C3: History of the Atom
Development of the atomic model
C3 History Of The Atom
📋 Key Definitions
Atomic model: A representation of the structure of an atom. Our understanding of the atom has changed over time as new evidence was discovered.
Nucleus: The tiny, dense centre of an atom containing protons and neutrons. It has a positive charge due to the protons.
Electron shell: A fixed energy level at a set distance from the nucleus in which electrons orbit the nucleus.
Isotope: Atoms of the same element that have the same number of protons but a different number of neutrons. They have the same atomic number but different mass numbers.
Electron configuration: The arrangement of electrons in shells around the nucleus, written as numbers (e.g. 2,8,1 for sodium).
📜 Development of the Atomic Model
Dalton (early 1800s) - Solid Spheres
John Dalton proposed that atoms were tiny, indivisible solid spheres. He suggested each element was made of a different type of sphere.
Dalton's Model
Atoms are like tiny billiard balls - solid spheres that cannot be divided. Different elements had different sized spheres. This was the first modern atomic theory.
Thomson (1897) - Plum Pudding Model
J.J. Thomson discovered the electron. He realised atoms were not solid spheres but contained even smaller negatively charged particles (electrons). He proposed the plum pudding model.
Plum pudding model: An atom is a positive sphere with negative electrons embedded in it, like plums in a plum pudding. The overall atom is neutral because the positive charge balances the negative electrons.
Thomson's Model
Imagine a spherical pudding with the positive charge spread throughout, and small negative electrons scattered inside like fruit in a Christmas pudding. The atom is overall neutral.
Rutherford (1909-1911) - Nuclear Model
Ernest Rutherford's team (Geiger and Marsden) carried out the alpha scattering experiment, which disproved the plum pudding model.
Alpha Scattering Experiment
A beam of alpha particles (positively charged) was fired at thin gold foil. Most passed straight through, some were deflected at large angles, and a very few bounced straight back.
- Most passed through: Atoms are mostly empty space
- Some deflected: The nucleus is positively charged (repelling the positive alpha particles)
- Very few bounced back: The nucleus is tiny but very dense and massive
Rutherford's nuclear model: The atom has a tiny, positive, dense nucleus at the centre, with electrons orbiting around it. The atom is mostly empty space.
Bohr (1913) - Electron Shells
Niels Bohr refined Rutherford's model by proposing that electrons orbit the nucleus in fixed energy levels (shells) at specific distances. Electrons can move between shells but cannot exist between them.
Bohr's Model
Rutherford's model could not explain why atoms were stable - orbiting electrons should lose energy and spiral into the nucleus. Bohr solved this by proposing that electrons can only exist in fixed shells at set distances, like planets orbiting the Sun at fixed distances.
Later Discoveries
Later experiments showed that the nucleus itself contains smaller particles:
- Protons were discovered by Rutherford (1919) - positive charge, found in the nucleus
- Neutrons were discovered by Chadwick (1932) - no charge, found in the nucleus
| Scientist | Year | Model | Key Idea |
|---|---|---|---|
| Dalton | ~1803 | Solid spheres | Atoms are tiny, indivisible solid spheres |
| Thomson | 1897 | Plum pudding | Positive sphere with electrons embedded |
| Rutherford | 1911 | Nuclear model | Tiny positive nucleus, electrons orbit, mostly empty space |
| Bohr | 1913 | Electron shells | Electrons in fixed energy levels/shells |
| Chadwick | 1932 | Neutrons added | Nucleus contains protons and neutrons |
📏 Size of the Atom
Size of an atom: The radius of a typical atom is about 1 × 10⁻¹⁰ m (0.1 nm). The radius of the nucleus is about 1 × 10⁻¹⁵ m - roughly 100,000 times smaller than the atom.
If an atom were the size of a football stadium, the nucleus would be the size of a pea at the centre. This illustrates how atoms are mostly empty space.
⚛️ Subatomic Particles
| Particle | Relative Charge | Relative Mass | Location |
|---|---|---|---|
| Proton | +1 | 1 | Nucleus |
| Neutron | 0 | 1 | Nucleus |
| Electron | -1 | ~0 (0.0005) | Orbiting in shells |
Key relationships: Atomic number = number of protons. Mass number = number of protons + number of neutrons. In a neutral atom, number of protons = number of electrons.
Number of neutrons = mass number − atomic number
Worked Example - Calculating Subatomic Particles
Lithium has atomic number 3 and mass number 7. How many protons, neutrons and electrons does it have?
Protons = atomic number = 3
Electrons = protons (neutral atom) = 3
Neutrons = mass number − atomic number = 7 − 3 = 4
So lithium has 3 protons, 4 neutrons and 3 electrons.
Worked Example - Chlorine
Chlorine has atomic number 17 and mass number 35. Calculate the subatomic particles.
Protons = 17
Electrons = 17
Neutrons = 35 − 17 = 18
So chlorine has 17 protons, 18 neutrons and 17 electrons.
🔄 Electron Configuration
Rules for electron shells: The first shell holds up to 2 electrons. The second and third shells each hold up to 8 electrons. Electrons fill the shells starting with the one closest to the nucleus.
Electron configuration for first 20 elements: 2, 8, 8, 2
Worked Example - Electron Configurations
- Hydrogen (1): 1
- Helium (2): 2
- Lithium (3): 2,1
- Carbon (6): 2,4
- Neon (10): 2,8
- Sodium (11): 2,8,1
- Argon (18): 2,8,8
- Calcium (20): 2,8,8,2
The number of electrons in the outer shell determines the group number in the periodic table. The number of occupied shells determines the period number.
Worked Example - Electron Configuration and the Periodic Table
Sodium has electron configuration 2,8,1.
Outer shell has 1 electron → Group 1
3 occupied shells → Period 3
Chlorine has electron configuration 2,8,7.
Outer shell has 7 electrons → Group 7
3 occupied shells → Period 3
🧬 Isotopes
Isotopes: Atoms of the same element with the same number of protons but a different number of neutrons. Isotopes have the same atomic number but different mass numbers.
Worked Example - Isotopes of Chlorine
Chlorine has two isotopes:
- Chlorine-35: 17 protons, 18 neutrons, 17 electrons (⁴⁵Cl)
- Chlorine-37: 17 protons, 20 neutrons, 17 electrons (³⁷Cl)
Both have atomic number 17 and the same electron configuration (2,8,7), so they have identical chemical properties. They differ only in mass.
More Isotope Examples
- Carbon-12: 6p, 6n, 6e and Carbon-14: 6p, 8n, 6e
- Hydrogen-1 (protium): 1p, 0n, 1e and Hydrogen-2 (deuterium): 1p, 1n, 1e and Hydrogen-3 (tritium): 1p, 2n, 1e
- Uranium-235: 92p, 143n, 92e and Uranium-238: 92p, 146n, 92e
Isotopes of the same element have identical chemical properties because they have the same electron configuration. Their physical properties differ slightly because they have different masses.
Relative atomic mass = (sum of: isotope mass × its abundance) ÷ 100
Worked Example - Calculating Relative Atomic Mass
Chlorine has two isotopes: 75% Cl-35 and 25% Cl-37. Calculate the relative atomic mass.
Ar = (35 × 75 + 37 × 25) ÷ 100 = (2625 + 925) ÷ 100 = 3550 ÷ 100 = 35.5
The relative atomic mass of chlorine is 35.5.
❓ Practice Questions
Q1: Foundation Describe the plum pudding model and explain how the alpha scattering experiment disproved it.
Q2: Foundation A boron atom has atomic number 5 and mass number 11. Calculate the number of protons, neutrons and electrons.
Q3: Foundation Write the electron configuration for: (a) nitrogen (atomic number 7), (b) magnesium (atomic number 12), (c) potassium (atomic number 19).
Q4: Higher Define the term isotope. Explain why isotopes of the same element have identical chemical properties but different physical properties.
Q5: Higher Magnesium has three isotopes: 79% Mg-24, 10% Mg-25 and 11% Mg-26. Calculate the relative atomic mass of magnesium. Give your answer to one decimal place.
Q6: Foundation Describe how the model of the atom has changed from Dalton's model to Bohr's model. Name the key scientist for each stage.
✅ Answers
- The plum pudding model proposed that an atom is a positive sphere with negative electrons embedded in it (like plums in a pudding). The alpha scattering experiment disproved it because: (1) most alpha particles passed straight through - showing atoms are mostly empty space, not a solid positive sphere; (2) some were deflected at large angles - showing a concentrated positive charge (nucleus) rather than a spread-out one; (3) a few bounced back - showing the nucleus is very small and dense.
- Protons = atomic number = 5. Electrons = protons (neutral atom) = 5. Neutrons = mass number − atomic number = 11 − 5 = 6.
- (a) Nitrogen (7): 2,5. (b) Magnesium (12): 2,8,2. (c) Potassium (19): 2,8,8,1.
- Isotopes are atoms of the same element with the same number of protons but a different number of neutrons. They have identical chemical properties because they have the same electron configuration (same number and arrangement of electrons). They have different physical properties because they have different masses (different numbers of neutrons).
- Ar = (24 × 79 + 25 × 10 + 26 × 11) ÷ 100 = (1896 + 250 + 286) ÷ 100 = 2432 ÷ 100 = 24.3.
- Dalton (~1803) proposed atoms as tiny, indivisible solid spheres. Thomson (1897) discovered electrons and proposed the plum pudding model - a positive sphere with negative electrons embedded in it. Rutherford (1911) discovered the nucleus through the alpha scattering experiment and proposed the nuclear model - a tiny, dense, positive nucleus with electrons orbiting in mostly empty space. Bohr (1913) refined this by proposing that electrons orbit in fixed energy levels (shells) at specific distances from the nucleus.
🎯 Exam Tips
- When describing Rutherford's experiment, always link each observation to the conclusion (e.g. "most passed through → atoms are mostly empty space")
- The phrase "mostly empty space" is key for Rutherford's model - examiners look for this specifically
- Electron configuration: remember the first shell only holds 2 electrons, then 8, 8, 2 for the first 20 elements
- For isotopes, always state that chemical properties are the same because the electron configuration is the same
- When calculating neutrons, students often forget to subtract: neutrons = mass number − atomic number
- Relative atomic mass calculations: multiply each isotope mass by its percentage, add them up, then divide by 100
🔢 Maths Skills
Mathematical Skills
Calculating relative atomic mass from isotope abundances: Multiply each isotope mass by its percentage abundance, add the results, then divide by 100.
Example: Copper has two isotopes: 69% Cu-63 and 31% Cu-65. Ar = (63 × 69 + 65 × 31) ÷ 100 = (4347 + 2015) ÷ 100 = 63.6.
Calculating neutrons: Neutrons = mass number − atomic number.
⚠️ Common Misconceptions
Watch Out!
Rutherford discovered the electron. Wrong: Rutherford discovered the electron Correct: Thomson discovered the electron in 1897; Rutherford discovered the nucleus in 1911
The nuclear model was immediately accepted by the scientific community. Wrong: the nuclear model was immediately accepted Correct: it took time to be accepted because the plum pudding model was well-established and new evidence had to be evaluated
✍️ 6-Mark Question
Extended Answer
6 marks: Describe how the atomic model has changed over time. Name the key scientist at each stage.
Dalton (early 1800s) proposed atoms as tiny, indivisible solid spheres. Thomson (1897) discovered electrons and proposed the plum pudding model — a positive sphere with negative electrons embedded in it. Rutherford (1911) discovered the nucleus through the alpha scattering experiment; his nuclear model had a tiny, dense, positive nucleus with electrons orbiting in mostly empty space. Bohr (1913) refined this by proposing that electrons orbit in fixed energy levels (shells) at specific distances. Chadwick (1932) discovered neutrons in the nucleus. Each change was driven by new experimental evidence that the previous model could not explain.
Mark scheme: 1 mark per scientist with correct model (up to 5); 1 mark for stating that new evidence led to changes.
📊 AO3: Analyse & Evaluate
Analysis and Evaluation
A sample of boron contains two isotopes: 20% B-10 and 80% B-11.
Question: Calculate the relative atomic mass of boron. Give your answer to 1 decimal place. Explain why the value is not a whole number.
Answer: Ar = (10 × 20 + 11 × 80) ÷ 100 = (200 + 880) ÷ 100 = 10.8. The value is not a whole number because it is a weighted average of the different isotope masses, reflecting the fact that boron exists as a mixture of two isotopes with different mass numbers.
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