Q.Why do gases always tend to be less soluble in liquids as the temperature is raised?

Concept: Solubility of gases in liquids is governed by thermodynamic principles and the nature of the dissolution process.

When a gas dissolves in a liquid, the process is typically exothermic — energy is released as gas molecules form interactions (dipole, dispersion forces, or hydrogen bonds) with solvent molecules. The enthalpy change $\Delta H_{\text{solution}} < 0$.

By Le Chatelier's principle, if we increase temperature (add heat) to an equilibrium system, the equilibrium shifts to counteract that change — it favors the endothermic direction. Since dissolution releases heat, the reverse process (gas escaping from solution) is endothermic and becomes favored at higher temperatures.

Consequently, raising the temperature drives dissolved gas molecules back into the gas phase, decreasing solubility. …

Why heating drives gases out of solution

When a gas dissolves in a liquid, the process nearly always releases energy. Think about what happens at the molecular level: gas molecules in the vapor phase are far apart, moving freely with essentially no intermolecular forces acting on them. When they enter the liquid, they become surrounded by solvent molecules and form new interactions—hydrogen bonds with water, dipole interactions, or at minimum van der Waals forces. Forming these attractions releases energy as heat.

This makes gas dissolution an exothermic process. The equilibrium we're dealing with is:

Le Chatelier's principle tells us exactly what happens when we add heat to an exothermic equilibrium: the system shifts to counteract the disturbance. If we raise the temperature, we're effectively adding heat to the right side of the equation. The equilibrium responds by shifting left—toward the gaseous state—to absorb that excess energy.

The thermodynamic picture

Henry's Law quantifies gas solubility:

where $p$ is the partial pressure of the gas, $x$ is its mole fraction in solution, and $K_H$ is Henry's constant. The key insight: $K_H$ increases with temperature for nearly all gas-liquid systems.

A larger $K_H$ means you need a higher partial pressure to maintain the same dissolved concentration. Equivalently, at fixed pressure, the mole fraction $x$ must decrease. The gas becomes less soluble.

We can see this through the van 't Hoff equation, which relates how an equilibrium constant changes with temperature:

For gas dissolution, $\Delta H^\circ < 0$ (exothermic), so $\frac{d \ln K_H}{dT} > 0$. Henry's constant climbs as temperature rises.

Everyday observations

1. Warm soda goes flat faster. Carbon dioxide escapes more readily from a warm soft drink than a cold one. Refrigeration keeps the $\text{CO}_2$ dissolved.

2. Boiling drives out dissolved air. When you heat water, you see bubbles forming well before it boils—that's dissolved air (mostly $\text{N}_2$ and $\text{O}_2$) being expelled as solubility drops.

3. Thermal pollution harms aquatic life. Power plants discharge warm water into rivers. The elevated temperature reduces dissolved oxygen, stressing fish and other organisms that depend on it. …

Method: Le Chatelier’s Principle Applied to Dissolution Equilibrium

This is the clearest way to explain the temperature effect on gas solubility.

Step 1 — Write the dissolution as an equilibrium

When a gas dissolves in a liquid, the process is:

This is an exothermic process for most gases — heat is released when gas molecules enter the liquid.

Step 2 — Identify the disturbance

When temperature is raised, you are adding heat to the system.

Step 3 — Apply Le Chatelier’s Principle

If a system at equilibrium is disturbed, it shifts in the direction that opposes the disturbance.

• The forward reaction (dissolution) releases heat.
• Adding heat favours the endothermic direction — the reverse reaction.
• So the equilibrium shifts backward: gas molecules leave the liquid.

Step 4 — Consequence

More gas escapes → less gas remains dissolved. …

Common Mistakes on Henry's Law & Temperature-Solubility Relationship

✗ Mistake 1: Confusing Henry's Law with Le Chatelier's Principle

The error: Students often say "Henry's Law states solubility decreases with temperature" — but Henry's Law itself says nothing about temperature. It only relates pressure to solubility at a fixed temperature.

Correct understanding:

• Henry's Law: $P = K_H \cdot x$ (at constant $T$)
• The temperature effect comes from thermodynamics, not Henry's Law directly

How to avoid: Memorise clearly:

• Henry's Law → pressure vs solubility (constant T)
• Temperature effect → Le Chatelier's Principle applied to dissolution equilibrium

✗ Mistake 2: Forgetting the Dissolution is Exothermic

The error: Students can't explain why solubility decreases, so they just memorise "it happens."

The correct reasoning (step-by-step):

1. Gas dissolving in liquid is generally exothermic (releases heat):

2. By Le Chatelier's Principle: raising temperature adds heat → system shifts to consume heat → shifts backward (toward undissolved gas)

3. Result: Less gas remains dissolved

How to avoid: Always write the equilibrium with "Heat" as a product, then apply Le Chatelier.

✗ Mistake 3: Thinking All Gases Behave Identically

The error: Assuming every gas becomes less soluble when heated.

The truth: Most gases follow this rule, but ammonia ($NH_3$) and HCl are exceptions — they react chemically with water, so their dissolution is more complex.

How to avoid: Remember the general trend applies to non-reactive gases ($O_2$, $CO_2$, $N_2$). For reactive gases, check the specific reaction enthalpy.

✗ Mistake 4: Vague or Wrong Explanations in Exams

The error: Writing "solubility decreases because of Henry's Law" — this gets zero marks.

A perfect exam answer structure: …