Metal Aqua Ions and Acidity Revision Notes | A Level AQA Chemistry

1. Why Dissolved Metal Ions Make Solutions Acidic

Dissolve iron(III) sulphate in water and measure the pH — you will find it is noticeably acidic, well below 7. No acid was added. So where does the acidity come from?

The answer lies in what happens at the molecular level when a metal ion enters water. Metal ions do not float freely in solution; they attract water molecules strongly and form metal aqua ions. These complexes then behave as Brønsted-Lowry acids, releasing H⁺ into solution and lowering the pH.

Understanding this requires two ideas working together: the geometry of the complex ion, and the effect of the metal's charge and size on the water molecules it holds. Build both of these clearly and the acidity argument follows naturally.

2. The Metal Aqua Ions You Need to Know

In aqueous solution, the four metal ions specified by AQA each form an octahedral complex with six water ligands. The water molecules coordinate through the oxygen atom's lone pair, pointing the O–H bonds outward away from the metal centre.

Notice the pattern: the 3+ ions (Al³⁺ and Fe³⁺) are both smaller and more highly charged than their 2+ counterparts. That combination is the key to understanding why 3+ aqua ions are more acidic — but to see why, you first need to understand what the metal ion actually does to the water molecules around it.

For AQA, the named [M(H₂O)₆]²⁺ ions you must know are limited to M = Fe and Cu, and the named [M(H₂O)₆]³⁺ ions are limited to M = Al and Fe. Make that explicit in written answers when you are asked to list or compare the specification examples.

The parallel species for 3+ metals are [Al(H₂O)₆]³⁺ and [Fe(H₂O)₆]³⁺. These are the only four complexes you are required to know for AQA A Level, so focus your recall on these specific ions.

3. How Metal Aqua Ions Act as Acids

Picture the [Fe(H₂O)₆]³⁺ ion. The Fe³⁺ centre carries a 3+ charge concentrated in a relatively small volume. This intense positive charge exerts a powerful pull on the electron density of the surrounding water molecules — specifically, it draws electron density away from the O–H bonds of the coordinated water.

When electron density is pulled away from an O–H bond, that bond becomes more polarised: the hydrogen atom becomes more δ⁺ and is more easily released as H⁺. A free water molecule in the surrounding solvent acts as a base and accepts this proton, generating H₃O⁺ and lowering the pH.

The equilibrium for this process can be written as follows, using iron(III) as the example:

The same type of equilibrium applies to all four aqua ions. The position of this equilibrium — how far it lies to the right — determines how acidic the solution is. For 3+ ions, the equilibrium lies further to the right than for 2+ ions, producing a lower pH.

4. Charge/Size Ratio: The Core Explanation

Here is where the real exam marks are earned. The question AQA asks most often is: why is [M(H₂O)₆]³⁺ a stronger acid than [M(H₂O)₆]²⁺? The answer must be framed in terms of charge/size ratio.

Charge/size ratio is a measure of how concentrated the positive charge is on the metal ion. A higher charge means a stronger electric field; a smaller ionic radius means that field is concentrated over a shorter distance. Both factors increase the polarising power of the metal ion on the coordinated water molecules.

Compare Fe²⁺ (charge 2+, radius ~0.78 Å) with Fe³⁺ (charge 3+, radius ~0.65 Å). The 3+ ion has both a higher charge and a smaller radius — its charge/size ratio is significantly greater. This means Fe³⁺ pulls electron density away from the O–H bonds of its coordinated water molecules much more strongly than Fe²⁺ does.

The result: the O–H bonds in [Fe(H₂O)₆]³⁺ are more polarised, the hydrogen atoms are more δ⁺, and proton donation to solvent water is more favourable. The acid dissociation equilibrium lies further to the right, and the solution is more acidic.

Worked example: Constructing a Charge/Size Ratio Explanation for an Exam Question

Explain why [Al(H₂O)₆]³⁺ produces a more acidic solution than [Fe(H₂O)₆]²⁺. [4 marks]

1. Al³⁺: charge = 3+, ionic radius ≈ 0.54 Å. Fe²⁺: charge = 2+, ionic radius ≈ 0.78 Å. Al³⁺ has a higher charge and a smaller radius, so a greater charge/size ratio. — [B1] — Identifying that Al³⁺ has a higher charge/size ratio than Fe²⁺ is the foundation of the argument.
2. A greater charge/size ratio means Al³⁺ has greater polarising power on the coordinated water ligands. — [M1] — The examiner needs to see the explicit link between charge/size ratio and polarising power.
3. Al³⁺ withdraws electron density from the O–H bonds of the coordinated water molecules more strongly, making those bonds more polarised and the H atoms more δ⁺. — [M1] — This is the mechanistic step: polarisation of the O–H bond is the direct cause of increased acidity.
4. The more polarised O–H bonds in [Al(H₂O)₆]³⁺ make proton donation to solvent water more favourable. The equilibrium lies further to the right, producing a higher concentration of H₃O⁺ and a lower pH. — [A1] — Completing the argument by connecting back to the observable acidity difference.

Answer: [Al(H₂O)₆]³⁺ is more acidic because Al³⁺ has a higher charge/size ratio than Fe²⁺, giving it greater polarising power. This withdraws more electron density from the O–H bonds of coordinated water, making proton donation more favourable and producing a higher [H₃O⁺] in solution.

This four-step structure — charge/size ratio → polarising power → O–H bond polarisation → proton donation — is the complete argument. Every mark in an AQA explain question on this topic maps onto one of these steps.

5. Comparing the Four Aqua Ions

Now that the charge/size ratio argument is established, you can apply it to compare any pair of the four specified ions. The general rule is clear: 3+ ions are more acidic than 2+ ions. Within each charge group, the ion with the smaller radius has the higher charge/size ratio and is therefore the stronger acid.

Al³⁺ sits at the top of this ranking because it combines the highest charge with the smallest ionic radius of the four. Fe²⁺ sits at the bottom because it has the lowest charge and the largest radius. When comparing Fe²⁺ and Fe³⁺ specifically, the charge difference is the dominant factor since both are iron — the 3+ ion is both more charged and smaller.

When comparing Cu²⁺ and Fe²⁺, both carry a 2+ charge, so the smaller radius of Cu²⁺ gives it a marginally higher charge/size ratio and slightly greater acidity. This level of detail is rarely tested directly, but the reasoning method is identical.

6. Exam Technique and Common Mistakes

AQA questions on this topic almost always use the command word 'explain', which means you must give reasons — not descriptions. Stating that '3+ ions are more acidic' without explaining why in terms of charge/size ratio will not earn the mechanistic marks.

The mark scheme for a 4-mark explain question on this topic typically awards one mark each for: identifying the higher charge/size ratio; linking this to greater polarising power; explaining the effect on O–H bond polarisation; and concluding that proton donation is more favourable. Missing any one of these steps costs a mark.

This topic rewards students who can construct a logical chain of reasoning rather than recalling a fact. The charge/size ratio argument is the thread that connects every mark in an explain question — learn to write it fluently and you will pick up full marks consistently.

These numerical values are not required by the AQA specification, but working through them confirms the ranking: Al³⁺ > Fe³⁺ > Cu²⁺ > Fe²⁺ in terms of charge/size ratio, which maps directly onto the acidity ranking of their hexaaqua complexes.

6a. Writing the Equation — Getting the Formula Right

One area where marks are frequently dropped is writing the equation for the acid dissociation. The product complex must be written correctly: one water ligand is replaced by a hydroxo (OH⁻) ligand, and the charge on the complex decreases by one.

For [Fe(H₂O)₆]²⁺ specifically, the product is [Fe(H₂O)₅(OH)]⁺ (charge drops from 2+ to 1+). For [Al(H₂O)₆]³⁺, the product is [Al(H₂O)₅(OH)]²⁺ (charge drops from 3+ to 2+). Keeping track of the charge change is a reliable way to check your equation is balanced.

7. Bringing It Together: The Full Argument at a Glance

Every element of this topic connects back to one central idea: the metal ion's charge/size ratio determines how strongly it polarises the O–H bonds of its coordinated water ligands, which in turn determines how readily those ligands donate H⁺ to solvent water.

• The four specified aqua ions are [Fe(H₂O)₆]²⁺, [Cu(H₂O)₆]²⁺, [Al(H₂O)₆]³⁺, and [Fe(H₂O)₆]³⁺ — all octahedral with six water ligands.
• All four act as Brønsted-Lowry acids by donating H⁺ from a coordinated water molecule to a solvent water molecule.
• 3+ ions have a higher charge/size ratio than 2+ ions, giving them greater polarising power and making their solutions more acidic.
• The acid dissociation equilibrium lies further to the right for 3+ ions, producing a higher [H₃O⁺] and a lower pH.
• The product of acid dissociation is always [M(H₂O)₅(OH)]^(n-1)+ — one hydroxo ligand replaces one water ligand, and the complex charge decreases by one.

Mastering this topic means being able to construct the charge/size ratio argument fluently in both directions: explaining why a named 3+ ion is more acidic than a named 2+ ion, and writing the correct equation for the acid dissociation of any of the four specified complexes. Both skills are regularly assessed in AQA A Level papers.