Getting Started
The solubility of an ionic compound is often treated as a fixed value at a given temperature, but this is an oversimplification. In reality, the dissolution of a salt is an equilibrium process that can be influenced by other reactions occurring in the same solution. This chapter explores how the pH of a solution—a measure of its acidity or basicity—can dramatically alter the solubility of certain salts by interacting with the ions they produce.
What You Should Be Able to Do
After completing this section, you should be able to:
Predict whether a salt's solubility will increase, decrease, or remain unchanged when acid (H⁺) is added.
Predict whether a salt's solubility will increase, decrease, or remain unchanged when base (OH⁻) is added.
Identify which ion in a salt, if any, will react with added acid or base.
Justify predictions about solubility changes using Le Châtelier’s principle and chemical equilibrium concepts.
Key Concepts & Analysis
The effect of pH on solubility is a classic example of coupled equilibria, where a change in one system affects another. We can analyze this relationship using the lens of dynamics and change, focusing on how an equilibrium system responds to a stress.
Baseline Condition: The Solubility Equilibrium
Consider a sparingly soluble salt, such as calcium carbonate (CaCO₃), placed in water. It establishes a dissolution equilibrium between the solid and its constituent ions.
Equilibrium Reaction:
CaCO₃(s) ⇌ Ca²⁺(aq) + CO₃²⁻(aq)
In a saturated solution, the rate of dissolution equals the rate of precipitation. The concentrations of the aqueous ions are constant and are governed by the solubility product constant (Ksp). This represents the initial, stable state of the system.
The Process or Stress: Changing the pH
The equilibrium can be disturbed, or "stressed," by changing the concentration of H⁺ or OH⁻ ions in the solution.
Stress 1: Adding an Acid (Lowering the pH)
When a strong acid like HCl is added, the concentration of H⁺ ions increases. We must now consider if either of the ions from the salt, Ca²⁺ or CO₃²⁻, can react with H⁺.
Ca²⁺ is the cation of a strong base (Ca(OH)₂) and is therefore a neutral ion with no tendency to react with H⁺.
CO₃²⁻ (carbonate) is the conjugate base of a weak acid, the hydrogen carbonate ion (HCO₃⁻). As a weak base, it will react with the added H⁺.
Acid-Base Reaction:
CO₃²⁻(aq) + H⁺(aq) ⇌ HCO₃⁻(aq)
This second reaction is directly coupled to the first. The added acid effectively consumes one of the products (CO₃²⁻) of the dissolution equilibrium.
Stress 2: Adding a Base (Raising the pH)
When a strong base like NaOH is added, the concentration of OH⁻ ions increases. This stress is most relevant for two types of salts:
Metal Hydroxides: For a salt like magnesium hydroxide, Mg(OH)₂, adding base increases the concentration of a product ion (OH⁻). This is the common ion effect, which will be discussed in more detail elsewhere. It decreases solubility.
Salts with Acidic Cations: For a salt like ammonium chloride (NH₄Cl), the ammonium ion (NH₄⁺) is the conjugate acid of a weak base (NH₃). It will react with the added OH⁻.
Acid-Base Reaction:
NH₄⁺(aq) + OH⁻(aq) ⇌ NH₃(aq) + H₂O(l)
The Resulting Change: The Equilibrium Shift
According to Le Châtelier’s principle, when a stress is applied to a system at equilibrium, the system will shift to counteract the stress.
Response to Added Acid:
By consuming the carbonate ion (CO₃²⁻), the addition of acid reduces the concentration of a product in the dissolution equilibrium. To counteract this loss, the equilibrium CaCO₃(s) ⇌ Ca²⁺(aq) + CO₃²⁻(aq) will shift to the right. This shift causes more of the solid CaCO₃ to dissolve to replenish the CO₃²⁻ ions.
- Outcome: The solubility of calcium carbonate increases as the pH decreases (becomes more acidic).
Response to Added Base (for a salt with an acidic cation):
By consuming the ammonium ion (NH₄⁺), the addition of base reduces the concentration of a product in the dissolution equilibrium of NH₄Cl. The equilibrium NH₄Cl(s) ⇌ NH₄⁺(aq) + Cl⁻(aq) will shift to the right, causing more solid to dissolve.
- Outcome: The solubility of ammonium chloride increases as the pH increases (becomes more basic).
Salts Unaffected by pH:
Consider a salt like silver chloride, AgCl. It dissolves according to the equilibrium: AgCl(s) ⇌ Ag⁺(aq) + Cl⁻(aq).
Ag⁺ is a cation that is not significantly acidic.
Cl⁻ is the conjugate base of a strong acid (HCl) and is therefore a neutral ion.
Since neither ion reacts significantly with added H⁺ or OH⁻, changing the pH does not remove a product from the system. Therefore, the solubility equilibrium is not disturbed.
- Outcome: The solubility of silver chloride is essentially independent of pH.
| Type of Salt | Ion that Reacts with... | Effect of Lowering pH (Adding Acid) | Effect of Raising pH (Adding Base) | Example |
|---|---|---|---|---|
| Salt with a Basic Anion | Anion reacts with H⁺ | Solubility Increases | No significant effect | CaCO₃, BaF₂, Ag₃PO₄ |
| Metal Hydroxide | OH⁻ reacts with H⁺ | Solubility Increases | Solubility Decreases (Common Ion Effect) | Mg(OH)₂, Al(OH)₃ |
| Salt with an Acidic Cation | Cation reacts with OH⁻ | No significant effect | Solubility Increases | NH₄Cl |
| Salt with Neutral Ions | Neither ion reacts | No significant effect | No significant effect | NaCl, KBr, AgCl |
Key Models & Representations
This flowchart helps determine how pH will affect a salt's solubility.
graph TD
A[Start: Consider a salt, MX] --> B{Is the anion X⁻ the conjugate base of a WEAK acid?};
B -- Yes --> C[Solubility INCREASES in acidic solution (lower pH)];
B -- No --> D{Is the salt a hydroxide, M(OH)ₙ?};
D -- Yes --> C;
D -- No --> E{Is the cation M⁺ the conjugate acid of a WEAK base?};
E -- Yes --> F[Solubility INCREASES in basic solution (higher pH)];
E -- No --> G[Solubility is largely UNAFFECTED by pH];
To use this, you must know the common strong acids (e.g., HCl, HNO₃, H₂SO₄). If an anion is the conjugate of a strong acid, it is neutral; otherwise, it is basic.
Key Terms, Quantities, & Concepts
Solubility: The maximum concentration of a solute that can dissolve in a solvent at a specific temperature to form a saturated solution.
Solubility Product Constant (Ksp): The equilibrium constant for the dissolution reaction of a sparingly soluble ionic compound. Its value is constant at a given temperature.
Le Châtelier’s Principle: A principle stating that if a stress (such as a change in concentration, pressure, or temperature) is applied to a system at equilibrium, the system will shift in a direction that counteracts the stress.
Saturated Solution: A solution in which the dissolved solute is in dynamic equilibrium with the undissolved solid solute.
Conjugate Acid-Base Pair: Two chemical species that differ from each other only by the presence or absence of a single proton (H⁺). For example, HCO₃⁻ is the conjugate acid of CO₃²⁻.
Weak Acid/Base: An acid or base that only partially dissociates or ionizes in water. The conjugate of a weak species is itself a weak species.
pH: A logarithmic scale used to specify the acidity or basicity of an aqueous solution. Lower pH values indicate higher acidity (more H⁺).
Skill Snapshots
Causation
Cause: Adding strong acid to a saturated solution of barium fluoride (BaF₂).
Effect: The H⁺ ions react with the basic F⁻ ions, reducing their concentration. This causes the dissolution equilibrium to shift right, increasing the solubility of BaF₂.
Cause: The chloride ion (Cl⁻) is the conjugate base of a strong acid (HCl).
Effect: The Cl⁻ ion is a neutral ion and does not react with H⁺, so the solubility of salts like AgCl or PbCl₂ is not significantly affected by adding acid.
Cause: Lowering the pH of a solution containing solid aluminum hydroxide, Al(OH)₃.
Effect: H⁺ ions neutralize the hydroxide ions (OH⁻) to form water. This removal of a product ion drives the dissolution of more Al(OH)₃.
Comparison
The solubility of AgCl is unaffected by pH, whereas the solubility of AgF increases in acidic solution because F⁻ is a weak base.
Adding acid to Mg(OH)₂ increases its solubility by removing the OH⁻ product, while adding a base like NaOH decreases its solubility due to the common ion effect.
The carbonate ion (CO₃²⁻) is a relatively strong weak base and reacts readily with H⁺, while the nitrate ion (NO₃⁻) is a neutral anion and does not.
Change and Continuity
Baseline: A saturated solution of lead(II) phosphate, Pb₃(PO₄)₂(s), is at equilibrium with its ions, Pb²⁺(aq) and PO₄³⁻(aq).
Change 1: Nitric acid is added, lowering the pH. The phosphate ion (PO₄³⁻), a weak base, reacts with H⁺. The equilibrium shifts right, and more solid dissolves.
Change 2: Sodium hydroxide is then added, raising the pH. This neutralizes the excess acid and can convert HPO₄²⁻ back to PO₄³⁻, potentially causing Pb₃(PO₄)₂ to re-precipitate.
Continuity: Throughout these pH changes, the value of the Ksp for Pb₃(PO₄)₂ remains constant, as long as the temperature does not change.
Common Misconceptions & Clarifications
Misconception: The solubility of every ionic compound is affected by pH.
Clarification: Only salts containing an ion that is appreciably acidic or basic will have their solubility affected by pH. Salts formed from a strong acid and a strong base (e.g., NaCl, KBr, Ca(NO₃)₂) have neutral ions, and their solubility is pH-independent.
Misconception: Adding acid always makes things more soluble.
Clarification: Adding acid increases solubility only when the salt contains a basic anion (like F⁻, CO₃²⁻, or S²⁻) or is a hydroxide. For a salt with an acidic cation (like NH₄Cl), adding a base increases its solubility.
Misconception: You can determine if an anion is basic just by looking at its formula.
Clarification: An anion is basic if its corresponding acid is a weak acid. The most effective way to identify these is to memorize the short list of strong acids (HCl, HBr, HI, HNO₃, HClO₄, H₂SO₄). If an anion is the conjugate of one of these, it is neutral. All others should be treated as basic.
One-Paragraph Summary
The solubility of an ionic compound is not always a fixed constant but is an equilibrium that can be influenced by solution pH. This effect is significant only when the salt contains an ion that is a weak acid or a weak base. By applying Le Châtelier’s principle, we can predict these changes qualitatively. Adding an acid will react with and remove basic anions (e.g., CO₃²⁻, F⁻) or hydroxide ions from the solution, causing the dissolution equilibrium to shift right and increase the salt's solubility. Conversely, adding a base can remove acidic cations (e.g., NH₄⁺), also increasing solubility. For salts composed of neutral ions derived from strong acids and bases, such as NaCl, changes in pH have a negligible effect on their solubility.