6.1. Equilibrium in Physical Process

The characteristics of system at equilibrium are better understood if we examine some physical processes. The most familiar examples are phase transformation processes, e.g., solid liquid, liquid gas, solid gas.

In physical processes, equilibrium is the state where the forward and reverse processes occur at equal rates. Key examples include changes in state, like the transition between solid and liquid or liquid and gas. At equilibrium, the amounts of each state do not change, even though molecular transitions are ongoing.

Consider melting ice in water. Although the amount of ice and water remains constant at 0°C, molecules continuously move from ice to water and vice versa. This is like a see-saw; when one side goes down, the other goes up, but overall balance is maintained.

Ice and water kept in a perfectly insulated thermos flask (no exchange of heat between its contents and the surroundings) at 273K and the atmospheric pressure are in equilibrium state and the system shows interesting characteristic features. We observe that the mass of ice and water do not change with time and the temperature remains constant. However, the equilibrium is not static. The intense activity can be noticed at the boundary between ice and water.

At 273K, ice and liquid water coexist in equilibrium. Although the total mass of ice and water stays constant, individual molecules are continuously transitioning between the solid and liquid phases. Water molecules collide with ice, some freeze, while some ice melts, ensuring that the amounts remain unchanged.

Think about a perfectly balanced see-saw; as one child gets on, another gets off to keep the see-saw level. In the same manner, when heat is applied to the ice, some sticks turn into water (melting), while others are being frozen again by water molecules.

This equilibrium can be better understood if we consider the example of a transparent box carrying a U-tube with mercury (manometer). After removing the drying agent by tilting the box on one side, a watch glass (or petri dish) containing water is quickly placed inside the box.

In the experimental setup, water evaporates, increasing the water vapour pressure inside the container until it reaches a constant value at equilibrium. The rate of evaporation equals the rate of condensation, leading to a stable vapour pressure that's dependent on temperature.

Imagine spraying perfume in a sealed room. Initially, the scent is strong, but over time it disperses evenly throughout the room; the rate of scent particles entering and leaving the air becomes constant, similar to how water vapour and liquid water reach equilibrium.

Let us now consider the systems where solids sublime to vapour phase. If we place solid iodine in a closed vessel, after sometime the vessel gets filled up with violet vapour and the intensity of colour increases with time.

When solid iodine is placed in a closed container, it slowly sublimates into iodine vapour, leading to an increase in the vapour's concentration until a dynamic equilibrium is established between the solid and vapour phases.

Think of sugar sitting at the bottom of a warm cup of tea. Initially, it dissolves quickly, but as saturation occurs, there's no more visible sugar left; it reaches a point where the sugar dissolves and recrystallizes at a balancing rate.

Solids in liquids: We know from our experience that we can dissolve only a limited amount of salt or sugar in a given amount of water at room temperature.

The concept of saturation occurs when no more solute can dissolve in the solvent. In a saturated solution, the rate of dissolution equals the rate at which the dissolved particles precipitate back out, maintaining a dynamic equilibrium.

Imagine making syrup. You can keep adding sugar to water until it no longer dissolves, forming a thick layer at the bottom. The amount of dissolved sugar balances with the undissolved sugar; this is a saturated solution.

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