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Spec Mapping — OCR H432 Module 2.1.1 — Atomic structure and isotopes, content statements covering the structure of the atom in terms of protons, neutrons and electrons; atomic number and mass number; isotopes and their notation; and the historical development of atomic models from Dalton through Thomson, Rutherford, Bohr and Chadwick (refer to the official OCR H432 specification document for exact wording). This is the foundational lesson of Module 2 and every subsequent piece of stoichiometry, mass-spectrometry and bonding work in Modules 2 and 3 rests on it.
Atomic structure is the conceptual entry-point of A-Level Chemistry. Every mole calculation you will perform across the next two years assumes a particular model of the atom — protons in a tiny dense nucleus, electrons in orbitals around it, with the proton count defining the element and the neutron count distinguishing isotopes. In this lesson you will build that model in stages: the historical experiments that established it, the three sub-atomic particles and their relative properties, the A/Z notation that lets us specify any nuclide, the treatment of ions as electron-modified atoms, and finally isotopes — the same-element-different-mass species that make natural relative atomic masses non-integer and that drive the entire next-lesson treatment of Ar. You also need a sense of the historical narrative so that you can frame the atomic model as a scientific construct, not a textbook fact.
Key Definition: An isotope is one of two or more atoms of the same element that have the same number of protons (and therefore the same atomic number Z) but different numbers of neutrons (and therefore different mass numbers A).
The idea that matter is made of indivisible particles traces back to the ancient Greek philosopher Democritus (c. 460 BCE), who coined the word atomos meaning "uncuttable". Scientific atomic theory, however, did not emerge until John Dalton published his work in 1808. Dalton's school of thought, paraphrased rather than quoted verbatim, proposed that elements are composed of tiny indivisible particles called atoms; that atoms of a given element are identical in mass and properties; that compounds form when atoms of different elements combine in simple whole-number ratios; and that chemical reactions rearrange atoms without creating or destroying them. Dalton's model is wrong in its first claim (atoms are divisible) and incomplete in its second (atoms of a given element can have different masses — isotopes), but it correctly captures the conservation of atoms that underpins balanced equations.
Dalton's model was refined dramatically over the next century by a sequence of experimental discoveries:
| Year | Scientist | Key experiment | Resulting model element |
|---|---|---|---|
| 1869 | Mendeleev | Periodic-table arrangement by atomic weight | Periodic pattern of properties (later explained by electron configuration) |
| 1897 | J. J. Thomson | Cathode-ray deflection | Discovery of the electron; "plum-pudding" model with electrons embedded in a uniform positive sphere |
| 1909 | Rutherford, Geiger, Marsden | Gold-foil α-particle scattering | Most α particles passed through, a few bounced back at large angles — implying a tiny dense positive nucleus with mostly empty space around it |
| 1913 | Niels Bohr | Hydrogen-spectrum analysis | Electrons occupy fixed energy levels (shells); transitions between levels emit photons of specific frequencies |
| 1932 | James Chadwick | Neutral-radiation beryllium experiment | Discovery of the neutron — uncharged nucleon explaining the mass discrepancy between atomic number and atomic mass |
Rutherford's gold-foil experiment is the historical pivot. Geiger and Marsden's beam of α particles striking a thin gold foil was expected (on the Thomson model) to pass through with at most tiny deflections; instead, a small fraction were scattered through large angles, with a handful bouncing almost straight back. Rutherford's interpretation — that the atom must contain a small, dense, positively charged core (the nucleus) — overturned the plum-pudding picture in a single stroke and established the nuclear-atom paradigm we still use. Chadwick's 1932 discovery of the neutron completed the trio of sub-atomic particles and finally explained why 12C has mass 12, not the 6 that its 6 protons alone would imply.
Atoms are the fundamental units of chemical identity. Despite being unimaginably small (around 10−10 m across), they are themselves composed of three sub-atomic particles: protons, neutrons and electrons. You must learn both the relative and absolute values of their masses and charges; OCR will examine the relative figures verbally and the absolute figures via calculation.
| Particle | Symbol | Relative mass | Relative charge | Location |
|---|---|---|---|---|
| Proton | p | 1 | +1 | Nucleus |
| Neutron | n | 1 | 0 | Nucleus |
| Electron | e⁻ | 18361 (≈0) | −1 | Orbitals around nucleus |
The absolute (SI) values, which you will need when you meet ionisation energies in Module 3 and electrochemistry in Module 5, are:
| Particle | Mass / kg | Charge / C |
|---|---|---|
| Proton | 1.673×10−27 | +1.602×10−19 |
| Neutron | 1.675×10−27 | 0 |
| Electron | 9.109×10−31 | −1.602×10−19 |
The nucleus contains the protons and neutrons (collectively the nucleons) and holds essentially all the mass of the atom. The electrons occupy the much larger volume around the nucleus, arranged in orbitals that you will quantise in detail when you meet electron configuration in Module 2.2.
Key Fact: If an atom were the size of a football stadium, the nucleus would be about the size of a pea at the centre circle. Matter is overwhelmingly empty space.
The density of the nucleus is around 2×1017 kg m−3 — the same density found in neutron stars. This astonishing concentration is what produces the strong-angle scattering Rutherford observed.
Two numbers fully describe any specific nuclide:
The standard isotope notation is:
ZAX
For example, 1123Na (sodium-23) has:
flowchart LR
A[Nuclide AZX] --> B[Nucleus]
A --> C[Electron cloud]
B --> D[Z protons defines element]
B --> E[A - Z neutrons distinguishes isotope]
C --> F[Z electrons in neutral atom]
F --> G[Modified by ion charge]
| Species | Protons | Neutrons | Electrons |
|---|---|---|---|
| 16O | 8 | 8 | 8 |
| 20Ne | 10 | 10 | 10 |
| 39K | 19 | 20 | 19 |
| 56Fe | 26 | 30 | 26 |
| 238U | 92 | 146 | 92 |
When an atom gains or loses electrons it becomes an ion. Crucially, the number of protons does not change — only the electron count changes. Chemistry happens through electrons; the proton inventory is fixed for a given element.
How many protons, neutrons and electrons are in 27Al3+?
Al³⁺ is isoelectronic with neon — both have 10 electrons and the configuration 1s22s22p6. This is one reason Al prefers a 3+ ion: doing so achieves a stable noble-gas configuration.
How many protons, neutrons and electrons are in 32S2−?
S²⁻ is isoelectronic with argon.
How many protons, neutrons and electrons are in 56Fe3+?
Fe³⁺ is not isoelectronic with a noble gas — transition-metal ions achieve stable half-filled or fully filled d-subshell configurations instead, as you will see in Module 5 on transition-metal chemistry.
Identify all the common ions and atoms that are isoelectronic with neon (1s22s22p6, 10 electrons). Then do the same for argon (18 electrons).
For neon (Z=10): the 10-electron club contains 16O2−, 19F−, Ne itself, 23Na+, 24Mg2+, 27Al3+ and (more rarely) 14N3−. Note how the proton count walks from 7 up to 13 across this isoelectronic series, while the electron count is fixed at 10 — so ionic radius decreases monotonically with Z across the series (more protons pulling on the same 10 electrons). N³⁻ is the largest (52 % bigger than Al³⁺ by Pauling radius), Al³⁺ the smallest. This is the conceptual basis of the Module 3 ionic-radius trend.
For argon (Z=18): the 18-electron club includes 32S2−, 35Cl−, Ar, 39K+, 40Ca2+ and 45Sc3+. Again the radius shrinks as Z climbs.
The isoelectronic concept is one of the most powerful pattern-recognition tools in inorganic chemistry: if two species have the same electron configuration, their chemistry — particularly the geometry and reactivity of any covalent bonding they engage in — closely mirrors each other. The classic A-Level move is to recognise that BF₃ (24 electrons) and CO₃²⁻ (24 electrons) are isoelectronic and both adopt trigonal-planar geometries; this kind of cross-cutting argument is examiner-favoured at top-band.
A nuclide has 30 neutrons and forms a 2+ ion with 20 electrons. Identify it.
Top candidates note that the Ti²⁺ ion has lost both its 4s electrons but retains all four 3d electrons, giving the configuration [Ar]3d2 — a key fact for the Module 5 colour-of-transition-metal-ions treatment.
Key Definition: Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons.
Because isotopes have the same number of protons (and so, in neutral atoms, the same number of electrons), they have identical chemical properties. Chemistry is governed by electron arrangement, and isotopes are electronically identical.
However, isotopes have slightly different physical properties — density, rate of diffusion, boiling point, bond strength — because their masses differ.
| Element | Isotope | Protons | Neutrons | Natural abundance |
|---|---|---|---|---|
| Hydrogen | 1H (protium) | 1 | 0 | 99.985 % |
| Hydrogen | 2H (deuterium, D) | 1 | 1 | 0.015 % |
| Hydrogen | 3H (tritium, T) | 1 | 2 | Trace (radioactive) |
| Carbon | 12C | 6 | 6 | 98.93 % |
| Carbon | 13C | 6 | 7 | 1.07 % |
| Carbon | 14C | 6 | 8 | Trace (radioactive — radiocarbon dating) |
| Chlorine | 35Cl | 17 | 18 | 75.78 % |
| Chlorine | 37Cl | 17 | 20 | 24.22 % |
| Uranium | 235U | 92 | 143 | 0.72 % |
| Uranium | 238U | 92 | 146 | 99.27 % |
Consider 35Cl and 37Cl. Both have 17 protons and 17 electrons arranged as 1s22s22p63s23p5. When chlorine reacts with sodium, both isotopes form Cl⁻ ions just as readily because the reaction involves the outer 3p electron only. The extra two neutrons in 37Cl make no difference to bonding behaviour.
Some isotopes are unstable and decay, emitting radiation. These radioisotopes have many applications:
Synoptic Links — Connects to:
ocr-alevel-chemistry-atoms-moles / relative-atomic-and-isotopic-mass— the existence of isotopes is the entire reason Ar is a weighted mean rather than a whole number, and the A/Z notation you learn here is the input format for every Ar calculation.ocr-alevel-chemistry-atoms-moles / mass-spectrometry— the mass spectrometer literally counts isotopes, producing the abundance data that the next lesson uses; everything in this lesson is the conceptual foundation of TOF-MS.ocr-alevel-chemistry-acids-redox-bonding / ionic-bonding-and-structure— the cation/anion electron-modification you have practised here is the same arithmetic underpinning lattice formation in Module 2.2.Practical Activity Group anchor: this lesson is non-PAG (theoretical foundation), but it underpins PAG 1 (moles determination) and PAG 2 (acid–base titration) by establishing the proton/neutron/electron picture that justifies why "balanced equation = balanced atoms".
Question (6 marks): State what is meant by the term isotopes. An aluminium ion 27Al3+ is investigated. Determine the number of protons, neutrons and electrons present in this ion, showing your reasoning. Explain why isotopes of an element have identical chemical properties.
| Mark | AO | Awarded for |
|---|---|---|
| 1 | AO1 | Defining isotopes as atoms of the same element / same number of protons |
| 2 | AO1 | Stating that isotopes have different numbers of neutrons (different A for same Z) |
| 3 | AO2 | Identifying 13 protons in 27Al3+ from the atomic number of aluminium |
| 4 | AO2 | Identifying 14 neutrons by A−Z=27−13 |
| 5 | AO2 | Identifying 10 electrons by recognising 3 electrons lost from neutral Al |
| 6 | AO3 | Explaining identical chemistry in terms of identical electron configuration / chemistry being determined by outer electrons |
AO split: AO1 =2, AO2 =3, AO3 =1.
Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons. The aluminium ion 27Al3+ has the atomic number 13, so it has 13 protons. To find the neutrons, I do mass number minus atomic number, so 27−13=14 neutrons. Because the ion has a 3+ charge, it has lost 3 electrons compared to a neutral atom of aluminium. A neutral Al atom has 13 electrons, so the ion has 13−3=10 electrons. Isotopes have the same chemical properties because they have the same number of electrons. Chemistry happens through the electrons, and the neutrons in the nucleus do not change how the atom reacts. So all isotopes of an element react the same way.
Examiner commentary: M1 awarded for the definition (same protons, different neutrons); M1 (AO2) for 13 protons; M1 (AO2) for 14 neutrons; M1 (AO2) for 10 electrons with the lost-three-electrons justification; M1 (AO1) implicit in correctly using Z=13 for Al. The candidate scores about 5/6 — the AO3 mark is partially earned by "chemistry happens through the electrons" but a top candidate would frame it explicitly as identical electron configuration in the outer shell. It secures the procedural AO2 marks but lacks the precise terminology required for the AO3 discriminator.
Isotopes are atoms of the same element — and so, by definition, with the same atomic number Z — that have different mass numbers A because they contain different numbers of neutrons. For the species 1327Al3+, the atomic number gives the proton count directly: Z=13, so there are 13 protons. The neutron count is A−Z=27−13= 14 neutrons. The 3+ charge denotes that three electrons have been removed from the neutral atom, so the electron count is 13−3= 10 electrons. The electron configuration of the Al³⁺ ion is 1s22s22p6, which is isoelectronic with neon — explaining why this is the thermodynamically preferred charge state.
Isotopes have identical chemical properties because chemistry is governed entirely by the outer-shell electron configuration. Reactions involve the making and breaking of bonds, which require electron rearrangement; the nucleus is a spectator. Two isotopes of an element have the same number of protons and (when neutral) the same number of electrons in the same configuration, so they present an identical interface to other atoms. The additional neutrons in the heavier isotope contribute only to mass, not to bonding behaviour. Physical properties such as rate of diffusion and bond-stretch frequency do depend on mass and so do differ slightly — a phenomenon exploited in the kinetic isotope effect to identify rate-limiting steps in organic mechanisms.
Examiner commentary: Full 6/6. M1 (definition with Z explicit), M1 (neutron count distinction in the definition), M1 (AO2 protons), M1 (AO2 neutrons), M1 (AO2 electrons with electron-configuration justification), M1 (AO3 outer-shell explanation plus kinetic-isotope-effect aside). The isoelectronic-with-neon move, the 1s22s22p6 configuration, and the explicit physical-vs-chemical separation are the discriminators that lift this answer from a strong A to A*.
Pedagogical observations — no fabricated candidate percentages:
At undergraduate level, the picture of the atom you have just learned is refined in two directions. Physical chemistry introduces quantum-mechanical orbitals as solutions to the Schrödinger equation; the Bohr "shells" become discrete energy levels indexed by quantum numbers n, ℓ, mℓ and ms, with shapes (s, p, d, f) you will preview later in Module 2.2. The first-year Physical Chemistry: Quantum Mechanics and Spectroscopy course at most UK universities formally develops these orbital wavefunctions and introduces the radial probability function 4πr2∣ψ(r)∣2 — a graph that shows where in space an electron is most likely to be found around a nucleus. The familiar A-Level pictures of dumbbell-shaped p-orbitals and four-lobed d-orbitals are simplified visualisations of these radial-and-angular probability distributions. Importantly, the Bohr model's "fixed orbits" are now understood as a useful pedagogical fiction: an electron in a 1s orbital is not orbiting at a fixed radius but exists as a probability cloud peaked at ∼0.53 A˚ from the proton (the Bohr radius). Nuclear chemistry treats the nucleus itself as a many-body quantum system with a "shell model" of nucleon energy levels — the same magic-numbers logic that explains why 4He, 16O, 40Ca and 208Pb are unusually stable. The magic numbers 2, 8, 20, 28, 50, 82 and 126 correspond to closed nuclear shells, analogous to noble-gas electron configurations. Analytical chemistry courses extend mass spectrometry from the simple TOF picture of the next lesson to high-resolution Orbitrap and FT-ICR techniques that resolve 12CH4 from 13CH3. An Oxbridge-style interview prompt: "Why is helium-4 so much more stable than helium-3, even though they are both 'helium'?" — the answer involves nuclear pairing energy, a topic absent from A-Level but motivated by the very neutron-count distinction you have just learnt. A second favourite prompt: "Iodine-131 and iodine-127 are both iodine. Why would a doctor inject a thyroid patient with one of them and not the other?" — the answer requires recognising that the isotopes are chemically identical (the thyroid concentrates iodine without isotope discrimination) but that 131I is a β and γ emitter, providing therapeutic radiation to thyroid tissue while 127I is stable and biologically inert beyond normal dietary roles.
The subtle errors that distinguish A from A*:
Atoms consist of a dense nucleus of protons and neutrons surrounded by electrons in orbitals. The atomic number Z — the proton count — defines the element; the mass number A is the sum of protons and neutrons. Isotopes are atoms of the same element (Z fixed) with different neutron counts (A different) and therefore different masses; because they have identical electron configurations, they have identical chemistry. Ions are atoms with electron counts modified by the ion's charge; protons are never altered. The historical model has been built up over two centuries from Dalton's indivisible-atom postulate through Thomson's plum-pudding, Rutherford's nuclear atom, Bohr's quantised shells and Chadwick's neutron — and every subsequent lesson in this course assumes this nuclear-atom picture.
Reference: OCR A-Level Chemistry A (H432) specification 2.1.1 (refer to the official OCR H432 specification document for exact wording).