Equilibrium in Physical Processes

Full Notes

Introduction to Equilibrium in Physical Processes

Equilibrium is not just a feature of chemical reactions – it is also crucial in physical changes, such as melting, boiling, and dissolving. These are all reversible processes.

When the rate of the forward change equals the rate of the reverse, the system is said to be in dynamic equilibrium. Though the macroscopic state appears unchanging, molecules are constantly shifting between phases or states.

6.1.1 Solid–Liquid Equilibrium

At 0 °C, when ice and water coexist in a beaker, the process of melting and freezing occurs simultaneously:

• If energy is supplied, more ice melts and if energy is removed, more water freezes.
• At constant temperature and pressure, the rate of melting = rate of freezing.
• No net change in mass of either phase is observed.
• This is an example of dynamic equilibrium in a closed system.

Key Point: Temperature must remain constant for equilibrium to be maintained.

6.1.2 Liquid–Vapour Equilibrium

When water is placed in a closed container, molecules escape into the air (evaporation) and simultaneously return to the liquid (condensation).

• Initially, evaporation dominates. As vapour pressure builds, condensation increases.
• Eventually, both rates equalise, reaching equilibrium.
• The pressure exerted by vapour at this point is called the equilibrium vapour pressure.

Vapour Pressure Depends On:

• Temperature (direct relation)
• Nature of liquid (more volatile = higher vapour pressure)

Important Observation:
In a closed system, the amount of vapour and liquid may remain constant at equilibrium, but individual molecules continue changing phase.

Boiling Point

The boiling point is the temperature at which a liquid’s vapor pressure equals the atmospheric pressure acting on it. At this point, molecules throughout the liquid – not just at the surface – have enough energy to overcome intermolecular forces and escape into the gas phase.

Below the boiling point, vaporization (evaporation) only occurs at the surface.

When vapor pressure equals atmospheric pressure, bubbles of vapor form within the liquid — this is boiling.

6.1.3 Solid–Vapour Equilibrium

Some solids, like camphor, iodine, ammonium chloride, and naphthalene, directly convert into vapour without becoming liquid. This process is called sublimation.

For Example If left in a closed beaker, a solid sample of iodine (I₂) will sublime and form a mixture of I₂(s) and I₂(g), with the gas being a purple vapour.

Solid ⇌ Vapour

• In a closed container, the vapour exerts equilibrium vapour pressure just like in liquid–vapour systems.
• The rate of sublimation = rate of deposition at equilibrium.
• Visible evidence includes constant mass of solid and stable vapour smell/intensity.

6.1.4 Equilibrium Involving Dissolution of Solids or Gases in Liquids

This category includes two important equilibrium systems:

(a) Solids in Liquids

When a solid like salt or sugar is added to water, it dissolves. However, after some time no more will dissolve when added — the solution becomes saturated.

Solid ⇌ Dissolved ions/molecules

• At saturation, the rate of dissolution (solid dissoling) = rate of crystallisation (solid reforming).
• This is a dynamic process — even though no more dissolves, particles are still exchanging between solid and solution.

Factors Affecting Equilibrium:

• Temperature (most salts are more soluble at higher temperatures).
• Nature of solvent and solute.

(b) Gases in Liquids

When a gas like CO₂ is dissolved in a liquid like soda water:

Gas ⇌ Dissolved Gas

• At equilibrium, the rate of gas entering the solution equals the rate of escape.
• The amount of gas dissolved is governed by Henry’s Law: Solubility ∝ Pressure of Gas Above Liquid
• Henry’s Law Constant varies with temperature and type of gas.

Everyday Example: Fizz in soda

Fizz in soda is due to CO₂ under high pressure. When you open the bottle, pressure drops and CO₂ escapes, causing the amount of fizz to reduce.

6.1.5 General Characteristics of Equilibria Involving Physical Processes

The following key features define physical equilibria:

• Dynamic Nature:
  • Even when the system appears static, particles continuously shift between phases.
  • No net change in measurable quantities (like pressure, concentration, volume) occurs at equilibrium.
• Achievable Only in a Closed System:
  • Physical equilibria require a closed system, so that matter (e.g., vapour, solute) is not lost to the surroundings.
  • In open systems, equilibrium may never be achieved.
• Measurable Quantities Remain Constant:
  • Parameters like temperature, pressure, and composition remain steady as long as the equilibrium is undisturbed.
• Specific Conditions:
  • Each equilibrium is defined by its own specific set of conditions.
  • Changing temperature or pressure can shift the position of equilibrium.
• Equilibrium is Reversible:
  • Processes can proceed in either direction depending on external changes.
  • This reversibility is a hallmark of equilibrium.

Summary

• Physical equilibrium is dynamic with forward and reverse rates equal.
• Vapour pressure and boiling depend on temperature and liquid identity.
• Saturated solutions have equal dissolution and crystallisation rates.
• Gas solubility in liquids follows Henry’s Law with pressure dependence.