Octahedral and Tetrahedral Complex Ions Revision Notes | A Level AQA Chemistry

1. Why Shape Matters in Complex Ions

Picture a transition metal ion sitting in solution, surrounded by a crowd of molecules and ions all trying to donate their lone pairs. The shape that crowd arranges itself into is not random — it is governed by one key factor: the size of the ligands relative to the central metal ion.

At A Level, you need to know two shapes in detail: octahedral and tetrahedral. The rule is straightforward — small ligands pack six around the metal ion; larger ligands can only fit four. Understanding why this happens is what separates a confident exam answer from a vague one.

2. Octahedral Complexes: Six Small Ligands

When the ligands are small — water ($H_2O$) and ammonia ($NH_3$) are the classic AQA examples — six of them can fit around the central metal ion without their electron clouds clashing significantly. The result is an octahedral complex with a coordination number of 6.

Think of the octahedron as two square pyramids joined at their bases. The metal ion sits at the centre, with four ligands in a square plane, one ligand above the plane and one below it. Adjacent ligands are 90° apart, while ligands directly opposite each other are 180° apart.

The hexaaquacopper(II) ion, $[Cu(H_2O)_6]^{2+}$, is the species responsible for the pale blue colour of copper sulphate solution. The hexaaquachromium(III) ion, $[Cr(H_2O)_6]^{3+}$, and the hexaaquairon(III) ion, $[Fe(H_2O)_6]^{3+}$, are further examples you will encounter across the AQA specification.

Ammonia forms octahedral complexes in exactly the same way. When excess ammonia is added to a copper(II) solution, the water ligands are displaced and the deep blue $[Cu(NH_3)_4(H_2O)_2]^{2+}$ ion forms — though note that this particular complex is actually a distorted octahedron with four ammonia ligands in the equatorial plane and two water ligands in the axial positions.

3. Tetrahedral Complexes: Four Larger Ligands

Now consider the chloride ion, $Cl^-$. It is significantly larger than water or ammonia — its ionic radius is much greater, and its electron cloud takes up considerably more space around the central metal ion. Because of this steric bulk, only four chloride ligands can fit around the metal ion before repulsion between their electron clouds becomes too great.

Four ligands arrange themselves to maximise the distance between their electron clouds, and the geometry that achieves this is the tetrahedron. The bond angles in a tetrahedral complex are 109.5°, which is larger than the 90° found in an octahedral complex — the ligands are further apart, reducing electron-cloud repulsion.

The tetrachlorocuprate(II) ion, $[CuCl_4]^{2-}$, is the standard AQA example of a tetrahedral complex. It forms when concentrated hydrochloric acid is added to a copper(II) solution — the high concentration of $Cl^-$ ions drives the ligand substitution. The colour changes from the pale blue of $[Cu(H_2O)_6]^{2+}$ to the yellow-green of $[CuCl_4]^{2-}$, which is a useful observation for exam questions on ligand substitution.

4. Comparing Octahedral and Tetrahedral Geometries

Now that both shapes are established, it is worth placing them side by side. The key variable is ligand size — everything else (coordination number, bond angle, and the examples you need to recall) follows from that single factor.

A useful way to remember the bond angles: the tetrahedral angle (109.5°) is larger than the octahedral angle (90°) because there are fewer ligands spread over the same space around the metal ion. Fewer ligands means each one can sit further from its neighbours.

5. The Ligand Substitution Link: Octahedral to Tetrahedral

The two geometries are not isolated facts — they connect directly to ligand substitution reactions, which appear frequently in AQA exam questions. When water ligands in an octahedral complex are replaced by larger ligands such as $Cl^-$, the coordination number drops from 6 to 4 and the geometry changes from octahedral to tetrahedral.

The equation for the conversion of the hexaaquacopper(II) ion to the tetrachlorocuprate(II) ion is:

The equilibrium sign is important here — the reaction is reversible. Adding water to $[CuCl_4]^{2-}$ shifts the equilibrium back to the left, regenerating the octahedral aqua complex. This reversibility, and the associated colour change, is a classic AQA observation question.

6. Worked Example: Identifying and Justifying Complex Ion Geometry

Exam questions on this topic often ask you to predict the shape of a complex ion and justify your answer. The following walkthrough models the reasoning chain an examiner expects.

Worked example: Predicting and Justifying the Shape of a Complex Ion

Predict the shape of the complex ion formed when $Fe^{3+}$ reacts with an excess of $NH_3$ molecules. Give the coordination number, bond angle, and justify the shape in terms of ligand size. [4 marks]

1. Ligand = $NH_3$ (ammonia). Ammonia is a small, neutral molecule. — Ligand size is the key factor that determines whether a complex is octahedral or tetrahedral. $NH_3$ is classified as a small ligand alongside $H_2O$. [B1]
2. Small ligands → octahedral geometry. Six $NH_3$ ligands can fit around $Fe^{3+}$ without excessive electron-cloud repulsion. — Because $NH_3$ is small, six molecules can coordinate to the metal ion. This gives a coordination number of 6 and an octahedral shape. [M1]
3. Coordination number = 6. Bond angle = 90°. — These values are fixed for all regular octahedral complexes. [B1]
4. $[Fe(NH_3)_6]^{3+}$ — $Fe^{3+}$ retains its 3+ charge; $NH_3$ is neutral, so the overall charge on the complex is 3+. [A1]

Answer: $[Fe(NH_3)_6]^{3+}$; octahedral; coordination number = 6; bond angle = 90°. Justified because $NH_3$ is a small ligand, so six molecules can fit around $Fe^{3+}$ without excessive repulsion between their electron clouds.

7. Exam Technique and Common Mistakes

Questions on complex ion geometry are reliable mark-earners if you know the two rules and can apply them precisely. The following guidance covers the most common ways students drop marks on this topic.

7a. High-Order Exam Practice

The following question mirrors the style of a 4-mark AQA 'explain' question. Use the marking points as a checklist — each one represents a distinct piece of chemical reasoning the examiner is looking for.

The two spec requirements for this topic — octahedral complexes with small ligands and tetrahedral complexes with larger ligands — are both addressed in that model answer. Notice how each point links ligand size explicitly to coordination number and geometry: that chain of reasoning is what earns marks, not naming the shapes.