Atomic Structure

An atom has a very small, positively charged nucleus containing protons and neutrons. Almost all its mass is concentrated there. Electrons occupy the surrounding space, which accounts for almost all the volume. A neutral atom has equal numbers of protons and electrons; an ion does not.

The proton number, Z, is the number of protons and defines the element. The nucleon number, A, counts protons plus neutrons. Therefore neutrons = A - Z. For an ion with signed charge q, electrons = Z - q: subtract a positive charge, but add the magnitude of a negative charge. Ion formation transfers electrons; it does not change A or Z.

Isotopes have the same proton number but different neutron numbers. Chlorine-35 and chlorine-37 both have 17 protons, but have 18 and 20 neutrons respectively. Neutral atoms of both have 17 electrons, so their electron arrangements and chemical behaviour are very similar. Their masses differ; isotopes are not different ions.

Opposite charges attract. A proton beam deflects towards the negative plate; an electron beam deflects towards the positive plate. Neutrons have no charge and are not deflected by the electric field. In the same field, force depends on charge and acceleration depends on force divided by mass.

A shell is labelled by principal quantum number n. Within a shell, subshells contain orbitals: one s orbital, three p orbitals or five d orbitals. Each orbital holds at most two electrons with opposite spins. Orbitals in the same subshell have the same energy in an isolated atom unless an external influence splits them.

Within a multi-electron shell, s is lower in energy than p, which is lower than d where present. The useful filling sequence for the atoms studied is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p. Energies also depend on occupancy: this filling order is not a claim that 4s is always lower than 3d in every ion.

The symbol [Ar] replaces the filled first 18 electrons. For Cr (Z = 24), use [Ar] 3d⁵ 4s¹; for Cu (Z = 29), use [Ar] 3d¹⁰ 4s¹. These are exceptions to simply assigning two electrons to 4s before 3d. When forming their positive ions, remove the 4s electron first.

The first ionisation energy is the energy needed to remove one electron from each atom in one mole of gaseous atoms, forming one mole of gaseous 1+ ions. For X: X(g) → X⁺(g) + e^-. Values are positive and usually quoted in kJ mol^-1: energy is supplied to overcome attraction to the nucleus.

Across Period 3, first ionisation energy generally increases: nuclear charge rises while added electrons enter the same principal shell, so shielding changes relatively little. Two dips need more precise reasoning. Al loses a higher-energy 3p electron, whereas Mg loses 3s; Al is easier to ionise. S loses an electron from a paired 3p orbital, whereas P has three singly occupied 3p orbitals; pair repulsion makes the S electron easier to remove.

Down a group, the outer electron occupies a higher shell farther from the nucleus and is more shielded. These effects usually outweigh the greater nuclear charge, so first ionisation energy decreases. Say which effects dominate; merely saying "there are more electrons" does not explain the attraction.

Successive ionisation energies refer to removing electrons one at a time from increasingly positive gaseous ions. The second process is X⁺(g) → X²⁺(g) + e^-. Values rise because the remaining electrons are held more strongly. A particularly large jump indicates that the next electron comes from a lower, inner shell much closer to the nucleus.