Topic 8.11 Notes – pH and Solubility

1. How pH Can Change the Solubility of a Salt

Every slightly soluble salt dissolves according to a solubility equilibrium:

$$\ce{MX(s) <=> M^{n+}(aq) + X^{m-}(aq)}$$

That equilibrium has a Ksp expression. If ion concentrations change, the reaction shifts according to Le Châtelier’s principle.

Here’s the key connection:

• pH only matters if one of the dissolved ions reacts with H⁺ or OH⁻.
• If H⁺ or OH⁻ removes (or adds) one of the ions in the equilibrium, then Q changes relative to Ksp.
• When $Q < K_{sp}$, more solid dissolves.
• When $Q > K_{sp}$, precipitation is favored.

If neither ion reacts with H⁺ or OH⁻, changing pH does nothing to solubility. (A common ion effect can still matter, but that’s separate from pH.)

So the whole question becomes: Does one of the ions behave as an acid or base?

2. When pH Affects Solubility

a. Salts with Basic Anions

These anions are the conjugate bases of weak acids, such as:

• $\ce{CO3^{2-}}$ (from $\ce{H2CO3}$)
• $\ce{F^{-}}$ (from HF)
• $\ce{CN^{-}}$ (from HCN)
• $\ce{CH3COO^{-}}$ (from acetic acid)

Because they come from weak acids, they are basic and react with H⁺:

$$\ce{A^{-} + H^{+} -> HA}$$

In an acidic solution:

• H⁺ removes $\ce{A^{-}}$.
• $[\ce{A^{-}}]$ decreases.
• $Q < K_{sp}$.
• The equilibrium shifts right.
• Solubility increases.

The weaker the original acid, the stronger the base, and the bigger this effect.

b. Metal Hydroxides

Metal hydroxides dissolve like this:

$$\ce{M(OH)_n(s) <=> M^{n+}(aq) + nOH^{-}(aq)}$$

In an acidic solution:

• H⁺ reacts with $\ce{OH^{-}}$ to form water.
• $[\ce{OH^{-}}]$ drops.
• $Q$ decreases.
• The reaction shifts right.
• Solubility increases.

This is why many metal hydroxides dissolve better in acid. The diagram below shows H⁺ removing $\ce{OH^{-}}$, which lowers $Q$ and drives the dissolution equilibrium to the right.

Effect of acid on metal hydroxide solubility

c. Salts with Acidic Cations

Some cations act as weak acids. The classic example is:

$$\ce{NH4^{+} <=> NH3 + H^{+}}$$

$\ce{NH4^{+}}$ is the conjugate acid of the weak base $\ce{NH3}$.

In a basic solution:

• OH⁻ reduces $[\ce{H^{+}}]$.
• This shifts the $\ce{NH4^{+}}$ equilibrium to the right.
• $[\ce{NH4^{+}}]$ decreases.
• $Q < K_{sp}$.
• The solid dissolves more.
• Solubility increases in basic solution.

So salts containing acidic cations become more soluble in base.

d. Ions from Strong Acids or Strong Bases

Examples:

• $\ce{Na^{+}}$, $\ce{K^{+}}$
• $\ce{Cl^{-}}$, $\ce{NO3^{-}}$

These are neutral spectators. They do not react with H⁺ or OH⁻.

Result:

• Solubility is unaffected by pH.

If solubility changes here, it’s because of a common ion, not pH itself.

3. Acidic vs Basic Conditions Summary

Also remember:

• Hydroxides are less soluble in basic solution due to the common ion $\ce{OH^{-}}$.
• Salts with basic anions are less soluble in basic solution for the same reason.

4. How to Answer AP Questions About pH and Solubility

This topic is qualitative. No math with pH-dependent Ksp is required.

When you see one of these questions:

1. Write the dissolution equilibrium.
2. Ask: does any ion react with H⁺ or OH⁻?
3. Decide whether that ion’s concentration increases or decreases.
4. State:
   • What happens to Q
   • Which direction equilibrium shifts
   • Whether solubility increases or decreases

Full-credit explanations always connect:

• Ion concentration change
• $Q$ vs $K_{sp}$
• Direction of shift

If you skip one of those links, you usually lose the point on FRQs.