Topic 1: Structure 1: Models of the particulate nature of matter

Chemistry begins by treating matter as countless tiny particles whose behaviour explains everything from melting ice to coloured flames. This theme builds the toolkit you need for the whole course: counting particles with the mole, describing the atom and its electrons, and reading the evidence that spectroscopy and gas behaviour provide. Each model is judged by how well it predicts what we actually observe.

States of matter and the particulate model

All matter is made of particles in constant motion, and the three states differ in how those particles are arranged and how fast they move. In a solid the particles are packed closely in fixed positions and only vibrate, giving a fixed shape and volume. In a liquid the particles still touch but can slide past one another, so a liquid flows and takes the shape of its container while keeping a fixed volume. In a gas the particles are far apart, move rapidly and randomly, and fill any container. Heating supplies kinetic energy, so the particles move faster; at the melting point and boiling point this energy goes into overcoming the forces holding particles together rather than raising temperature. Changes of state (melting, freezing, vaporization, condensation, sublimation) are physical changes because the particles themselves are unchanged.

The mole, molar mass and Avogadro's constant

Because atoms are far too small and numerous to count individually, chemists group them into the mole. One mole is the amount of substance that contains Avogadro's constant of particles, where Avogadro's constant L is 6.02 times 10 to the power 23 per mole. The number of particles equals the number of moles multiplied by L. Molar mass M is the mass of one mole of a substance in grams per mole, and it is numerically equal to the relative atomic or relative molecular mass. The central relationship is n equals m divided by M, where n is the amount in moles and m is the mass in grams. The mole acts as a bridge between the mass you can weigh on a balance and the number of particles that actually react.

Empirical and molecular formulas

An empirical formula gives the simplest whole-number ratio of atoms in a compound, while a molecular formula gives the actual number of each atom in one molecule. To find an empirical formula, convert the mass or percentage of each element to moles by dividing by its molar mass, then divide every result by the smallest value to obtain the ratio; multiply up if needed to reach whole numbers. The molecular formula is a whole-number multiple of the empirical formula, found by dividing the relative molecular mass by the empirical formula mass. For example, both ethyne C2H2 and benzene C6H6 share the empirical formula CH, so empirical data alone cannot identify a compound without its molar mass.

The nuclear atom and isotopes

The current model places almost all of an atom's mass in a tiny central nucleus made of positively charged protons and neutral neutrons, surrounded by negatively charged electrons. The atomic number Z equals the number of protons and defines which element an atom is, while the mass number A equals the number of protons plus neutrons. A neutral atom has equal numbers of protons and electrons. Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons, so they share chemical properties but differ in mass. Ions form when atoms gain or lose electrons, becoming negative anions or positive cations respectively while keeping their proton count unchanged.

Relative atomic mass and mass spectra

Relative atomic mass Ar is the weighted average mass of an element's atoms compared with one twelfth of the mass of a carbon-12 atom, taking the natural abundance of each isotope into account. A mass spectrometer measures this directly: a sample is ionized, the ions are accelerated and deflected by a magnetic field according to their mass-to-charge ratio, and a detector records the relative abundance of each ion. The resulting spectrum shows a peak for each isotope. To calculate Ar, multiply each isotope mass by its fractional abundance and add the results. For example, chlorine has isotopes of mass 35 and 37 in roughly a 3 to 1 ratio, giving an Ar of about 35.5.

Electron configurations and orbitals

Electrons occupy regions called orbitals, each holding a maximum of two electrons, grouped into subshells labelled s, p, d and f. An s subshell has one orbital, a p subshell three, and a d subshell five, holding 2, 6 and 10 electrons respectively. Electrons fill the lowest available energy levels first (the Aufbau principle), each orbital gains one electron before any pairs up (Hund's rule), and paired electrons have opposite spins (the Pauli exclusion principle). Configurations are written by listing subshells with their electron counts, for example sodium as 1s2 2s2 2p6 3s1. Note that the 4s subshell fills before 3d because it is slightly lower in energy.

Emission spectra and evidence for energy levels

When atoms absorb energy their electrons jump to higher energy levels, and when these excited electrons fall back they emit photons of light with specific energies. Because only certain energy differences are possible, the emitted light appears as discrete coloured lines rather than a continuous rainbow, producing a line emission spectrum unique to each element. This discreteness is direct evidence that electron energy is quantized into fixed levels. In the hydrogen spectrum the lines converge at higher energy (shorter wavelength) because the energy levels themselves get closer together as they approach the point where the electron escapes the atom. Emission spectra confirm the quantized model of the atom and underpin flame tests and astronomical analysis.

Ideal gases and the ideal gas equation

The kinetic model treats an ideal gas as particles of negligible volume that move randomly, exert no forces on each other, and collide elastically. For such a gas, pressure P, volume V, amount n and temperature T are linked by the ideal gas equation PV equals nRT, where R is the universal gas constant and temperature must be in kelvin. From this follow the simple laws: at fixed temperature pressure is inversely proportional to volume, and at fixed pressure volume is proportional to absolute temperature. To convert Celsius to kelvin, add 273. Real gases deviate from ideal behaviour at high pressure and low temperature, where particle volume and intermolecular forces can no longer be ignored.

(HL) Ionization energy trends

The first ionization energy is the energy needed to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous positive ions. Across a period ionization energy generally rises because increasing nuclear charge pulls electrons in more strongly while the shielding stays similar. Down a group it falls because outer electrons are further from the nucleus and better shielded, so they are easier to remove. Small dips, such as the drop from beryllium to boron and from nitrogen to oxygen, reveal subshell structure: removing an electron from a higher-energy p subshell, or breaking a newly formed electron pair, requires less energy. Successive ionization energies of one atom rise sharply when a full inner shell is broken into, giving direct evidence for shells and the periodic arrangement of electrons.