States of Matter: Solid, Liquid, Gas

In a solid, particles (atoms, ions, or molecules) are packed closely together and held in fixed positions by strong intermolecular or interionic forces. The particles vibrate in place but do not translate or flow past one another. This gives solids two defining macroscopic properties: a fixed shape and a fixed volume. Crystalline solids such as sodium chloride ($\ce{NaCl}$) have a highly ordered, repeating lattice structure, whereas amorphous solids such as glass lack long-range order. Solids are essentially incompressible because the particles are already as close together as repulsive forces allow. Density is high relative to the corresponding gas phase — ice ($\ce{H2O(s)}$) has a density of approximately 0.92 g cm$^{-3}$, while water vapour at 100 °C is roughly 1600 times less dense.

In a liquid, particles are still close together and in contact (giving a fixed volume comparable to the solid), but they have enough kinetic energy to break free from fixed lattice positions and move past one another. This gives liquids fluidity — they flow and take the shape of their container — while retaining a definite volume. Surface tension arises because molecules at the liquid surface experience a net inward attraction, since they are only surrounded by neighbours on one side. Viscosity describes the resistance to flow and depends on the strength of intermolecular forces; for example, glycerol ($\ce{C3H8O3}$) is far more viscous than water because of its three hydroxyl groups forming an extensive hydrogen-bond network. Liquids are nearly incompressible: water at 1000 atm compresses by only about 4%.

In a gas, particles are separated by large distances — on average about ten molecular diameters for air at room temperature and pressure — and move rapidly in random directions. Intermolecular forces are negligible compared with kinetic energy, so gas particles interact only during brief collisions. Gases have no fixed shape and no fixed volume; they expand to fill any container. They are highly compressible because the space between particles can be reduced by applying pressure. The pressure exerted by a gas arises from the collective force of particles colliding with container walls. Real gases deviate from ideal behaviour at high pressure and low temperature when intermolecular attractions become significant, as described by the van der Waals equation: $\left(P + \frac{an^2}{V^2}\right)(V - nb) = nRT$, where $a$ accounts for attractions and $b$ for the finite volume of particles.

The state adopted by a substance at a given temperature and pressure depends critically on the type and strength of its intermolecular forces. London dispersion forces (present in all molecules, stronger with larger electron clouds), dipole–dipole forces (between polar molecules), and hydrogen bonds (between molecules with N–H, O–H, or F–H bonds) all act to hold particles together. Water ($\ce{H2O}$) is liquid at room temperature despite its low molar mass (18 g mol$^{-1}$) because of strong hydrogen bonding (~20 kJ mol$^{-1}$ per bond). In contrast, methane ($\ce{CH4}$, molar mass 16 g mol$^{-1}$) has only weak London forces and is a gas at room temperature, with a boiling point of −161.5 °C. Understanding these forces allows chemists to predict and engineer the physical state of materials.

Transitions between states involve the transfer of thermal energy. Melting (solid → liquid) and vaporisation (liquid → gas) are endothermic — energy must be supplied to overcome intermolecular attractions. Freezing and condensation are the reverse exothermic processes. Sublimation is the direct solid-to-gas transition, seen in dry ice ($\ce{CO2(s)}$) at atmospheric pressure. The enthalpy of vaporisation ($\Delta H_{\text{vap}}$) is always larger than the enthalpy of fusion ($\Delta H_{\text{fus}}$) for the same substance, because vaporisation requires completely separating particles from liquid-range contact distances to gas-phase separation, whereas melting only disrupts the ordered lattice. For water: $\Delta H_{\text{fus}} = 6.01$ kJ mol$^{-1}$ and $\Delta H_{\text{vap}} = 40.7$ kJ mol$^{-1}$.

• P — Pressure (Pa (pascals) or atm): Force exerted per unit area by gas particles colliding with container walls.
• V — Volume (m³ or L): The space occupied by the gas; gases expand to fill available volume.
• n — Amount of substance (mol): Number of moles of gas particles present in the sample.
• R — Universal gas constant (8.314 J mol⁻¹ K⁻¹): Proportionality constant linking energy scale to temperature and moles.
• T — Absolute temperature (K (kelvin)): Must be in kelvin; proportional to average kinetic energy of gas particles.

Problem. A sample of nitrogen gas ($\ce{N2}$) occupies a volume of 5.00 L at a temperature of 27 °C and a pressure of 2.00 atm. (a) Calculate the number of moles of $\ce{N2}$ present. (b) If the temperature is raised to 127 °C at constant volume, what is the new pressure?

Answer. n = 0.406 mol; new pressure P₂ = 2.67 atm

• Kinetic Molecular Theory — Provides the molecular-level explanation for why gases, liquids, and solids behave differently — linking particle speed and energy to macroscopic properties such as pressure, temperature, and volume.
• Intermolecular Forces (London Dispersion, Dipole–Dipole, Hydrogen Bonding) — The type and strength of intermolecular forces determine which state a substance adopts at a given temperature, and govern properties like boiling point, viscosity, and surface tension.
• Phase Changes and Enthalpy (Heating Curves) — Directly extends states-of-matter knowledge by quantifying the energy absorbed or released during transitions between solid, liquid, and gas using enthalpy of fusion and vaporisation.
• Ideal Gas Law (PV = nRT) — The ideal gas law mathematically describes the behaviour of the gaseous state, connecting pressure, volume, temperature, and amount of substance in a single equation.
• Density and Physical Properties of Matter — Density differences between states (solids and liquids are far denser than gases) follow directly from particle spacing and packing, reinforcing the structural picture of each state.