Evaporation: Surface phenomenon, occurs at any temperature, slow, causes cooling, molecules need sufficient kinetic energy to escape. Boiling: Bulk phenomenon, occurs at a fixed temperature (boiling point) at given pressure, rapid, formation of vapour bubbles throughout liquid, requires continuous energy input. Evaporation rate increases with wind, surface area, and temperature; boiling rate depends on heat supply.
Boiling point is the temperature at which vapour pressure equals external pressure. Higher external pressure → molecules need more kinetic energy to match it → boiling point increases. Lower pressure (e.g., mountains) → boiling occurs at lower temperature. This is why pressure cookers increase pressure to raise boiling point and cook faster; at high altitudes, water boils below 100°C, requiring longer cooking times.
Evaporation uses the internal energy of the liquid itself. The molecules with higher kinetic energy overcome attractive forces and escape from the surface. The average kinetic energy of remaining liquid decreases → cooling effect. Energy is absorbed from surroundings (or liquid), but no external heater is required. That’s why sweating cools the body: latent heat taken from skin.
Phase changes are reversible and involve the same latent heat. During heating, energy breaks intermolecular bonds (melting, vaporization). During cooling, exactly the same amount of energy is released when bonds form (freezing, condensation). Thus the plateaus occur at identical temperatures (e.g., 0°C and 100°C for water), assuming equilibrium and no superheating/cooling. The path is essentially mirrored.
Ice (solid water) has an unusual property: its density is lower than liquid water. Applying pressure favours the denser phase (water) → ice melts at temperature below 0°C. Once pressure is released, water may refreeze. This is called regelation, demonstrated by a wire slicing through an ice block without cutting it. Most substances contract on freezing, but water expands, making ice anomalous.
Phase changes occur when the kinetic energy (temperature) of particles crosses a threshold relative to intermolecular potential energy. Adding heat increases particle motion until they overcome attractive forces (solid→liquid→gas). Removing heat decreases motion, allowing forces to pull particles into ordered arrangements (gas→liquid→solid). Essentially phase changes balance between energy and intermolecular forces.
According to kinetic molecular theory, gas particles are in constant random motion, and the average kinetic energy (KE) depends only on temperature: \( KE = \frac{1}{2} m v^2 \). At the same temperature, lighter gases (lower molar mass) have higher root-mean-square speed: \( v_{rms} = \sqrt{\frac{3RT}{M}} \).
Diffusion rate ∝ average speed ∝ \( \frac{1}{\sqrt{M}} \) (Graham’s law). For example, hydrogen (M=2) diffuses about 4× faster than oxygen (M=32). Collision frequency also depends on size, but mass dominates. Thus, in Q7, H₂S (34 g/mol) diffuses faster than CO₂ (44) and SO₂ (64).
🖨️ Dye‑sublimation printing: Solid dye particles are heated until they sublimate directly into gas without passing through liquid. The gaseous dye penetrates the polymer coating of paper or fabric and then solidifies (deposition), creating vibrant, permanent, high‑resolution images. Used for printing on mugs, T‑shirts, and banners.
🌸 Air fresheners: Solid gel or crystalline fresheners contain fragrant organic compounds that sublime slowly at room temperature. The solid turns into vapour, releasing pleasant smell over weeks. No liquid phase means no mess. Naphthalene and paradichlorobenzene work similarly (mothballs). Sublimation provides controlled, gradual release.
Plateaus: At 0°C (fusion) and 100°C (vaporization). Slopes represent temperature increase in single phase.
Explanation: Segment A→B: ice warms. B→C: ice melts at constant 0°C (latent heat of fusion). C→D: liquid water warms. D→E: water boils at constant 100°C (latent heat of vaporization). E→F: steam warms. The curve shape is universal for pure substances.