Liquids, solids, and phase changes
Matter changes between solid, liquid, and gas as energy is added or removed. Phase diagrams map exactly when each state is stable.
Properties of liquids
Liquids have a fixed volume but take the shape of their container. Their molecules are close together but can slide past one another. Intermolecular forces determine viscosity, surface tension, and how easily the liquid evaporates.
Types of solids
Crystalline solids have ordered, repeating structures: ionic solids (salt), molecular solids (ice, sugar), covalent network solids (diamond, quartz), and metallic solids. Amorphous solids like glass lack long-range order. Diamond's network of covalent bonds makes it extremely hard and high-melting.
Phase changes
Melting, vaporisation, and sublimation absorb energy; freezing, condensation, and deposition release it. During a phase change, temperature stays constant: added heat breaks intermolecular attractions rather than speeding molecules up. That is why a pot of boiling water stays at \(100^\circ\text{C}\) at sea level.
Heating curves and latent heat
A heating curve plots temperature against heat added. Sloped segments warm a single phase (\(q = mc\Delta T\)); flat segments are phase changes (\(q = n\Delta H\)). Water's heat of vaporisation (about 40.7 kJ/mol) is much larger than its heat of fusion (about 6.0 kJ/mol), because boiling separates molecules completely.
q = mc\Delta TVapour pressure and boiling
Molecules escaping a liquid's surface create vapour pressure, which rises with temperature. A liquid boils when its vapour pressure equals the surrounding pressure. On a high mountain, lower air pressure means water boils below \(100^\circ\text{C}\), so food cooks more slowly; pressure cookers do the reverse.
Phase diagrams
A phase diagram shows which phase is stable at each pressure and temperature. The triple point is where solid, liquid, and gas coexist; the critical point marks where liquid and gas become indistinguishable. Water's solid-liquid line slopes backward, so pressure can melt ice, which is unusual.
Worked example
How much heat melts 36.0 g of ice at \(0^\circ\text{C}\)? That is 2.00 mol, so \(q = 2.00 \times 6.01 = 12.0\) kJ. Warming the resulting water to \(100^\circ\text{C}\) takes \(q = 36.0 \times 4.18 \times 100 \approx 15.0\) kJ, and boiling it takes about 81 kJ, the largest step by far.
Sublimation
Some solids turn directly into gas without melting, such as dry ice (solid \(\text{CO}_2\)) at atmospheric pressure. Freeze-drying food uses sublimation of ice under low pressure to remove water while preserving structure and flavour.
Key takeaways
Solids may be ionic, molecular, network, metallic, or amorphous. Phase changes absorb or release latent heat at constant temperature. Heating curves combine \(q = mc\Delta T\) and \(q = n\Delta H\). Boiling occurs when vapour pressure equals external pressure. Phase diagrams show stable phases, the triple point, and the critical point.
Flashcard
Ordered repeating structure
Check your understanding
Why does water boil at a lower temperature on a high mountain?