QCE Year 11 Chemistry practice exams with worked answers

Every atom is built from a tiny, dense, positively charged nucleus surrounded by a much larger region occupied by electrons. The nucleus holds protons (charge +1, mass about 1 u) and neutrons (no charge, mass about 1 u). Electrons carry a charge of −1 and have a mass roughly 1/1836 of a proton, so almost all of an atom's mass sits in the nucleus while almost all of its volume is electron cloud.

The Geiger–Marsden gold foil experiment supplied the key evidence: most alpha particles passed straight through thin gold foil, but a small fraction bounced back sharply, which only makes sense if positive charge and mass are concentrated in a very small centre. Later, Maria Goeppert Mayer's nuclear shell model showed that protons and neutrons are themselves arranged in energy levels, explaining why some nuclei are especially stable.

The atomic number (Z) is the number of protons and defines the element. The mass number (A) is the total number of protons plus neutrons. In nuclear symbol notation the mass number is written as a superscript and the atomic number as a subscript before the symbol, for example ³⁵₁₇Cl; the same isotope can be written chlorine-35 or Cl-35. Neutrons = A − Z.

For a neutral atom, electrons = protons; for an ion, subtract the charge from Z (a 2+ ion has two fewer electrons, a 3− ion has three more).

Isotopes are atoms of the same element with different numbers of neutrons. Because they have identical electron configurations, isotopes react in the same way chemically, but they differ in mass-dependent physical properties such as density, rate of diffusion and nuclear stability. Unstable isotopes (radioisotopes) decay, which is useful: carbon-14 dating estimates the age of once-living material, and cobalt-60 or iodine-131 are used in radiotherapy.

The relative atomic mass (Aᵣ) of an element is the weighted average mass of its naturally occurring isotopes, measured relative to 1/12 of the mass of a carbon-12 atom. A mass spectrometer ionises a sample, separates the ions according to their mass-to-charge ratio and detects them; the resulting spectrum shows a peak for each isotope, and the peak heights give relative abundances.