Nuclear atom

• An atom is the smallest particle of an element that retains that element’s chemical properties. • At its centre is the nucleus — a tiny, dense region containing protons and neutrons (collectively called nucleons). • Electrons move around the nucleus in regions of space called energy levels or orbitals.

• Proton (p⁺): in the nucleus • Neutron (n): in the nucleus • Electron (e⁻): in orbitals around the nucleus Together they form the structure of the atom.

Exam tip: Always give relative values — no units.

Atomic number (Z): the number of protons in the nucleus. It defines the element. Mass number (A): the total number of protons + neutrons (nucleons) in the nucleus.

• Proton: determines the identity of the element (its atomic number). • Neutron: contributes to mass and nuclear stability. • Electron: controls the chemical properties and bonding behaviour.

p=Z, n=A−Z, e=Z−q where q is the ionic charge (+ for cations, – for anions). Example: 40 20Ca2+ : p = 20, n = 20, e = 18.

Ions form when electrons are lost or gained. - Cations (+): electrons lost. - Anions (–): electrons gained. In chemical reactions, only electrons change; protons and neutrons remain constant.

Rutherford’s gold-foil experiment showed that most α-particles passed straight through the foil, while a few were deflected. - This means the nucleus is extremely small and dense, and the rest of the atom is mostly empty space where electrons move.

• Neutrons stabilize the nucleus. They experience the strong nuclear force, which attracts both protons and neutrons together. • Because neutrons have no charge, they help offset proton–proton repulsion without adding extra repulsion. • If there are too few or too many neutrons, the nucleus becomes unstable and may undergo radioactive decay. “Neutrons act as a nuclear glue, providing attraction through the strong nuclear force that balances proton repulsion.”

Isotopes are atoms of the same element (same Z) but with different numbers of neutrons (different A). - They have identical chemical properties (same electron configuration) but different physical properties such as density, diffusion rate, or melting point.

Ar=(35×0.75)+(37×0.25)=35.5

Isotopic Mixtures: Elements exist as mixtures of isotopes, each with different numbers of neutrons. The listed atomic mass is a weighted average of these isotopic masses. Example: Cl = 35.5 (average of ³⁵Cl and ³⁷Cl).

The weighted mean of the masses of the isotopes of an element relative to 1/12 of the mass of a carbon-12 atom.

Rutherford’s α-particle scattering showed that: - Most particles passed straight through → atoms are mostly empty space. - A few deflected strongly → a small, dense, positively charged nucleus exists. This replaced the “plum-pudding” model and led to the nuclear model.

• Electrostatic attraction between the nucleus (+) and electrons (–) keeps electrons bound. • Electrostatic repulsion between protons tends to push them apart. • Strong nuclear force binds nucleons in the nucleus and overcomes proton repulsion at very short distances.

• In a neutral atom, the number of electrons = number of protons, so overall charge is zero. • The number of neutrons can vary (different isotopes) without affecting neutrality.

Na atom: 11 p, 12 n, 11 e → neutral. Na⁺ ion: 11 p, 12 n, 10 e → +1 charge (lost one electron).

An atom is the smallest particle of an element that retains its properties. - It consists of a dense, positively charged nucleus containing protons and neutrons, surrounded by negatively charged electrons in energy levels. - Protons determine the element’s identity, - electrons determine its chemistry, and neutrons stabilise the nucleus by reducing proton–proton repulsion. - Most atomic masses are not whole numbers because natural elements are mixtures of isotopes and because a small mass defect occurs when nucleons bind together.

• Neutrons have no electrical charge but are vital for nuclear stability. • Protons in the nucleus repel each other because they are all positively charged. • Neutrons add attractive strong nuclear forces between all nucleons (p–p, p–n, n–n) that overcome this repulsion at very short distances (≈ 10⁻¹⁵ m). • Because neutrons don’t contribute to electrostatic repulsion, they act as a nuclear glue, holding the nucleus together. A balanced neutron : proton ratio keeps the nucleus stable. - Too few neutrons → proton repulsion dominates → β⁺ decay. - Too many neutrons → excess stability → β⁻ decay. “Neutrons stabilise the nucleus by providing strong-force attraction without adding charge, preventing protons from repelling each other apart.”

Carbon-12 was chosen because it is: - Stable (non-radioactive) and abundant, - Found in pure form, easy to measure precisely, - Provides a convenient integer reference (12 u) that gives almost whole-number relative masses for most other elements. - All relative atomic masses are defined against it: 1 atomic mass unit (u) = 1/12 of the mass of one atom of 12C - Exam phrase: “The carbon-12 isotope was adopted as the universal reference because it is stable, widely available, and allows atomic masses to be measured on a consistent, reproducible scale.”

• A fixed standard ensures that relative masses of all elements are measured consistently worldwide. • It provides a common scale for comparing atomic, molecular, and formula masses and allows data from different experiments and instruments to be directly comparable.

Experiment: - A thin sheet of gold foil was bombarded with α-particles (positively charged helium nuclei). - Most α-particles passed straight through; a few were deflected at small angles, and very few bounced straight back. Observations and Conclusions: Observation Conclusion Most α-particles passed through → Atoms are mostly empty space. Some deflected slightly→ There is a small, dense, positively charged nucleus. A few bounced back → The nucleus contains most of the atom’s mass and positive charge. Impact: This experiment disproved Thomson’s “plum pudding” model and led to the nuclear model, where a dense nucleus is surrounded by orbiting electrons. Exam phrase: “Rutherford’s experiment demonstrated that atoms consist of a tiny, dense, positively charged nucleus containing most of the mass, surrounded by mostly empty space where electrons are found.”

Isotopes are atoms of the same element (same atomic number Z) but with different numbers of neutrons and therefore different mass numbers A. Example: 17 17 35Cl and 37Cl • Both have 17 protons & 17 electrons → same chemical behaviour, but masses = 35 and 37. Exam phrase: “Isotopes are atoms with the same proton number but different numbers of neutrons.”