Atom Structure


Dalton’s Atomic Theory

Key ideas:

  • Matter is made of tiny, indivisible particles called atoms

  • Atoms of the same element are identical

  • Atoms combine in simple whole-number ratios

  • Atoms are rearranged, not created or destroyed, in reactions

Limitations:

  • Atoms are divisible (subatomic particles exist)

  • Atoms of the same element can differ (isotopes)

📌 Exam tip: Be able to state + critique Dalton’s theory.


2. Discovery of Subatomic Particles

Electron

  • Discovered by J.J. Thomson

  • Cathode ray tube experiment

  • Charge: –1

  • Very small mass (≈ 1/1836 of a proton)

Proton

  • Discovered by Rutherford

  • Charge: +1

  • Relative mass: 1

Neutron

  • Discovered by Chadwick

  • Charge: 0

  • Relative mass: 1

📊 Summary Table

Particle

Charge

Relative Mass

Location

Proton

+1

1

Nucleus

Neutron

0

1

Nucleus

Electron

–1

1/1836

Outside nucleus

3. Atomic Models (Very Exam-Heavy)

Thomson’s Model

  • “Plum pudding” model

  • Electrons embedded in a positive sphere

  • ❌ Could not explain scattering results

Rutherford’s Nuclear Model

Gold foil experiment:

  • Most α-particles passed straight through

  • Some deflected

  • Few rebounded

Conclusions:

  • Atom is mostly empty space

  • Dense, positively charged nucleus

Limitations:

  • Could not explain atomic stability or line spectra

Bohr’s Model

Key ideas:

  • Electrons move in fixed energy levels (shells)

  • Energy is absorbed/emitted when electrons move between levels

Energy equation:

ΔE=hf\Delta E = hfΔE=hf

Limitations:

  • Works mainly for hydrogen

  • Fails for multi-electron atoms

📌 Exam tip: Know experiments, observations, conclusions, and limitations.

4. Atomic Number, Mass Number & Isotopes

Definitions

  • Atomic number (Z): Number of protons

  • Mass number (A): Protons + neutrons

Neutrons=A−Z\text{Neutrons} = A - ZNeutrons=A−Z

Isotopes

  • Same number of protons

  • Different number of neutrons

  • Same chemical properties, different physical properties

Examples:

  • Carbon-12, Carbon-13, Carbon-14

Uses:

  • Carbon-14 dating

  • Medical tracers

  • Nuclear energy

5. Relative Atomic Mass (Ar)

Definition

Weighted average mass of an element’s isotopes compared to 1/12 of carbon-12.

Calculation

Ar=∑(isotopic mass×abundance)100A_r = \frac{\sum (\text{isotopic mass} \times \text{abundance})}{100}Ar​=100∑(isotopic mass×abundance)​

📌 Exam tip: Expect full calculations with working and units.

6. Electromagnetic Radiation

Key Terms

  • Wavelength (λ)

  • Frequency (f)

  • Speed of light (c = 3.0 × 10⁸ m s⁻¹)

Relationship:

c=fλc = f\lambdac=fλ

Energy of radiation:

E=hfE = hfE=hf

Where:

  • h=6.63×10−34 J sh = 6.63 \times 10^{-34} \, \text{J s}h=6.63×10−34J s

7. Atomic Emission Spectra

  • Produced when excited electrons fall to lower energy levels

  • Each element has a unique line spectrum

  • Evidence for quantised energy levels

📌 Exam favourite: Link line spectra → Bohr model → energy levels.

8. Electron Arrangement

Shells & Sub-shells

  • Shells: n = 1, 2, 3, 4…

  • Sub-shells: s, p, d, f

Orbital Capacities

  • s: 2 electrons

  • p: 6 electrons

  • d: 10 electrons

  • f: 14 electrons

Filling Order

Aufbau principle:
1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p …

Rules:

  • Pauli Exclusion Principle

  • Hund’s Rule

9. Ionisation Energy (Intro Link)

Definition:

  • Energy required to remove one mole of electrons from one mole of gaseous atoms to form 1+ ions

Trends (basic):

  • Increases across a period

  • Decreases down a group

Reasons:

  • Nuclear charge

  • Atomic radius

  • Shielding effect