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This lesson covers the development of atomic models, the properties of subatomic particles, isotopes, and the concepts of mass number, atomic number, and relative atomic mass. A solid grasp of these fundamentals is essential for everything that follows in A-Level Chemistry.
Spec mapping (AQA 7405): This lesson maps directly to §3.1.1.1 (fundamental particles), §3.1.1.2 (mass number and isotopes), and §3.1.2.1 (relative atomic mass and relative molecular mass). Electron configuration (§3.1.1.3) is introduced briefly and developed in detail in the next-but-one lesson. Refer to the official AQA specification document for the exact wording of each section.
Assessment objectives in this lesson: Knowledge recall of subatomic particles, isotope definitions and the structure of the atom is tested on Paper 1 and Paper 2 (AO1). Mass-spectrometry calculations and ion-identification problems most often appear on Paper 1 Section A as short-answer items (AO2). Extended-response questions on the historical development of atomic models — Rutherford's gold foil experiment in particular — appear on Paper 3 and test AO3 (analysis and evaluation of experimental evidence).
Our understanding of the atom has evolved over centuries through key experiments and theoretical breakthroughs.
John Dalton proposed that all matter is made of indivisible atoms. He suggested that atoms of the same element are identical and that chemical reactions involve the rearrangement of atoms. While groundbreaking, this model had no concept of internal structure.
J.J. Thomson discovered the electron using cathode ray tubes. He measured the charge-to-mass ratio of electrons and proposed the "plum pudding" model: a sphere of positive charge with negatively charged electrons embedded within it, like plums in a pudding.
Ernest Rutherford directed alpha particles at a thin gold foil. Most passed straight through, but a small fraction were deflected at large angles and some bounced back. This could only be explained if:
Niels Bohr refined the nuclear model by proposing that electrons orbit the nucleus in fixed energy levels (shells). Electrons can move between energy levels by absorbing or emitting specific amounts of energy. This model successfully explained the line spectrum of hydrogen.
The modern model treats electrons as existing in orbitals — regions of space where there is a high probability of finding an electron. Electrons do not follow fixed orbits but are described by wave functions. This model underpins A-Level electron configuration.
Atoms are made of three subatomic particles:
| Particle | Relative mass | Relative charge | Location |
|---|---|---|---|
| Proton | 1 | +1 | Nucleus |
| Neutron | 1 | 0 | Nucleus |
| Electron | 1/1836 (≈ 0.00055) | −1 | Orbitals around nucleus |
The actual masses are:
Key Point: For most A-Level calculations, the mass of the electron is considered negligible compared to protons and neutrons. The relative mass of an electron is approximately 1/1836 that of a proton.
Every element is defined by two key numbers:
The number of neutrons = A − Z.
In a neutral atom, the number of electrons equals the number of protons. When an atom forms an ion, it gains or loses electrons but the number of protons remains unchanged.
An atom is represented as:
ᴬ_Z X
For example, sodium-23: ²³₁₁Na has 11 protons, 12 neutrons (23 − 11), and 11 electrons in the neutral atom.
Exam Tip: If asked for the number of electrons in an ion, remember to adjust: a 2+ ion has lost 2 electrons, a 2− ion has gained 2 electrons.
Isotopes are atoms of the same element (same number of protons / same atomic number) that have different numbers of neutrons (different mass numbers).
| Isotope | Protons | Neutrons | Mass number | Natural abundance (%) |
|---|---|---|---|---|
| ¹H (protium) | 1 | 0 | 1 | 99.98 |
| ²H (deuterium) | 1 | 1 | 2 | 0.02 |
| ¹²C | 6 | 6 | 12 | 98.9 |
| ¹³C | 6 | 7 | 13 | 1.1 |
| ³⁵Cl | 17 | 18 | 35 | 75.8 |
| ³⁷Cl | 17 | 20 | 37 | 24.2 |
Isotopes of the same element have:
Common Misconception: Students sometimes say isotopes have "different chemical properties." This is incorrect. Chemical properties depend on the electron configuration, which is the same for all isotopes of an element.
Because the actual masses of atoms are incredibly small, we use a relative scale based on carbon-12.
Key Definition: Relative atomic mass is a weighted mean because it takes into account the natural abundances of each isotope, not just their masses.
The formula for Aᵣ is:
Aᵣ = Σ (isotopic mass × percentage abundance) / 100
Chlorine has two isotopes: ³⁵Cl (75.8%) and ³⁷Cl (24.2%).
Aᵣ = (35 × 75.8 + 37 × 24.2) / 100
Aᵣ = (2653.0 + 895.4) / 100
Aᵣ = 3548.4 / 100
Aᵣ = 35.5 (to 1 decimal place)
Boron has two isotopes: ¹⁰B (19.9%) and ¹¹B (80.1%).
Aᵣ = (10 × 19.9 + 11 × 80.1) / 100
Aᵣ = (199.0 + 881.1) / 100
Aᵣ = 1080.1 / 100
Aᵣ = 10.8 (to 1 decimal place)
Silicon has three isotopes: ²⁸Si (92.2%), ²⁹Si (4.7%), ³⁰Si (3.1%).
Aᵣ = (28 × 92.2 + 29 × 4.7 + 30 × 3.1) / 100
Aᵣ = (2581.6 + 136.3 + 93.0) / 100
Aᵣ = 2810.9 / 100
Aᵣ = 28.1 (to 1 decimal place)
Exam Tip: Always show your working in relative atomic mass calculations. A common error is to simply average the mass numbers without weighting by abundance. The Aᵣ of chlorine is 35.5, NOT 36 (which would be the simple average of 35 and 37).
When atoms gain or lose electrons, they form ions:
How many protons, neutrons, and electrons are in ⁵⁶₂₆Fe³⁺?
An ion X²⁻ has 10 electrons and 8 protons. Identify the element and write its notation.
This lesson connects to several other AQA Chemistry topics — A-Level papers test cross-topic understanding through extended-response questions:
Question 1. [13 marks total]
(a) Define the term isotope. [2 marks]
(b) A sample of an unknown element X produced the following mass spectrum:
| Mass / (m/z) | Relative abundance / % |
|---|---|
| 24 | 78.99 |
| 25 | 10.00 |
| 26 | 11.01 |
Calculate the relative atomic mass of X to two decimal places. State which element X is, justifying your answer using a periodic table. [4 marks]
(c) Rutherford's gold foil experiment (1911) provided evidence that overturned Thomson's plum pudding model. State two observations Rutherford made, and explain what each observation tells us about atomic structure. [4 marks]
(d) Explain why an ion of ²⁶₁₂Mg²⁺ has the same number of electrons as a neutral atom of neon. State the electron configuration of this ion. [3 marks]
(a) Definition of isotope [2 marks, AO1]
Accept "atoms with the same number of protons but a different number of neutrons". Do not accept "atoms with different masses" alone (too vague — must mention neutrons explicitly).
(b) Relative atomic mass calculation and identification [4 marks, AO2]
Tolerance: accept 24.31 or 24.32. Award all 4 marks for correct final answer with working shown; lose 1 mark for correct method but arithmetic error; lose 2 marks for using a simple average (24 + 25 + 26)/3 = 25.
(c) Rutherford observations and conclusions [4 marks, AO3]
Award 1 mark per observation and 1 mark per linked explanation, up to 4 marks total.
Do not award marks for observations alone (without conclusion) or conclusions alone (without observation). The mark scheme rewards the explicit linking of evidence to claim — this is an AO3 question.
(d) Mg²⁺ electron count and configuration [3 marks, AO1 + AO2]
Common error: candidates correctly identify the electron count but fail to write out the configuration. Read both halves of the question before answering.
The three responses below cover the meaningful A-Level range: Grade C (the borderline-pass floor), Grade B (solid mark-scheme coverage), and Grade A* (top-band synthesis). No Grade D or E responses are shown — no A-Level student is aiming for those bands, and modelling failure adds nothing pedagogically. The commentary after each response (editorial, not a real examiner report) names the marks earned and the specific moves that differentiate from adjacent bands.
(a) Isotopes are atoms of the same element (same number of protons, same atomic number) that have different numbers of neutrons, and therefore different mass numbers.
(b) Aᵣ = (24 × 78.99 + 25 × 10.00 + 26 × 11.01) / 100 = (1895.76 + 250.00 + 286.26) / 100 = 24.32. Comparing 24.32 to the periodic table, this matches magnesium (Aᵣ = 24.31). The small discrepancy is within rounding tolerance for the given abundances. X = Mg.
(c) Observation 1: most alpha particles passed straight through the foil — this indicates the atom is mostly empty space, with very few particles encountering anything to deflect them. Observation 2: a small fraction were deflected through large angles, including some that bounced back almost in the direction of the incident beam — this requires a small, dense, positively charged region (the nucleus) capable of producing sufficient electrostatic repulsion on the positive alpha particles.
(d) Mg²⁺ is formed by Mg losing 2 electrons from the neutral atom (which has 12 protons and 12 electrons). The ion therefore has 10 electrons. Neon (Z = 10) also has 10 electrons in its neutral atom — so Mg²⁺ is isoelectronic with neon. Electron configuration: 1s² 2s² 2p⁶ (or [Ne]).
Examiner commentary: Approximately 9-10/13. Mark scheme coverage is broad but thin in places. Part (a) earns both marks for a precise definition (M1 same-element, M1 different-neutron-count). Part (b) gets all 4 marks (weighted-mean expression, correct arithmetic, two-decimal-place answer, magnesium identified). Part (c) gets 3-4 (observations linked to conclusions; "positively charged region" mentioned). Part (d) gets 2-3 (correct electron count and "isoelectronic" used). The response is correct throughout but lacks the precision and depth that lifts it to Grade B or above — no quantitative scattering ratio in (c), no mention of Aᵣ tolerance in (b), no electron-configuration notation extended to neutral Mg in (d).
(a) Isotopes are atoms of the same element — atoms with the same atomic number Z — that differ in mass number A because they contain different numbers of neutrons. They have identical chemical properties (same electron configuration) but different physical properties such as density, diffusion rate, and natural abundance.
(b) Aᵣ = Σ(isotopic mass × % abundance) / 100 = (24 × 78.99 + 25 × 10.00 + 26 × 11.01) / 100. Step by step: 1895.76 + 250.00 + 286.26 = 2432.02; divide by 100 → Aᵣ = 24.32. Comparison: the periodic table gives Aᵣ(Mg) = 24.31; the calculated 24.32 sits within the rounding tolerance of the two-decimal-place input data. X is magnesium. A common student misconception is the simple average 25, which would (incorrectly) identify sodium — the weighted mean is essential because abundances differ markedly.
(c) Observation 1: the majority of alpha particles (positively charged helium nuclei) passed undeflected straight through the gold foil. This is evidence that the atom is mostly empty space, since alpha particles encountering a substantial barrier would deflect. Observation 2: a small fraction were deflected through angles greater than 90°, with a few rebounding almost back along the incident path. This is only possible if the positive charge and the majority of the mass of the atom are concentrated in a very small, dense region — the nucleus — whose electrostatic repulsion of an incoming positive alpha particle is sufficient to reverse its trajectory.
(d) The neutral magnesium atom has 12 protons and 12 electrons (1s² 2s² 2p⁶ 3s²). Forming Mg²⁺ requires the loss of two electrons from the 3s orbital, leaving 10 electrons in the 1s² 2s² 2p⁶ configuration. Neon (Z = 10) has 10 electrons in its neutral atom, also configured 1s² 2s² 2p⁶. Therefore the two species are isoelectronic.
Examiner commentary: Approximately 11-12/13. Substantially better than Grade C: chemistry-of-isotopes physical-property contrast in (a); the simple-average misconception explicitly called out in (b); the alpha-particle charge named in (c); the neutral-Mg electron configuration shown explicitly in (d). The student is consistently using A-Level-precise vocabulary (isoelectronic, electrostatic repulsion, weighted mean). To progress to A*, the response would benefit from a quantitative scattering ratio in part (c) (Rutherford's 1-in-8000 figure), kinetic-isotope-effect commentary in part (a), and an extension of the isoelectronic series argument to Na⁺ / F⁻ / O²⁻ / N³⁻ in part (d).
(a) Isotopes are atoms of the same element — same Z, same proton count, same electron count when neutral — that differ in mass number A because they contain different numbers of neutrons. The chemistry of an element is governed by its electron configuration, which is identical for all isotopes; physical properties (density, diffusion rate per Graham's law, natural abundance, vibrational frequencies of bonds via the reduced mass) differ measurably. Hydrogen exemplifies this most clearly: protium (¹H), deuterium (²H) and tritium (³H) have the same chemistry but differ markedly in bond strengths and reaction rates (the kinetic isotope effect).
(b) Aᵣ = Σ(mᵢ × pᵢ) / 100 = (24 × 78.99 + 25 × 10.00 + 26 × 11.01) / 100 = 24.3202 → 24.32 (to 2 d.p.). The IUPAC standard atomic weight of magnesium is 24.305 ± 0.006 — the calculated value sits within this stated uncertainty range. X is magnesium. The simple arithmetic mean of 25.00 would mis-identify sodium and would be inconsistent with the weighted-mean axiom: when abundances differ significantly, only the weighted mean is physically meaningful.
(c) Rutherford's gold foil experiment (Geiger and Marsden, supervised by Rutherford, 1909–1911) gave two key observations. (1) The vast majority of alpha particles passed undeflected through the foil — evidence that the atom is mostly empty space, since gold foil is approximately 1000 atoms thick yet alpha particles traverse it without significant electrostatic interaction. (2) Approximately 1 in 8000 alpha particles were deflected by more than 90°, with a tiny fraction rebounding. This scattering ratio is consistent with a very small, dense, positively charged nucleus of diameter approximately 10⁻¹⁴ m (compared to atomic diameter approximately 10⁻¹⁰ m, a factor of 10⁻⁴). The observation directly disproved Thomson's plum pudding model, which predicted only small-angle deflections, and established the nuclear model that Bohr (1913) and Schrödinger (1926) subsequently refined.
(d) The neutral magnesium atom (Z = 12) has 12 protons and 12 electrons (configuration 1s² 2s² 2p⁶ 3s²). Ionisation removes the two 3s electrons sequentially; the resulting Mg²⁺ ion has 10 electrons (1s² 2s² 2p⁶ = [Ne]). Neon is also Z = 10, configuration 1s² 2s² 2p⁶ — the two species are isoelectronic. This argument extends naturally to a wider 10-electron isoelectronic family: Na⁺, F⁻, O²⁻, N³⁻ all share the [Ne] configuration. Across this series, ionic radius decreases with increasing nuclear charge (since the same 10 electrons are increasingly attracted to a larger Z), a periodicity trend you will meet again in §3.2.1.
Editorial commentary (Grade A):* Reaches A* by adding genuinely undergraduate-adjacent reasoning: the kinetic isotope effect (a connection to physical chemistry), IUPAC standard atomic weight with formal uncertainty notation (a connection to scientific data conventions), the quantitative size ratio of nucleus to atom (10⁻⁴) with proper orders of magnitude, the historical chain Rutherford → Bohr → Schrödinger, and the isoelectronic-series argument extended to ionic-radius trends. Each of these moves is what separates a student who has mastered the spec from one who has internalised the underlying physics.
Many candidates lose marks in this topic through a predictable set of slips. Pedagogical observations from chemistry teaching, not from any specific examiner report:
A frequent pitfall in the Rutherford observation question is to describe the observations without explaining what each tells us about atomic structure. The mark scheme awards 1 mark per linked observation–conclusion pair; observations alone score nothing.
These subtler errors distinguish a B-grade from an A-grade response:
Three undergraduate-adjacent extensions of this material that strengthen Oxbridge or top-university interview preparation:
| Concept | Key facts |
|---|---|
| Atomic number (Z) | Number of protons; defines the element |
| Mass number (A) | Protons + neutrons |
| Isotopes | Same Z, different A (different neutrons) |
| Relative atomic mass | Weighted mean mass relative to ¹²C / 12 |
| Neutral atom | Protons = electrons |
| Positive ion | Fewer electrons than protons |
| Negative ion | More electrons than protons |
Exam Tip: In multiple-choice questions, you may be asked to identify the correct definition of isotopes. The key phrase is "atoms of the same element with different numbers of neutrons." Do not say "different mass" — be specific about neutrons.
This content is aligned with the AQA A-Level Chemistry (7405) specification.