C3: History of the Atom
Development of the atomic model from Dalton to Bohr
📋 Key Definitions
Atom: The smallest particle of an element that retains its chemical properties. Atoms consist of a nucleus containing protons and neutrons, with electrons orbiting in shells.
Proton: A subatomic particle found in the nucleus with a relative charge of +1 and a relative mass of 1.
Neutron: A subatomic particle found in the nucleus with a relative charge of 0 and a relative mass of 1.
Electron: A subatomic particle that orbits the nucleus in shells with a relative charge of −1 and a relative mass of approximately 0 (1/1836).
Isotope: Atoms of the same element that have the same number of protons but a different number of neutrons. Isotopes have the same atomic number but different mass numbers.
Relative atomic mass (Aᵣ): The weighted mean mass of an atom of an element relative to 1/12th the mass of an atom of carbon-12.
📊 Subatomic Particle Properties
| Particle | Relative Mass | Relative Charge | Location |
|---|---|---|---|
| Proton | 1 | +1 | Nucleus |
| Neutron | 1 | 0 | Nucleus |
| Electron | ≈ 0 (1/1836) | −1 | Electron shells |
Key rule: Atomic number = number of protons. Mass number = number of protons + neutrons. Number of neutrons = mass number − atomic number. In a neutral atom, number of electrons = number of protons.
🕰️ Development of the Atomic Model
Our understanding of the atom has changed over time as new evidence has been discovered. Each new model replaced or refined the previous one. This is a key example of how scientific theories develop through experimentation and evidence.
| Scientist | Year | Model | Evidence |
|---|---|---|---|
| Dalton | 1803 | Solid sphere — atoms are indivisible solid spheres, different elements have different types of atom | Based on experimental observations of mass ratios in chemical reactions |
| Thomson | 1897 | Plum pudding — atom is a positive sphere with negative electrons embedded within it | Discovered the electron using cathode ray experiments; realised atoms could be divided |
| Rutherford | 1911 | Nuclear model — tiny dense positive nucleus surrounded by a cloud of electrons, mostly empty space | Gold foil experiment: most alpha particles passed through (empty space), some deflected (nucleus), very few bounced back (dense positive centre) |
| Bohr | 1913 | Electron shell model — electrons orbit the nucleus in fixed shells at specific distances, each shell has a fixed energy | Explained why atoms emit light at specific frequencies; mathematical model matched spectral line observations |
| Chadwick | 1932 | Discovered the neutron — neutral particles in the nucleus that explain isotopes | Bombarded beryllium with alpha particles and detected neutral radiation (neutrons) |
Examiners often ask you to explain how evidence led to changes in the atomic model. Always link the specific experimental result to the specific feature of the new model.
🔵 Thomson's Plum Pudding Model
J.J. Thomson discovered the electron in 1897 through cathode ray experiments. He found that cathode rays were negatively charged particles much smaller than atoms, which he called "corpuscles" (now called electrons). This showed that atoms were not indivisible as Dalton had thought.
Thomson proposed the "plum pudding" model: the atom is a sphere of positive charge with negative electrons embedded within it, like plums in a pudding. The overall charge of the atom was neutral because the positive and negative charges balanced.
Why the Plum Pudding Model Was Accepted
The plum pudding model explained two key observations: (1) atoms are neutral overall, because the positive and negative charges balance; (2) atoms can lose or gain electrons to form ions, because the electrons are separate particles within the atom. It was the best model available until Rutherford's experiment provided contradictory evidence.
⚡ Rutherford's Gold Foil Experiment
In 1911, Ernest Rutherford, Hans Geiger and Ernest Marsden conducted the gold foil experiment. They fired alpha particles (positively charged) at a thin sheet of gold foil and observed how they scattered.
The three observations and what they proved:
| Observation | What it showed |
|---|---|
| Most alpha particles passed straight through the foil | Atoms are mostly empty space |
| Some alpha particles were deflected at large angles | The centre of the atom (nucleus) is positively charged, repelling the positive alpha particles |
| A very few alpha particles bounced straight back | The nucleus is very dense and contains most of the atom's mass |
Why the Results Contradicted the Plum Pudding Model
If the plum pudding model were correct, the positive charge would be spread throughout the atom. Alpha particles should have been slightly deflected by the diffuse positive charge, with none bouncing back. The observation that most passed through (empty space) and some bounced back (concentrated positive charge) could only be explained by a tiny, dense, positive nucleus. This completely disproved the plum pudding model.
🔵 Bohr's Electron Shell Model
Rutherford's nuclear model had a problem: according to classical physics, electrons orbiting the nucleus should continuously lose energy and spiral into the nucleus. Niels Bohr refined the model in 1913 by proposing that electrons orbit in fixed shells at specific distances from the nucleus.
Bohr's key ideas: Electrons can only exist in certain fixed energy levels (shells). Electrons can move between shells by absorbing or emitting a specific amount of energy. This explained the line spectra observed for elements — each line corresponds to an electron moving between two specific shells.
Electron Shell Configuration
The first shell holds up to 2 electrons, the second holds up to 8, and the third holds up to 8 (for the first 20 elements). For example:
Sodium (Na, atomic number 11): Electron configuration = 2, 8, 1
Chlorine (Cl, atomic number 17): Electron configuration = 2, 8, 7
Calcium (Ca, atomic number 20): Electron configuration = 2, 8, 8, 2
🔬 Chadwick and the Neutron
By 1932, scientists knew that the nucleus contained protons, but the measured mass of nuclei was greater than the total mass of the protons alone. This suggested there must be another particle in the nucleus. James Chadwick proved the existence of the neutron in 1932 by bombarding beryllium with alpha particles and detecting a previously unknown type of neutral radiation.
Why neutrons matter: Neutrons explain the existence of isotopes — atoms of the same element (same number of protons) can have different numbers of neutrons and therefore different mass numbers. They also help explain nuclear stability and radioactive decay.
🧪 Isotopes
Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons. They have the same atomic number but different mass numbers.
Examples of Isotopes
Hydrogen isotopes:
- Protium (¹H): 1 proton, 0 neutrons, mass number = 1
- Deuterium (²H): 1 proton, 1 neutron, mass number = 2
- Tritium (³H): 1 proton, 2 neutrons, mass number = 3
Carbon isotopes:
- Carbon-12: 6 protons, 6 neutrons
- Carbon-13: 6 protons, 7 neutrons
- Carbon-14: 6 protons, 8 neutrons
Isotopes of the same element have identical chemical properties because they have the same electron configuration. They have slightly different physical properties because of their different masses.
📐 Calculating Relative Atomic Mass
Aᵣ = (sum of: isotope mass × isotope abundance) ÷ 100
The relative atomic mass is a weighted average that takes into account the abundance of each isotope in a naturally occurring sample.
Worked Example - Calculating Relative Atomic Mass
Chlorine has two naturally occurring isotopes: ³⁵Cl (75.8% abundance) and ³⁷Cl (24.2% abundance). Calculate the relative atomic mass of chlorine.
Step 1: Multiply each isotope mass by its abundance:
35 × 75.8 = 2653
37 × 24.2 = 895.4
Step 2: Add these together:
2653 + 895.4 = 3548.4
Step 3: Divide by 100:
Aᵣ = 3548.4 ÷ 100 = 35.5
The relative atomic mass of chlorine is 35.5 (to 1 decimal place).
Worked Example 2 - Calculating Relative Atomic Mass
Copper has two isotopes: ⁶³Cu (69.2% abundance) and ⁶⁵Cu (30.8% abundance). Calculate the relative atomic mass of copper.
Step 1: 63 × 69.2 = 4359.6 and 65 × 30.8 = 2002
Step 2: 4359.6 + 2002 = 6361.6
Step 3: 6361.6 ÷ 100 = 63.6
Aᵣ of copper = 63.6
❓ Practice Questions
Q1: Foundation Describe Rutherford's gold foil experiment and state the three key observations.
Alpha particles were fired at thin gold foil. Three observations: (1) Most passed straight through — atoms are mostly empty space. (2) Some were deflected at large angles — the nucleus is positively charged. (3) A very few bounced back — the nucleus is very dense and contains most of the mass.Q2: Higher Explain how the results of the gold foil experiment disproved the plum pudding model.
The plum pudding model predicted that positive charge was spread throughout the atom, so alpha particles should only be slightly deflected. The observation that most passed through showed the atom is mostly empty space, and the fact that some bounced back showed the positive charge is concentrated in a tiny, dense nucleus — contradicting the diffuse positive charge in the plum pudding model.Q3: Foundation Define the term isotope and give one example.
Isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons. Example: carbon-12 (6 protons, 6 neutrons) and carbon-14 (6 protons, 8 neutrons) are isotopes of carbon.Q4: Higher Magnesium has three isotopes: ²⁴Mg (78.6%), ²⁵Mg (10.1%) and ²⁶Mg (11.3%). Calculate the relative atomic mass of magnesium.
Aᵣ = (24 × 78.6 + 25 × 10.1 + 26 × 11.3) ÷ 100 = (1886.4 + 252.5 + 293.8) ÷ 100 = 2432.7 ÷ 100 = 24.3. The relative atomic mass of magnesium is 24.3.Q5: Foundation State the relative mass and relative charge of a proton, neutron and electron.
Proton: mass = 1, charge = +1. Neutron: mass = 1, charge = 0. Electron: mass ≈ 0, charge = −1. Protons and neutrons are found in the nucleus; electrons orbit in shells.🎯 Exam Tips
When describing Rutherford's experiment, always link each observation to its conclusion. Three observations, three conclusions — examiners award marks for each pair.
For isotope questions, remember that isotopes have identical chemical properties (same electron configuration) but different physical properties (different mass). Examiners often ask this distinction.
In relative atomic mass calculations, show all working clearly: multiply each mass by its percentage, add them up, then divide by 100. Rounding errors lose marks — keep intermediate values to at least one decimal place.
When asked how the atomic model has changed over time, structure your answer chronologically and always explain WHY each model was accepted or rejected based on evidence.
🧮 Maths Skills
Calculating Relative Atomic Mass from Isotope Abundances
Formula: Aᵣ = (sum of: isotope mass number × percentage abundance) ÷ 100. This gives a weighted mean that accounts for how common each isotope is in a naturally occurring sample.
Worked example: Chlorine has two isotopes: ³⁵Cl (75.8%) and ³⁷Cl (24.2%). Aᵣ = (35 × 75.8 + 37 × 24.2) ÷ 100 = (2653 + 895.4) ÷ 100 = 35.5. Always show each step clearly in the exam.
Working with three or more isotopes: Magnesium has three isotopes: ²⁴Mg (78.6%), ²⁵Mg (10.1%) and ²⁶Mg (11.3%). Aᵣ = (24 × 78.6 + 25 × 10.1 + 26 × 11.3) ÷ 100 = (1886.4 + 252.5 + 293.8) ÷ 100 = 24.3. Extend the same method for any number of isotopes.
Reverse calculations: If you know Aᵣ and all but one isotope abundance, you can work backwards. E.g. if Aᵣ = 35.5 and ³⁵Cl is 75.8%, then ³⁷Cl must be 100 − 75.8 = 24.2%. The abundances always sum to 100%.
Rounding and significant figures: Give Aᵣ to the same number of decimal places as the data provided. If abundances are given to 1 d.p., give your answer to 1 d.p. Do not round intermediate values — only round the final answer. Carry all figures through the calculation and round only at the end.
Standard form: The mass of an electron is approximately 9.11 × 10⁻³¹ kg. Using standard form avoids writing many zeros and makes comparison of subatomic particle masses clearer.
❌ Common Misconceptions
Misconceptions About the Atomic Model
Wrong: Rutherford discovered the electron
Correct: Thomson discovered the electron in 1897 using cathode ray experiments. Rutherford discovered the nucleus in 1911 using the gold foil experiment. Chadwick discovered the neutron in 1932. Each scientist made a different key discovery.
Wrong: The current atomic model is definitely correct and will not change
Correct: Scientific models are always being refined as new evidence is discovered. The current model is our best explanation so far, but it may still be improved. Bohr's model was refined by the discovery of sub-shells and quantum mechanics. Models are representations, not perfect descriptions of reality.
Wrong: Isotopes have different chemical properties because they have different numbers of neutrons
Correct: Isotopes have identical chemical properties because they have the same electron configuration. Only physical properties (e.g. density, rate of diffusion) differ slightly due to the different mass.
✍️ 6-Mark Extended Question
Question
Describe how the model of the atom has changed over time. Explain the evidence that led to each change.
Dalton (1803): Proposed that atoms are indivisible solid spheres, different elements have different atoms. Based on observations of mass ratios in chemical reactions [1 mark].
Thomson (1897): Discovered the electron using cathode ray experiments, showing atoms could be divided. Proposed the plum pudding model — a sphere of positive charge with electrons embedded within it. Evidence: cathode rays were deflected by electric and magnetic fields, proving they were negatively charged particles smaller than atoms [1 mark].
Rutherford (1911): Gold foil experiment showed most alpha particles passed through (atoms are mostly empty space), some were deflected (nucleus is positive) and a few bounced back (nucleus is dense). This disproved the plum pudding model because the positive charge was concentrated in a tiny nucleus, not spread throughout the atom. Proposed the nuclear model [1 mark].
Bohr (1913): Refined the model by proposing electrons orbit in fixed shells at specific distances. Evidence: atoms emit light at specific frequencies (line spectra), which matches electrons moving between fixed energy levels. Explained why electrons do not spiral into the nucleus [1 mark].
Chadwick (1932): Discovered the neutron by bombarding beryllium with alpha particles and detecting neutral radiation. Evidence: nuclei were heavier than the total mass of protons alone, suggesting another particle. Explained isotopes — same element with different numbers of neutrons [1 mark].
Mark scheme: 1 mark per scientist for model description, 1 mark per scientist for evidence. Max 6 marks from any combination covering at least three scientists with both model and evidence.🔍 AO3: Analyse and Evaluate
Calculating Ar from Isotope Data and Comparing to Published Values
Scenario: A student looks up the isotopes of copper and finds: ⁶³Cu (69.17% abundance) and ⁶⁵Cu (30.83% abundance). The published Aᵣ of copper is 63.546. Calculate the Aᵣ from the isotope data and comment on how well it matches the published value.
Calculation: Aᵣ = (63 × 69.17 + 65 × 30.83) ÷ 100 = (4357.71 + 2003.95) ÷ 100 = 63.6 (to 1 d.p.). The calculated value of 63.6 is very close to the published value of 63.546. The small difference is because the isotope masses are not exactly whole numbers — ⁶³Cu has a mass of 62.9296 and ⁶⁵Cu has a mass of 64.9278. Using more precise isotope masses gives a result closer to the published value.
Evaluation: The calculation using whole-number mass numbers gives an approximation. For more accurate results, the exact isotope masses should be used. The student should recognise that mass numbers are integers (protons + neutrons) but actual isotope masses are slightly different due to the mass-energy equivalence (binding energy). This explains why Aᵣ values on the Periodic Table are often not whole numbers.
Evaluating atomic models: Each model had strengths and weaknesses. Dalton's model explained conservation of mass but could not explain ions or electricity. Thomson's model explained neutral atoms and ions but was disproved by Rutherford's experiment. Rutherford's model explained the nucleus but could not explain why electrons did not spiral in. Bohr's model explained line spectra but could not explain heavier elements fully. Scientific models are refined, not replaced — each builds on the previous one.
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