# Development of the Atomic Model
Our understanding of the atom has evolved dramatically over the past 200 years. From Dalton's simple solid sphere to the modern quantum mechanical model, each new discovery has refined our picture of this tiny particle. For GCSE Chemistry, you need to understand the key models, the evidence that led to each change, and how electrons are arranged in atoms.
1. Timeline of Atomic Models
1.1 Dalton's Model (early 1800s)
John Dalton proposed that:
- All matter is made of tiny, indivisible particles called atoms
- Atoms of the same element are identical
- Atoms of different elements have different masses
- Atoms combine in simple whole-number ratios to form compounds
- Atoms cannot be created, destroyed, or divided
Dalton imagined atoms as solid spheres — like tiny billiard balls.
Limitation: Dalton didn't know about subatomic particles (protons, neutrons, electrons).
1.2 Thomson's "Plum Pudding" Model (1897)
J.J. Thomson discovered the electron using cathode ray tubes. He showed that electrons were negatively charged particles much smaller than atoms.
Thomson proposed the plum pudding model:
- The atom is a sphere of positive charge
- Negative electrons are embedded throughout, like plums in a pudding
- The positive and negative charges balance, making the atom neutral
1.3 Rutherford's Nuclear Model (1911)
Ernest Rutherford directed his students Geiger and Marsden to fire alpha particles () at a thin gold foil.
Expected results (plum pudding): All alpha particles should pass straight through with minimal deflection.
Actual results:
- Most alpha particles passed straight through
- Some were deflected at small angles
- A very few bounced back (deflected > 90°)
Conclusions:
- Most of the atom is empty space (most particles passed through)
- There is a small, dense, positive nucleus at the centre (particles bounced back)
- Electrons orbit the nucleus at a distance (like planets around the Sun)
This was the nuclear model of the atom.
1.4 Bohr's Model (1913)
Niels Bohr improved Rutherford's model by proposing that:
- Electrons orbit the nucleus in fixed energy levels (shells)
- Electrons can only exist in specific orbits, not anywhere
- Electrons can jump between energy levels by absorbing or emitting energy
- Each energy level can hold a maximum number of electrons
The evidence came from atomic emission spectra — when elements are heated, they emit light at specific wavelengths (line spectra), not a continuous spectrum. This suggested that energy changes in atoms are quantised (come in fixed amounts).
1.5 Chadwick's Discovery (1932)
James Chadwick discovered the neutron — a particle with no charge and a mass of 1 (similar to the proton). This explained why atoms were heavier than expected based on protons alone.
2. Summary of Atomic Models
| Scientist | Date | Model | Key Discovery |
|---|---|---|---|
| Dalton | ~1803 | Solid sphere | Atoms as indivisible particles |
| Thomson | 1897 | Plum pudding | Discovery of electrons |
| Rutherford | 1911 | Nuclear model | Small, dense, positive nucleus |
| Bohr | 1913 | Electron shells | Electrons in fixed energy levels |
| Chadwick | 1932 | Modern nuclear | Discovery of neutrons |
3. Electron Configuration
Electrons are arranged in energy levels (shells) around the nucleus. Each shell can hold a maximum number of electrons:
| Shell | Maximum Electrons |
|---|---|
| 1st (closest to nucleus) | 2 |
| 2nd | 8 |
| 3rd | 8 (at GCSE level) |
Rules for Filling Shells
- Fill the innermost shell first
- When a shell is full, start filling the next one
- The number of electrons equals the number of protons (for neutral atoms)
Examples
| Element | Atomic Number | Electron Configuration | Notation |
|---|---|---|---|
| Hydrogen | 1 | 1 | 1 |
| Helium | 2 | 2 | 2 |
| Lithium | 3 | 2, 1 | 2.1 |
| Carbon | 6 | 2, 4 | 2.4 |
| Oxygen | 8 | 2, 6 | 2.6 |
| Neon | 10 | 2, 8 | 2.8 |
| Sodium | 11 | 2, 8, 1 | 2.8.1 |
| Chlorine | 17 | 2, 8, 7 | 2.8.7 |
| Calcium | 20 | 2, 8, 8, 2 | 2.8.8.2 |
Electron Configuration and the Periodic Table
- The group number tells you the number of electrons in the outer shell
- Group 1: 1 outer electron
- Group 7: 7 outer electrons
- Group 0: full outer shell (8, except He with 2)
- The period number tells you the number of occupied shells
- Period 1: 1 shell occupied
- Period 3: 3 shells occupied
4. The Importance of Outer Electrons
The electrons in the outermost shell (valence electrons) determine an element's:
- Chemical reactivity — how easily it reacts
- Bonding behaviour — ionic, covalent, or metallic
- Position in the periodic table
Atoms are most stable when their outer shell is full. This explains why:
- Noble gases (Group 0) are unreactive — they already have full outer shells
- Metals tend to lose electrons to achieve a full outer shell
- Non-metals tend to gain or share electrons
Worked Example: Drawing Electron Configurations
Question: Draw the electron configuration of a sodium atom (atomic number 11).
- Total electrons = 11
- 1st shell: 2 electrons
- 2nd shell: 8 electrons
- 3rd shell: 1 electron
- Configuration: 2, 8, 1
Worked Example: Identifying the Model
Question: Which atomic model proposed that electrons are embedded in a sphere of positive charge?
Thomson's plum pudding model (1897).
Worked Example: Alpha Particle Scattering
Question: In Rutherford's experiment, most alpha particles passed straight through the gold foil. What does this tell us about the structure of the atom?
This shows that the atom is mostly empty space. The nucleus is very small compared to the overall size of the atom, so most particles miss it entirely.
6. Practice Questions
- Place these atomic models in chronological order: Bohr, Dalton, Rutherford, Thomson.
- Describe the key features of the plum pudding model.
- Explain why a small number of alpha particles bounced back in Rutherford's experiment.
- Write the electron configuration of: (a) nitrogen, (b) argon, (c) potassium.
- How did Bohr's model differ from Rutherford's model?
Want to check your answers and get step-by-step solutions?
7. Common Misconceptions
| Misconception | Reality |
|---|---|
| Electrons orbit like planets in neat circles | Electrons exist in energy levels but their exact position is uncertain |
| Rutherford's experiment used atoms | Rutherford fired alpha particles at gold foil (many layers of atoms) |
| The plum pudding model was wrong and useless | It was an important step that explained the existence of electrons |
| Atoms are solid | Atoms are mostly empty space |
| Dalton's model is completely wrong | Many of Dalton's ideas are still valid (atoms of same element are similar, combine in ratios) |
8. Exam Tips
- When describing the alpha scattering experiment, state the observation and then the conclusion separately
- Learn the maximum electrons per shell: 2, 8, 8 (for GCSE)
- Know the link between electron configuration and position in the periodic table
- Questions often ask you to compare two models — state what is similar and what is different
- You don't need to know the quantum model for GCSE, but understanding that science progresses through evidence is important
Frequently Asked Questions
Why did scientists change their model of the atom?
Each new model was based on new experimental evidence. When experiments produced results that couldn't be explained by the current model, a new model was developed. This is how science works — models are refined as new evidence emerges.
Is the Bohr model the final model?
No. The modern quantum mechanical model is more accurate, but Bohr's model is a good approximation for GCSE-level chemistry. The quantum model describes electrons in terms of probability clouds rather than fixed orbits.
Why does the 3rd shell hold only 8 electrons at GCSE but 18 at A-Level?
At GCSE, you only need to consider the first three shells with the 2, 8, 8 rule. At A-Level, you learn about sub-shells (s, p, d, f) which allow the 3rd shell to hold up to 18 electrons.
Summary
- The atomic model evolved: Dalton → Thomson → Rutherford → Bohr → Chadwick
- Each new model was supported by experimental evidence
- Rutherford's alpha scattering proved the existence of a small, dense, positive nucleus
- Bohr introduced fixed energy levels for electrons
- Electron configuration follows the 2, 8, 8 pattern
- The number of outer electrons determines an element's group and chemical properties