VCE Year 11 Chemistry practice exams with worked answers

Everything in chemistry starts with the atom. An atom has a tiny, dense nucleus containing protons (charge +1, mass about 1 u) and neutrons (no charge, mass about 1 u). Around the nucleus move electrons, which carry a charge of −1 and have almost no mass (about 1/1836 of a proton).

Because almost all of the mass sits in the nucleus, the mass of an atom is set by its protons and neutrons, while its chemical behaviour is set by its electrons.

An element is a substance made of atoms that all have the same number of protons. That number is the atomic number, Z, and it is the identity card of the element: every atom with Z = 6 is carbon, no matter what else is true about it. The mass number, A, is the total number of protons plus neutrons.

Standard notation writes the mass number as a superscript and the atomic number as a subscript to the left of the symbol, for example ²³Na with 11 written below, often shortened to sodium-23 or Na-23. The number of neutrons is simply A − Z.

Isotopes are atoms of the same element (same Z) with different numbers of neutrons, and therefore different mass numbers. Carbon-12, carbon-13 and carbon-14 all have 6 protons but 6, 7 and 8 neutrons respectively. Isotopes of an element have essentially identical chemical properties because they have the same electron arrangement, but they differ in mass and sometimes in nuclear stability (carbon-14 is radioactive).

An ion forms when an atom gains or loses electrons. Losing electrons gives a positive ion (cation); gaining electrons gives a negative ion (anion). The number of protons never changes when an ion forms. For an ion, electrons = Z − charge, so Mg²⁺ has 12 − 2 = 10 electrons and S²⁻ has 16 + 2 = 18 electrons.

Electrons occupy shells (n = 1, 2, 3, 4…). In the simple shell model, the maximum number of electrons in shell n is 2n², giving 2, 8, 18, 32. The more detailed subshell model divides shells into s, p, d and f subshells that hold 2, 6, 10 and 14 electrons. Subshells fill in order of increasing energy: 1s 2s 2p 3s 3p 4s 3d 4p. Note that 4s fills before 3d. Chromium and copper are exceptions (3d⁵4s¹ and 3d¹⁰4s¹).

When transition metals form ions, the 4s electrons are removed first.