Atoms, elements and compounds, the development of the atomic model, the structure of the atom and electronic structure, isotopes and relative atomic mass, the periodic table and its development, and the properties of metals, non-metals, Group 1, Group 7 and Group 0.

What this topic is asking

AQA wants you to define atoms, elements and compounds, describe how the atomic model developed, give the structure of the atom and electronic structures, explain isotopes and relative atomic mass, and describe the periodic table and the properties of Groups 1, 7 and 0.

Atoms, elements and compounds

The atomic model developed as new evidence appeared. John Dalton pictured atoms as tiny solid spheres. J. J. Thomson discovered the electron and proposed the plum pudding model (a ball of positive charge with electrons embedded in it). The alpha-scattering experiment, in which most alpha particles passed straight through gold foil but a few were deflected back, led Rutherford to the nuclear model: a small, dense, positive nucleus surrounded by mostly empty space. Niels Bohr then showed that electrons occupy fixed energy levels (shells) at set distances, and later work identified the proton and (by James Chadwick) the neutron. This sequence is a classic example of how a scientific model changes as experimental evidence accumulates.

Structure of the atom

Atoms have a radius of about $1 \times 10^{-10}$ m, and the nucleus is around 10000 times smaller again, so the atom is mostly empty space. The nucleus contains protons (relative charge +1, relative mass 1) and neutrons (charge 0, mass 1); electrons (charge -1, almost no mass) occupy shells around it.

Isotopes are atoms of the same element (same number of protons) with different numbers of neutrons, so they have the same atomic number but different mass numbers, for example carbon-12 and carbon-14. Because they have the same electron arrangement, isotopes react chemically in the same way. The relative atomic mass of an element is the mean mass of its atoms, weighted by the abundance of each isotope, which is why values such as 35.5 for chlorine are not whole numbers.

The periodic table and the groups

Dmitri Mendeleev arranged the known elements in order of atomic weight but left gaps and even swapped some so that elements with similar properties lined up; the gaps correctly predicted undiscovered elements. The modern periodic table lists elements in order of atomic number. Periods are the rows; groups are the columns, and elements in the same group have the same number of outer-shell electrons, which is why they have similar chemical properties. Metals (most of the table, on the left and centre) react by losing electrons to form positive ions; non-metals (upper right) tend to gain or share electrons.

• Group 1 (alkali metals): soft, reactive metals with one outer electron. They react with water to form an alkaline hydroxide and hydrogen, and become more reactive down the group because the outer electron is further from the nucleus, more shielded by inner shells, and so lost more easily.
• Group 7 (halogens): reactive non-metals with seven outer electrons. They become less reactive down the group because the outer shell is further from the nucleus, so an electron is gained less easily. A more reactive halogen will displace a less reactive one from a solution of its salt.
• Group 0 (noble gases): unreactive (inert) because they have full outer shells (8 electrons, except helium with 2), so they have no tendency to gain, lose or share electrons. Their boiling points increase down the group.