Atomic Structure

An atom contains a tiny, positively charged nucleus surrounded by electrons. Protons and neutrons supply almost all its mass. Electrons occupy the much larger surrounding region and provide negative charge. A neutral atom has equal numbers of protons and electrons; neutral does not mean that it contains no charged particles.

In ^A_ZX, the proton number Z is the number of protons and the nucleon number A is 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. Gaining or losing electrons does not change A or Z.

Isotopes are atoms of the same element with the same proton number but different neutron numbers. Chlorine-35 and chlorine-37 both have 17 protons, but have 18 and 20 neutrons. Their neutral atoms have the same electron configuration, so their chemical properties are very similar; their masses differ. Do not confuse an isotope difference with an ionisation change.

The electric force is F = qE. A neutron has no net charge and is not deflected by the uniform electric field. A proton and electron have equal charge magnitudes, so the force magnitudes are equal but opposite. The electron has much less mass, giving much greater acceleration under that force.

For particles entering at the same speed, the time inside the field is the same. Deflection is then proportional to charge-to-mass ratio, |q|/m. If speeds or initial kinetic energies differ, that comparison alone is insufficient. Read the beam conditions before choosing the greatest deflection.

The principal quantum number n labels a shell. Within a shell, subshells are labelled s, p and d. An orbital describes a region where an electron is likely to be found; it is not a circular track. Each orbital can hold at most two electrons with opposite spins.

Orbitals within the same subshell have equal energy in an isolated atom. In many-electron atoms, penetration and shielding split the energies of different subshells. This differs from hydrogen, where orbitals with the same n have the same energy in the simple atomic model.

A drawing of a two-lobed p orbital is one orbital, not two orbitals. Similarly, the four lobes of most d orbitals belong to one orbital. An orbital shape is a probability description; it is not the shape of an electron. Mathematical wave functions are outside this syllabus.

Count electrons first. Fill lower-energy orbitals before higher-energy orbitals. Two electrons in one orbital have opposite spins (Pauli exclusion). In a set of equal-energy orbitals, put one electron into each with parallel spins before pairing (Hund's rule). Thus the three 2p electrons of nitrogen occupy three different p orbitals, rather than pairing in one immediately.

Superscripts count electrons: 3d⁶ means six electrons distributed over five d orbitals. [Ar] abbreviates the 18-electron argon core. Chromium and copper have the ground-state exceptions shown above; a rigid filling diagram does not capture every balance of orbital energies and electron repulsions.

First ionisation energy is the energy needed to remove one mole of electrons from one mole of gaseous atoms, forming one mole of gaseous singly positive ions: X(g) → X⁺(g) + e^-. It is positive because energy is needed to separate an electron from the attraction of the nucleus. The gas state prevents lattice or intermolecular effects being included.

Across a period, proton number rises while added electrons enter the same principal shell. Shielding does not increase enough to cancel the increased nuclear charge, so the general trend is increasing ionisation energy. Down a group, extra shells increase distance and shielding, usually outweighing the increased nuclear charge.

Successive ionisation energies remove electrons one at a time from increasingly positive gaseous species. The second process is X⁺(g) → X²⁺(g) + e^-. Values rise because the remaining electrons are attracted more strongly to the increasingly positive ion. A particularly large jump means the next electron is removed from a more tightly held inner shell.

For a main-group atom, one outer electron suggests group 1; two suggest group 2; three to eight suggest groups 13 to 18. Helium is the familiar two-electron exception in group 18. Do not apply this simple outer-count rule uncritically to transition elements, where 3d and 4s energies and occupations also matter.

Keep the distinctions: isotope versus ion; orbital versus subshell; electron count versus proton number; filling order versus removal order; general periodic trend versus a subshell or pairing exception.