Everything from the 6 lessons of this topic on a few pages: the key idea of each lesson, the diagram to remember, the points to learn and the traps examiners set. Read it, cover it, test yourself.
Protons (relative charge +1, relative mass 1) and neutrons (0, 1) sit in the nucleus, so the nucleus carries nearly all the mass and all the positive charge. Electrons (−1, 1/1836) fill the space around it. Proton number Z counts protons; nucleon number A counts protons plus neutrons. Neutrons = A − Z, and electrons = Z − (charge on the ion).
The charge on an ion changes the electron count only. Protons and neutrons are untouched: an ion is still the same element with the same nucleus. And the nucleon number on a symbol is a whole-number count for one atom; it is not the relative atomic mass in the Periodic Table, which is an average over isotopes.
Across a period the radius decreases: each step adds one proton and one electron to the same shell, shielding barely changes, so the outer electrons are pulled in harder. Down a group the radius increases: each period adds a whole new shell, further out and more shielded. A cation is smaller than its atom (the outer shell is lost), an anion is larger (extra repulsion), and in an isoelectronic series the ion with the most protons is the smallest.
Ionic radius across Period 3 is not a smooth decrease. It falls from Na+ to Mg2+ to Al3+ (all 10 electrons, n = 2 outer shell), then jumps up at P3−, S2− and Cl−, which have 18 electrons and an n = 3 outer shell, and then falls again along that second isoelectronic set.
Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons, so the same proton number Z and a different nucleon number A. They have the same electron configuration and therefore the same chemical properties. Physical properties that depend on mass, such as density and rate of diffusion, differ slightly. Relative atomic mass is the weighted mean of the isotopic masses, which is why it is rarely a whole number.
Isotopes have the same number of protons and electrons; only the neutron number differs. Saying that isotopes have a different atomic number, or a different number of electrons, loses the mark. And the mass number is a property of one atom; the relative atomic mass is a property of the natural mixture.
Shell n holds s, p, d and f sub-shells up to n of them; an s sub-shell has one orbital, p has three, d has five, and each orbital holds two electrons with opposite spins. Sub-shells fill in order of increasing energy, 1s 2s 2p 3s 3p 4s 3d 4p, so 4s fills before 3d. Within a sub-shell, electrons occupy orbitals singly before pairing. Chromium and copper break the pattern to give half-filled and full 3d sub-shells.
4s is filled before 3d and emptied before 3d. Writing Fe2+ as [Ar] 3d4 4s2 is the classic error. And chromium is 3d5 4s1, not 3d4 4s2; copper is 3d10 4s1, not 3d9 4s2.
First ionisation energy is the energy needed to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions: X(g) → X+(g) + e−. It is larger when the nuclear charge is greater, the outer electron is closer, and shielding is lower. Across Period 3 it generally rises, with dips at aluminium (3p sits above 3s) and sulfur (spin-pair repulsion in 3p4). Down a group it falls because the outer electron is further away and more shielded.
Three things lose marks in a definition: forgetting gaseous, forgetting one mole, and writing a 2+ ion for a first ionisation. In explanations, say shielding or outer electron further from the nucleus; "more electrons" on its own is not an explanation.
Successive ionisation energies always increase: each electron leaves an ion that is more positive, so the same nuclear charge holds fewer electrons more tightly and the ion shrinks. A large jump appears when the next electron comes from an inner shell, closer to the nucleus and less shielded. The number of electrons removed before the first large jump equals the number of outer-shell electrons, which is the group number. Electron affinity, X(g) + e− → X−(g), is exothermic for the first electron but endothermic for a second, because a negative ion repels the incoming electron.
Count the electrons removed before the jump, not the position of the jump itself. A large increase between the 2nd and 3rd ionisation energies means two outer electrons, Group 2. And successive ionisation energies are always positive and always increase; a smaller number than the one before is a sign you have misread the table.