2.3 Evaporation and Boiling
Evaporation is the process by which molecules at the surface of a liquid gain enough kinetic energy to escape into the gaseous phase. Only molecules at the liquid's surface participate in evaporation. These molecules must have sufficient kinetic energy to overcome intermolecular forces. Energy for evaporation comes from the surrounding environment. This absorption makes surroundings cooler than before. Evaporation can occur at any temperature below the liquid's boiling point. Higher temperatures increase the rate of evaporation since more molecules have sufficient energy to escape.
As high- energy molecules leave the liquid, the average kinetic energy of the remaining molecules decreases, leading to a cooling effect. This is why sweating helps cool the body.
Examples:
Drying of clothes on a line. Evaporation of water from lakes and oceans contributing to the water cycle. Sweat evaporating from the skin, providing a cooling effect.
Boiling is the process in which a liquid turns into vapor when it is heated to its boiling point, resulting in the formation of vapour bubbles within the liquid. Boiling occurs throughout the entire liquid, not just at the surface. Vapour bubbles form within the liquid, rise to the surface, and burst, releasing gas.
Boiling occurs at a specific temperature known as the boiling point. This temperature varies depending on the liquid and the external pressure. For water at sea level, the boiling point is (212°F).
Continuous heat input is required to maintain the liquid at its boiling point. The temperature of the liquid remains constant during boiling, as the added energy is used to convert the liquid into gas. Boiling is characterized by the vigorous formation of bubbles and steam, making it a highly visible process.
The boiling point changes with external pressure. Higher pressure increases the boiling point, while lower pressure decreases it. This is why water boils at lower temperatures at high altitudes.