Atomic Structure Glossary: 60 Key Terms Defined

Atomic structure has a lot of vocabulary, and many terms sound alike: atomic number, atomic mass, mass number, relative atomic mass. This glossary collects the 60 terms students meet most often, grouped by theme so related ideas sit together. Each entry gives a plain-English definition, and many link to a full article if you want to go deeper.

Use it for revision, bookmark it for homework, or test yourself by covering the definitions.

Particles and the nucleus

1. Atom: the smallest particle of an element that keeps the element’s chemical identity. A nucleus surrounded by electrons. See what is an atom?
2. Subatomic particle: any particle smaller than an atom, especially protons, neutrons and electrons.
3. Proton: a positively charged particle in the nucleus, with relative mass 1 and charge +1. See what is a proton?
4. Neutron: a neutral particle in the nucleus with relative mass 1. See what is a neutron?
5. Electron: a negatively charged particle outside the nucleus, with about 1/1836 the mass of a proton. See what is an electron?
6. Nucleus: the tiny, dense centre of an atom containing protons and neutrons, which holds almost all the mass.
7. Nucleon: a proton or neutron, the particles found in the nucleus.
8. Quark: a fundamental particle that makes up protons and neutrons. A proton is two up quarks and one down quark. See quarks and atoms
9. Strong nuclear force: the very short-range attraction between nucleons that holds the nucleus together despite proton repulsion.
10. Positron: the antiparticle of the electron, with the same mass and a positive charge. See antimatter and positrons

Numbers that describe an atom

11. Atomic number (Z): the number of protons in the nucleus. It defines the element: every carbon atom has Z = 6.
12. Mass number (A): the total number of protons plus neutrons. Carbon-14 has A = 14.
13. Neutron number (N): the number of neutrons, N = A − Z.
14. Isotope: atoms of the same element (same Z) with different numbers of neutrons (different A), such as carbon-12 and carbon-14.
15. Isotope notation: writing a nuclide as its symbol with A as a superscript and Z as a subscript on the left, e.g. ¹⁴₆C. See isotope notation
16. Nuclide: a particular kind of nucleus defined by its Z and N.
17. Atomic mass unit (u or Da): one-twelfth of the mass of a carbon-12 atom, about 1.66 × 10⁻²⁷ kg. See the atomic mass unit
18. Relative atomic mass (Ar): the weighted average mass of an element’s atoms, compared with 1/12 of carbon-12. Chlorine’s is about 35.45.
19. Natural abundance: the percentage of an element’s atoms that are a particular isotope in nature. About 75.8% of chlorine atoms are chlorine-35.
20. Mass defect: the difference between the mass of a nucleus and the total mass of its separate nucleons. See mass defect and binding energy
21. Binding energy: the energy needed to split a nucleus into separate nucleons, equivalent to the mass defect by E = mc².

Charged atoms

22. Ion: an atom or group of atoms with an overall electric charge because it has gained or lost electrons. See what is an ion?
23. Cation: a positive ion, formed when an atom loses electrons, e.g. Na⁺.
24. Anion: a negative ion, formed when an atom gains electrons, e.g. Cl⁻.
25. Isoelectronic: having the same number of electrons. O²⁻, F⁻, Ne and Na⁺ all have 10 electrons. See isoelectronic species

Electron arrangement

26. Energy level / shell: a set of electron states at roughly the same energy and distance from the nucleus, numbered n = 1, 2, 3… See electron shells
27. Subshell: a group of orbitals of the same type within a shell: s, p, d or f.
28. Orbital: a region of space where there is a high probability of finding an electron. Each orbital holds up to two electrons. See atomic orbitals
29. Electron configuration: the arrangement of an atom’s electrons in orbitals, e.g. 1s² 2s² 2p⁶ for neon.
30. Noble-gas (shorthand) configuration: writing the configuration using the previous noble gas as a core, e.g. [Ne] 3s¹ for sodium.
31. Valence electrons: the electrons in the outermost shell, which take part in bonding.
32. Core electrons: the inner electrons that are not involved in bonding.
33. Ground state: the lowest-energy arrangement of an atom’s electrons.
34. Excited state: an arrangement in which one or more electrons have moved to a higher energy level. See excited vs ground state
35. Aufbau principle: electrons fill the lowest-energy available orbitals first. See the Aufbau principle
36. Pauli exclusion principle: no two electrons in an atom can have the same four quantum numbers, so an orbital holds at most two electrons, with opposite spins.
37. Hund’s rule: electrons fill orbitals of equal energy singly, with parallel spins, before pairing up.
38. Unpaired electron: an electron alone in an orbital. Atoms with unpaired electrons are paramagnetic.
39. Paramagnetic: weakly attracted to a magnetic field because of unpaired electrons. The opposite is diamagnetic.
40. Shielding: the reduction in nuclear attraction felt by outer electrons because inner electrons repel them.
41. Effective nuclear charge (Z_eff): the net positive charge an electron actually feels after shielding. See effective nuclear charge

Quantum ideas

42. Quantum number: one of four numbers (n, l, mₗ, mₛ) that describe an electron’s state. See quantum numbers explained
43. Principal quantum number (n): gives the shell and main energy level.
44. Angular momentum quantum number (l): gives the subshell shape: 0 = s, 1 = p, 2 = d, 3 = f.
45. Magnetic quantum number (mₗ): gives the orientation of an orbital.
46. Spin quantum number (mₛ): gives the electron’s spin, +½ or −½.
47. Node: a place in an orbital where the probability of finding an electron is zero. See nodes in orbitals
48. Wave–particle duality: the idea that electrons and light behave as both waves and particles. See wave–particle duality
49. Heisenberg uncertainty principle: position and momentum cannot both be known exactly at the same time.
50. Quantum mechanical model: the modern model of the atom, describing electrons with wavefunctions and probabilities rather than fixed orbits.

Light and spectra

51. Photon: a single packet (quantum) of light energy, with E = hf.
52. Emission spectrum: the set of wavelengths of light given out by excited atoms, seen as bright lines.
53. Absorption spectrum: dark lines in a continuous spectrum, where atoms have absorbed particular wavelengths. See emission vs absorption spectra
54. Ionization energy: the energy needed to remove the outermost electron from a gaseous atom.
55. Electron affinity: the energy change when a gaseous atom gains an electron. See electron affinity
56. Zeeman effect: the splitting of spectral lines in a magnetic field. See the Zeeman effect

Models and history

57. Plum pudding model: J. J. Thomson’s 1904 model of electrons embedded in a sphere of positive charge. See Thomson’s model
58. Nuclear model: Rutherford’s 1911 model of a tiny dense nucleus with electrons around it, based on the gold foil experiment.
59. Bohr model: Niels Bohr’s 1913 model in which electrons occupy fixed circular orbits with set energies. See the Bohr model
60. Radioactive decay: the spontaneous change of an unstable nucleus, releasing radiation (alpha, beta or gamma). See alpha, beta and gamma radiation

Terms students often mix up

Pair	The difference
Atomic number vs mass number	Protons only vs protons + neutrons
Mass number vs relative atomic mass	A whole number for one isotope vs a weighted average for the element
Shell vs orbital	A whole energy level vs one region holding up to two electrons
Isotope vs ion	Different neutrons vs different electrons
Emission vs absorption spectrum	Bright lines on dark vs dark lines on bright

For more of these, see atomic mass vs mass number vs molar mass and isotopes, ions, isomers and allotropes.

How to use this glossary for revision

• Cover and recall. Hide the definitions and try to explain each term aloud in one sentence.
• Make links. Pick two terms (e.g. “shielding” and “ionization energy”) and explain how one affects the other.
• Use examples. For every definition, name a real element that shows it; the periodic table makes this quick.
• Test yourself with the atomic structure practice questions.

Key takeaways

• Atomic structure vocabulary falls into a few themes: particles, numbers, ions, electron arrangement, quantum ideas, spectra and history.
• Many confusing pairs differ in just one detail. Learn them side by side.
• Each definition here links to a full explanation for deeper study.