2-cie-atomic-structure

Atomic Structure

Particle: Proton
- Position: Nucleus
- Relative Mass: 1
- Relative Charge: +1
Particle: Neutron
- Position: Nucleus
- Relative Mass: 1
- Relative Charge: 0
Particle: Electron
- Position: Orbitals
- Relative Mass: 1/1840
- Relative Charge: -1

Example: Lithium (Li)
- Atomic Number (Protons, Z): 3
- Nucleon (Mass) Number (Protons + Neutrons, A): 7
- Formula to Find Neutrons:
  - Number of neutrons = A - Z
  - Example: 7 - 3 = 4 neutrons

Definition: Atoms with the same number of protons but differing numbers of neutrons.
Chemical Properties: Similar due to identical electronic structure, resulting in similar reactivity and bonding behavior.
Physical Properties: May vary due to differing masses, affecting properties like boiling point and density.

Neutron: No charge, unaffected by electric field.
Proton: Attracted to the negative plate but deflected less due to larger mass.
Electron: Attracted to the positive plate, moving with significant acceleration due to its small mass.

Models of the Atom:
- Early Model: Bohr Model, introduced quantized orbits.
- Electron Shells:
  - 2 electrons in the first shell, 8 in the second, following the 2n² rule.
- Principle Energy Levels: Numbered 1, 2, 3, 4... (with 1 closest to the nucleus)
- Sub-Energy Levels:
  - s: 2 electrons
  - p: 6 electrons
  - d: 10 electrons
  - f: 14 electrons
- Orbitals hold up to 2 electrons of opposite spin
- Shapes of Orbitals:
  - s: spherical
  - p: dumbbell shape

Filling Order: Electrons fill sub-shells in order of increasing energy
Sequence: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p
Example: Calcium (Ca) Configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²
Spin Diagrams: Important for understanding orbital occupation, showing electron arrangements and spin orientations.

Formation of Ions:
- Positive Ion: Electrons are lost from the outermost shell.
  - Example: Mg → Mg²⁺ (losing 2 electrons)
- Negative Ion: Electrons are gained.
  - Example: O → O²⁻ (gaining 2 electrons)
Example Electronic Configurations of Elements:
- Sc: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹
- Fe: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶
- Cu: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d¹⁰
Important Note: 4s electrons are lost before 3d when forming ions.

First Ionisation Energy (IE):
- Definition: Energy required to remove one mole of gaseous electrons.
- Generic Equation: H(g) → H⁺(g) + e⁻
- Nature: All ionisation energies are positive and endothermic.
Second Ionisation Energy (IE):
- Definition: Energy required when one mole of +1 ions forms +2 ions.
- Equation: Ti⁺(g) → Ti²⁺(g) + e⁻
Factors Affecting Ionisation Energy:
- Attraction of the nucleus: More protons = greater attraction for electrons.
- Distance of electrons from the nucleus: Larger atoms have weaker attraction due to increased distance.
- Shielding effect: Outer electrons are repelled by complete inner shells, reducing effective nuclear charge felt by outer electrons.
Trends in Successive IEs:
- Always larger due to increased attraction post positive ion formation, indicated by jumps in energy.
- Patterns in Ionisation Energy Graph Trends:
  - Large jumps indicate change in electron shell.
  - Example: Group 2 to Group 3.
  - Influence of electron energy levels on ionisation.
  - Noble gases peak every period with decreasing ionisation down groups.

First Electron Affinity:
- Change when 1 mole of gaseous atoms gains an electron.
- Exothermic for elements forming negative ions: O(g) + e⁻ → O⁻(g)
Second Electron Affinity:
- Change when 1 mole of gaseous 1- ions gains an electron: O⁻(g) + e⁻ → O²⁻(g)
- Endothermic due to repulsion between electrons.