶ 3s² 3p¹.

Fig. 2 - Aluminium electron configuration

When it forms an Al³⁺ ion, its electronic structure is now 1s² 2s² 2p⁶, leaving all orbitals in the 3rd level empty. Al³⁺ can now accept lone pairs of electrons from a molecule/ion to stabilise it. In the example below, Al³⁺ accepts six lone pairs from six water molecules.

Fig. 3 - Hexaquaaluminium

We write the formula for this complex as [ Al (H₂O)₆ ]³⁺. Six water ligands each give the aluminium ion one lone pair of electrons. So, we say the complex [Al (H₂O)₆ ]³⁺ has a coordination number of 6.

Coordination number tells us the number of coordinate bonds in a complex ion. Coordination numbers also determine the shape of complex ions. Keep reading to learn more!

Common complex shapes of metal ions

Overall, there are four complex shapes of transition metal ions:

• Linear
• Square planar
• Tetrahedral
• Octahedral

The two most common ones are tetrahedral and octahedral.

Tetrahedral complex ions

Tetrahedral complexes have four coordinate bonds with bond angles of 109.5º. Large ligands like Cl^- form tetrahedral complexes.

Two examples of tetrahedral complexes are copper chloride [CuCl₄]^2-, and cobalt chloride [CoCl₄]^2-, shown below.

Fig. 7 - Tetrahedral complex ions

Notice how four Cl^- ions form a bond with the central metal ion. Both Cu (II) and Co (II) have a 2+ charge. Each Cl ion carries a 1- charge, thus this gives the overall charge of 2- on both complex ions.

(4 X Cl^- = 4- ; 4- + 2+ = 2- overall charge)

Octahedral complex ions

Octahedral complexes form when there are six coordinate bonds formed around the central metal. They have a bond angle of 90º.

Small ligands like H₂O or NH₃ form octahedral complexes. A small ligand like water or ammonia leaves space for more ligands to fit around the complex ion. On the other hand, large ligands like chlorine leave space only for four ligands.

There are many ways these six coordinate bonds can be formed. Maybe you have already thought of a few! Let us take a look at the possibilities.

• Six small monodentate ligands - each ligand forms one coordinate bond with the central metal. H₂O and NH₃ are examples of ligands that form one coordinate bond.

Fig. 8 - Octahedral complex ions

• Three bidentate ligands - each ligand forms two coordinate bonds with the central metal. The oxalate (ethanedioate) ion (ox) or ethylenediamine (en) are examples of an ion and molecule respectively, that can donate two lone pairs of electrons and form two coordinate bonds.
• One multidentate ligand - one ion/molecule can form three or more coordinate bonds to the metal ion. EDTA^4- is the most common example of a multidentate ligand.

In addition to the two shapes we have looked at, linear and square-shaped complexes can also form. Let us consider these now.

Linear complex ions

Linear complex ions form when there are two coordinate bonds. They have a bond angle of 180°.

A common example of a linear complex ion is diamminesilver(I) [Ag(NH₃)₂]⁺, which is used in Tollens’ reagent.

Fig. 9 - Linear complex ion

Square planar complexes

Square planar complexes have four coordinate bonds too! Unlike tetrahedral complexes, they have a bond angle of 90º.

An early drug used in the treatment of cancer, cis-platin, forms a square planar complex. It has the platinum ion, Pt²⁺, as its central ion, and two ammonia molecules and two chloride ions as ligands.

Fig. 10 - Cis-platin

You might have guessed from the name that cis-platin is a geometric isomer. Square planar and tetrahedral complexes are special because they can show stereoisomerism. We will explore what that means next.

Isomerism in transition metal complexes

Occasionally, we see stereoisomerism in octahedral and square planar complexes. Stereoisomers are molecules with the same order of atoms, but different spatial arrangements. There are two types of stereoisomerism: geometric E-Z isomerism and optical isomerism. Let us examine how they display in metal complexes.

Isomerism in transition metal complexes: Geometric isomerism

Square planar complexes can display geometric cis-trans isomerism (also called E/Z isomerism). These complexes have their high-priority ligands either adjacent to each other, or across from each other. For example, cis-platin and trans-platin are geometric isomers.

Fig. 11 - Cis and trans platin

Notice how cis-platin has the chlorine and ammonia ligands right next to each other. On the other hand, trans-platin has the chlorine and ammonia ligands opposite each other.

The isomer with the high priority ligands adjacent to each other is the cis-isomer, in this case, cis-platin. The isomer with the high priority ligands across from each other is the trans-isomer.

We can also see geometric isomerism in octahedral complex ions with monodentate ligands. Two of the ligands must be different from the other four. For example, cis and trans-tetraamminedichlorocobalt(III) ions shown below are octahedral geometric isomers.

Fig. 12 - Cis and trans-tetraamminedichlorocobalt(III) ion

You will notice how the two chlorine ions are adjacent to each other in the cis-isomer, while they are opposite each other in the trans-isomer.

Isomerism in transition metal complexes: Optical isomerism

Optical isomers are non-superimposable mirror images of each other. They show no plane of symmetry. We can see this type of stereoisomerism in octahedral complexes with bidentate ligands. Remember, bidentate means they donate two lone pairs of electrons.

Ethylenediamine (or ‘en’ for short) is an octahedral complex ion that shows optical isomerism. You might not be able to tell by looking at the structural formulas below, so making some physical models might be helpful!

Fig. 13 - Cis and trans-tetraamminedichlorocobalt(III) ion

With a little imagination, you can see that both molecules are mirror images of each other. But you cannot superimpose them no matter which way you turn them.

Before we conclude, let us take a look at how you might draw the structure of a complex ion.

How can we deduce the overall polarity of complexes that show isomerism?

You may notice that ligands are within two square brackets, and you also may notice that sometimes outside the brackets there is charge. It is important to remember that only some complexes may have an overall charge and this can be calculated using two things:

1. The charge of the ligands within the complex.
2. The charge on the central transition metal ion of the complex