AQA Chemistry A-Level - Atomic Structure

3.1.1: Atomic Structure

Fundamental Particles

• Evolution of Atomic Models:

  • Plum Pudding Model: Atoms thought to be a sphere of positive charge with small embedded negative charges.
  • Electron Shell Model: Discovered in 1911 from Rutherford scattering experiment.
    • Atoms consist of a dense nucleus with protons and neutrons.
    • Electrons orbit in shells.
    • Nucleus is positively charged and contains most of the atom's mass.
    • Neutral atoms have equal numbers of protons and electrons.

• Particles and Properties:

  Particle	Relative Charge	Relative Mass
  Proton	+1	1
  Neutron	0	1
  Electron	-1	1/1840

• Electron Shells:

  • Maximum electrons in a shell: calculated by (2n^2) where (n) is the shell number.
  • Example: Shell 2 can hold 8 electrons. Shells must fill in order.

Mass Number and Isotopes

• Definitions:

  • Mass Number (A): Sum of protons and neutrons.
  • Atomic Number (Z): Number of protons.
  • Relative Atomic Mass (Ar): Mean mass of an atom relative to 1/12 of carbon-12 isotope.

• Isotopes:

  • Atoms with the same atomic number but different neutron numbers.
  • Chemically similar due to same proton number and electron configuration.
  • Different physical properties due to varying mass numbers.
  • Example: Hydrogen isotopes (Protium, Deuterium, Tritium).

• Ions:

  • Formed by gaining or losing electrons, resulting in non-neutral charge.

Mass Spectrometry

• Purpose: Identifies isotopes and determines relative atomic mass.

• Time of Flight (TOF) Mass Spectrometry:

  1. Ionisation: Sample vaporized, high voltage removes electrons, forming +1 ions.
  2. Acceleration: Ions accelerated towards negative detection plate.
  3. Ion Drift: Magnetic field deflects ions based on mass and charge.
  4. Detection: Ions gain electrons at detection plate, generating charge flow.
  5. Analysis: Flight times and current values produce spectra, showing isotope abundances.

• Example Calculation: Mass to charge ratio (m/z) can be halved for 2+ ions.

• Chlorine Spectra: Displays distinct patterns due to isotope abundance differences.

Electron Configuration

• Orbitals and Electron Filling:

  • Orbitals (s, p, d, f) correlate with periodic table blocks.
  • Max electrons: s = 2, p = 6, d = 10.
  • Filled in order of increasing energy (s -> p -> d).

• Rules for Electron Configuration:

  1. Fill lowest energy orbital first.
  2. Electrons with same spin fill first before pairing.
  3. Maximum 2 electrons with opposite spin per orbital.

• Spin: Electrons in same orbital have opposite spins to enhance stability.

• Exceptions: Unpaired spins cause instability; may rearrange for stability.

Ionisation Energy

• Definition: Minimum energy to remove one mole of electrons from gaseous atoms, measured in kJ/mol.

• Trends:

  • Across Period: Ionisation energy increases (decreasing atomic radius, stronger attraction).
  • Down Group: Ionisation energy decreases (increasing atomic radius, more shielding).

• Successive Ionisation Energies:

  • Larger energy required for electrons closer to nucleus.
  • Sudden energy increase indicates change in energy level, supporting atomic orbital theory.

• Examples:

  • Aluminium's first ionisation energy lower due to electron repulsion.
  • Graphs show large increases signaling orbital changes.