In 1897, J.J. Thomson discovered the electron and shattered the 2,400-year-old idea of an indivisible atom. By 1911, Rutherford's gold foil experiment revealed a tiny, dense, positive nucleus surrounded mostly by empty space. By 1913, Bohr had quantised electron energies into fixed orbits. Each model was revolutionary — and each was later proved incomplete. Science advances not by being right, but by making better models as new evidence arrives. This is IQ3: how do scientists develop understanding of atomic structure from experimental evidence?
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In 1911, Ernest Rutherford fired positively charged alpha particles at an extremely thin sheet of gold foil. Most passed straight through, but a tiny fraction bounced back. At the time, scientists believed atoms were like smooth, solid balls of positive charge with electrons embedded in them (the "plum pudding" model). Why did the bouncing particles force scientists to completely change their view of what an atom looks like?
Before reading on, write your best answer. What must be true about the structure of the atom if most particles pass through but some are deflected straight back?
Core Content
Wrong: Rutherford's nuclear model explained why electrons do not fall into the nucleus.
Right: Rutherford's model proposed a dense positive nucleus with orbiting electrons but could not explain electron stability. Bohr later proposed quantised energy levels to explain why electrons remain in stable orbits without radiating energy and collapsing.
| Model | Scientist (year) | Key features | Evidence base | Limitation / why it was superseded |
|---|---|---|---|---|
| Solid sphere (Dalton) | Dalton, 1803 | Atom as indivisible solid sphere. Elements have unique atomic masses. Compounds form from fixed ratios. | Law of definite proportions (fixed mass ratios in compounds); law of conservation of mass. | Assumed atoms were indivisible. Discovery of electrons (1897) showed substructure exists. |
| Plum pudding (Thomson) | Thomson, 1904 | Atom = sphere of positive charge with electrons embedded throughout (like plums in a pudding). Overall neutral. | Discovery of electrons by Thomson (1897) via cathode ray tube — showed negative particles existed inside atoms. | Rutherford's gold foil experiment (1909–1911) showed most of the mass was concentrated in a small nucleus — not spread out. The plum pudding model predicted alpha particles should pass through uniformly. |
| Nuclear model (Rutherford) | Rutherford, 1911 | Tiny, dense, positively charged nucleus surrounded by mostly empty space. Electrons orbit the nucleus at large distances. | Gold foil experiment: alpha particles fired at gold foil; most passed straight through but a small fraction were deflected at large angles, some reflected back. Only a concentrated positive charge could explain these results. | Classical physics: orbiting electrons should continuously radiate energy and spiral into the nucleus within nanoseconds — atoms would be unstable. Also could not explain atomic emission spectra (discrete lines, not continuous). |
| Bohr model | Bohr, 1913 | Electrons occupy fixed circular orbits (shells) at specific energy levels. Electrons can jump between levels by absorbing/emitting photons of specific energy. Each orbit has a fixed energy. | Hydrogen emission spectrum: discrete coloured lines (Balmer series) at specific wavelengths. Bohr calculated these matched the energy differences between his proposed energy levels exactly. | Only worked precisely for hydrogen (one-electron atom). Could not explain multi-electron spectra or the fine structure of spectral lines. Superseded by quantum mechanical model (Schrödinger, 1926) using orbitals (probability clouds) instead of fixed circular orbits. |
This is the most important single experiment in the history of atomic theory and is frequently examined.
When atoms are excited (by heat or electrical energy), electrons jump to higher energy levels. When they fall back to lower levels, they release photons of light. The energy of the photon matches the energy difference between the two levels:
This is why a sodium street lamp emits a characteristic yellow-orange colour, and why hydrogen emits a specific set of red, blue-green, blue, and violet lines (Balmer series). Each element has a unique spectral fingerprint — used in spectroscopy to identify elements.
Worked Examples
Step 1 — Thomson's model and what it explained
Thomson discovered electrons in 1897 via cathode ray tube experiments.
Cathode ray = beam of negative particles deflected by electric/magnetic fields.
Evidence explained: (1) atoms contain negative particles (electrons), (2) atoms are overall neutral → must contain compensating positive charge.
Thomson's model: positive charge spread uniformly throughout atom; electrons embedded within.
Step 2 — New evidence that challenged Thomson's model
Rutherford (1909–1911): fired alpha particles (positive, heavy, fast) at thin gold foil.
Expected result from plum pudding: alpha particles should pass through with minor deflections as they interact with the diffuse positive sphere.
Actual result:
• Most alpha particles passed straight through (consistent with mostly empty space)
• ~1 in 8,000 deflected at angles > 90° (impossible if positive charge was diffuse)
• ~1 in 20,000 reflected straight back (nucleus must be extremely small and dense)
Thomson's model CANNOT explain large-angle deflections — a diffuse positive sphere would only produce small deflections.
Step 3 — Rutherford's nuclear model addresses the evidence
Concentrated positive charge (nucleus): explains large-angle and back-scatter deflections — only a tiny, very dense, very positive region could repel alpha particles so strongly.
Mostly empty space: explains why the vast majority of alpha particles passed through undeflected.
Electrons outside nucleus at large distances: explains the overall neutral atom.
Step 1 — Read the nuclide notation
Mass number A = 56 (top number = protons + neutrons)
Atomic number Z = 26 (bottom number = protons)
Charge = 2+ (ion that has lost 2 electrons)
Step 2 — Calculate subatomic particles
Protons = Z = 26
Neutrons = A − Z = 56 − 26 = 30
Electrons: neutral Fe would have 26 electrons.
Fe²⁺ has lost 2 electrons: electrons = 26 − 2 = 24
Step 3 — Check charge
Protons (26) − Electrons (24) = +2 ✓ (confirms the 2+ charge)
Activities
1 For each experimental observation, identify which atomic model it supported or disproved, and explain why: (a) Cathode rays were deflected by electric and magnetic fields. (b) Most alpha particles passed through gold foil undeflected. (c) Hydrogen emits light at specific wavelengths (discrete spectral lines).
2 State one key limitation of the Rutherford nuclear model and explain how Bohr's model addressed it.
A Complete the table:
| Nuclide | Z | A | Protons | Neutrons | Electrons |
|---|---|---|---|---|---|
| ¹²₆C | |||||
| ³⁵₁₇Cl⁻ | |||||
| ²³₁₁Na⁺ | |||||
| ¹⁹⁷₇₉Au |
Look back at what you wrote in the Think First section. What has changed? What did you get right? What surprised you?
Multiple Choice
5 random questions from a replayable lesson bank — feedback shown immediately
Short Answer
6. Describe Rutherford's gold foil experiment, including the experimental design, observations, and the conclusions drawn about atomic structure. 5 MARKS
7. Explain how the development of atomic models illustrates the nature of science — specifically the idea that models are revised when new evidence emerges. Use at least two specific historical examples. 4 MARKS
1. (a) Cathode rays deflected by fields → supported Thomson's discovery of electrons (negative particles in the atom); disproved Dalton's solid sphere (indivisible atoms cannot contain subparticles). (b) Most alpha particles through undeflected → supported Rutherford's nuclear model (mostly empty space); incompatible with Thomson's model (diffuse positive sphere should cause uniform deflection). (c) Discrete spectral lines → supported Bohr's model (fixed quantised energy levels produce specific photon energies); incompatible with Rutherford (no explanation for specific energies).
2. Rutherford's limitation: Classical physics predicted orbiting electrons would continuously lose energy (accelerating charges radiate), causing them to spiral into the nucleus within nanoseconds — atoms would be unstable and collapse. Bohr addressed this by proposing electrons exist in fixed, allowed energy levels (orbits) where they do not radiate energy. Electrons only emit or absorb energy when jumping between levels.
¹²₆C: Z=6, A=12, protons=6, neutrons=6, electrons=6. ³⁵₁₇Cl⁻: Z=17, A=35, protons=17, neutrons=18, electrons=18 (Cl⁻ gains 1 electron). ²³₁₁Na⁺: Z=11, A=23, protons=11, neutrons=12, electrons=10 (Na⁺ loses 1 electron). ¹⁹⁷₇₉Au: Z=79, A=197, protons=79, neutrons=118, electrons=79 (neutral atom).
1. B — Large-angle deflections are the key disproof. The plum pudding model predicts only small, uniform deflections from diffuse positive charge.
2. C — Z=16 (protons), neutrons=32−16=16, electrons=16+2=18 (S²⁻ gains 2 electrons).
3. D — Bohr's model worked well for hydrogen (one electron) but failed for multi-electron atoms. Option C describes Rutherford's limitation, not Bohr's.
4. A — Discrete lines = quantised energy levels. Continuous spectrum would come from continuously variable electron energies.
5. B — Rutherford's nuclear model: small dense nucleus + mostly empty space. Thomson's had positive charge throughout; Dalton's was solid.
Q6 (5 marks): Design: Rutherford directed a beam of alpha particles (positively charged, from a radioactive source) through a very thin gold foil (~100 nm thick). A zinc sulfide screen surrounding the apparatus detected alpha particles by scintillation (flashes of light) (1 mark). Observations: (1) Most alpha particles passed straight through the foil with little or no deflection (1 mark). (2) A small fraction (~1 in 8,000) were deflected at angles greater than 90° (1 mark). (3) A very small fraction (~1 in 20,000) were reflected almost straight back (1 mark). Conclusions: (1) Most of the atom is empty space (most particles pass through). The nucleus is tiny, dense, and positively charged — concentrating the repulsive force for near-misses. The deflections increase as alpha particles pass closer to the nucleus; near-direct hits produce back-scatter (1 mark).
Q7 (4 marks): In science, models are tentative explanations consistent with current evidence; they are revised when new evidence cannot be explained (1 mark). Example 1: Thomson's plum pudding model (1904) explained the existence of electrons and overall neutrality of atoms. However, Rutherford's gold foil experiment (1911) produced large-angle deflections of alpha particles — impossible if the positive charge was diffuse (as in the plum pudding). This new evidence necessitated the nuclear model (1 mark). Example 2: Rutherford's nuclear model correctly described the nucleus but could not explain why orbiting electrons didn't spiral inward (classical physics) nor why hydrogen emits discrete spectral lines. Bohr (1913) revised the model by introducing quantised energy levels, which explained both the stability of electrons and the discrete spectral lines (1 mark). Both revisions show that models are not "right or wrong" — they are progressively refined as evidence expands our understanding, always keeping the core ideas that worked while adding new explanatory power (1 mark).
Return to your Think First response. You should now be able to explain why Rutherford's experiment overturned the plum pudding model:
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