4. States of matter

Gas pressure is produced by gas particles colliding with the walls of their container. Each collision transfers momentum to the wall; the combined effect of many collisions over an area gives the measured pressure.

At constant temperature, reducing the container volume packs the particles into a smaller space. Wall collisions become more frequent, so pressure increases: pressure is inversely related to volume under this condition.

At constant volume, increasing temperature raises the particles' average kinetic energy. They move faster and collide with the walls more frequently and with greater effect, so pressure increases: pressure is directly related to temperature under this condition.

State the controlled variable with every relationship: the volume–pressure relationship requires constant temperature, while the temperature–pressure relationship requires constant volume. Do not treat either proportionality as an unrestricted rule when the condition changes.

An ideal gas is a model whose particles move rapidly and randomly, occupy negligible volume, exert no intermolecular attraction or repulsion, and collide elastically. Its temperature is linked to the particles' average kinetic energy.

Because the particles are treated as point-like and non-attracting, the model predicts gas behaviour using only measurable pressure, volume, temperature and amount. At constant pressure, heating gives the particles more kinetic energy and the volume increases so the collision effect on the walls remains consistent.

Real gases approach the model under suitable conditions but deviate at very high pressure or low temperature. Particles are then closer: attractions can pull particles inward and lower the measured wall pressure, while the particles' own volume reduces the free space available for movement.

Do not treat ‘ideal’ as the description of every real gas or forget which assumptions fail. Keep this card on model assumptions and their limits; the numerical use of pV = nRT, unit conversion and Mr calculations belong to the neighbouring quantitative objective.

Use the ideal-gas equation pV = nRT to connect pressure, volume, amount of gas and temperature. Identify the unknown and keep pressure, volume, amount and temperature consistent with the units required by the chosen gas constant.

Before substituting, convert pressure to Pa, volume to m³ and temperature to kelvin; keep the amount in mol. A value in kPa, cm³ or dm³ is not ready for direct use with the SI value of R, and Celsius must not be used as an absolute temperature.

Use an ordered calculation: write pV = nRT, rearrange for the unknown, audit units, substitute, calculate, and convert the final output only when needed. For molar mass, first find the amount from n = pV/RT, then divide mass by amount in moles.

Do not mix pressure, volume or temperature units, confuse mass with amount of substance, or use the wrong amount of gas. Keep this card on quantitative use of the ideal-gas equation; deviations from ideal behaviour belong to the neighbouring model-and-limitations objective.

A crystalline lattice is a regular, repeating arrangement of particles. Classify the lattice by asking what repeats: ions in a giant ionic lattice, atoms in a giant covalent or metallic lattice, or discrete molecules in a simple molecular lattice.

In a giant ionic lattice, oppositely charged ions alternate through the structure and are held by electrostatic attraction in all directions. The arrangement and relative ion sizes determine the packing; sodium chloride and magnesium oxide are source-supported examples of cubic ionic lattices.

Covalent substances may form simple molecular lattices, such as iodine, buckminsterfullerene and ice, where discrete molecules are arranged together, or giant molecular lattices, such as diamond, graphite and silicon(IV) oxide, where covalent bonding extends through a large network. Do not treat all covalent substances as one lattice type.

A giant metallic lattice contains positive metal ions arranged regularly and surrounded by a sea of delocalised electrons. The ions may form layered or cubic arrangements; copper is a source-supported example. The delocalised electrons are part of the metallic structure, not separate molecules between the ions.

Giant ionic substances have high melting and boiling points because strong electrostatic attractions hold oppositely charged ions throughout the lattice. They are brittle because shifting layers can bring like charges together and split the crystal. Ions are fixed in a solid but mobile when molten or dissolved, so conductivity depends on state; ion–dipole interactions can support solubility in water.

Giant metallic substances are strong and generally have high melting and boiling points because positive metal ions attract delocalised electrons. Mobile electrons allow conduction in both solid and liquid states. Layers can slide while the metallic attraction is maintained, giving malleability; more delocalised electrons and smaller cations strengthen the attraction in the source-supported comparison.

Simple molecular substances generally have low melting and boiling points because only weak intermolecular forces need to be overcome between molecules; they usually do not conduct because they lack mobile charged particles. Polarity or hydrogen bonding can change solubility and attraction strength. The covalent bonds inside each molecule are not broken during melting or boiling.

Giant covalent substances have high melting and boiling points because many strong covalent bonds extend through the network. Diamond and silicon(IV) oxide are hard and lack mobile charge carriers, whereas graphite is soft because its layers attract weakly and can slide, and it conducts along the layers because it has delocalised electrons. Use these as structure-based exceptions, not as a rule that all carbon forms behave alike.

Explain a property by naming the structure, the relevant attraction or mobile particle, and the condition being tested. Do not infer ionic structure from melting point alone, say that every ionic solid conducts, or treat ‘insoluble’ and ‘non-conducting’ as interchangeable evidence.

Start by recording the conditions of the evidence: melting or boiling behaviour, electrical conductivity in the stated state, solubility in water, hardness or malleability, and any physical state. Do not compare a solid conductivity result with a molten or aqueous result as if they were the same test.

Use the evidence to narrow the structure: high melting behaviour suggests strong attractions or an extended network; mobile charge in a solid supports a metallic structure, while mobile charge only when molten or in solution supports an ionic structure. Low melting behaviour and weak conductivity support a simple molecular structure, subject to polarity and hydrogen-bonding exceptions.

Check the remaining observations against the candidate. Distinguish giant covalent materials from simple molecular substances using the extent of covalent bonding and the resulting melting behaviour; use graphite’s delocalised electrons and layered softness, and diamond or silicon(IV) oxide’s hard network, as source-supported contrasts.

State the best-supported structure and the linked causal evidence: particle type and bonding/attraction → mobility or energy needed → observed properties. A single property is rarely decisive; if evidence is incomplete or an exception applies, state the limitation rather than forcing a classification.

Do not identify a structure from melting point alone, call every high-melting substance ionic, or treat insolubility as proof of one structure. Keep the deduction on the four syllabus structure classes and their source-supported property signatures; detailed reaction or energetics analysis belongs elsewhere.