CH2: Electrons In Atoms

What is the charge of a single electron?

1.6 × 10^-19 C

Isotopes definition

Isotopes are atoms of the same element with different mass numbers.

Ions definition

Ions are charged particles formed by the loss/gain of electrons from an atom or group of covalently bonded atoms.

Order of Shells

Principle Quantum Shell → Sub-Shell → Orbitals

Number of orbitals for each sub-shell

1. S - 1 Orbital (2 Electrons)

2. P - 3 Orbitals (6 Electrons)

3. D - 5 Orbitals (10 Electrons)

4. F - 7 Orbitals (14 Electrons)

Ionisation Energy Definition

Ionisation energy is the energy required to remove 1 mole of electron from 1 mole of atoms of an element in gaseous state to form 1 mole of gaseous ion.

Equations for ionisation energies 

1. X_(g) → X_(g)⁺¹ + 1e^-

2. X_(g)⁺¹ → X_(g)⁺² + 1e^-

3. X_(g)⁺² → X_(g)⁺³ + 1e^-

Factors affecting ionisation energy

1. Proton Number (Effective Nuclear Charge)
   As the atomic number increases, the positive nuclear charge increases which causes a greater attractive force between the positively charged nucleus and the negatively charged electrons, requiring more energy.

2. Distance From Nucleus
   As the distance from the nucleus increases, the attractive force decreases.

3. Shielding Effect
   The electrons in the inner shells repel the outer electrons because of like-charges, lowering the ionisation energy. 

4. Spin-Pair Repulsion
   Electrons in the same orbital repel each other, causing ionisation energy to decrease.

Shape of S and P orbitals

• S Orbital: Circular

• P Orbital: X/Y/Z infinity sign shape

What is the most stable electronic configuration?

The ones with the lowest energy.

• Half-filled orbitals

• Full orbitals

Order of Orbitals 

1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s²

Exceptions to the electronic configuration pattern

• Copper:
  [Ar] 4s² 3s⁹ → [Ar] 4s¹ 3d¹⁰

• Chromium:
  [Ar] 4s² 3s⁴ → [Ar] 4s¹ 3d⁵

This is to make the the energy level more stable.

Free Radical Definition

A free radical is a species with one or more unpaired electrons.

Atomic/Ionic Radius Trends 

• Down the group:
  Atomic radius increases down the group because of shielding effect from inner electrons

• Across the group:
  The effective nuclear charge increases
  Shielding effect is constant

Positively charged ions are smaller than their atom 

Negatively charged ions are bigger than their atom 

Patterns in ionisation Energy

1. Ionisation energy increases across a period due to:
   Increased Nuclear Charge
   = Atomic Radius
   = Shielding Effect
   

2. Ionisation energy decreases between periods due to:
   Increased Atomic Radius
   Increased Shielding Effect
   

3. Ionisation energy decreases between Beryllium and Boron due to:
   Increased Shielding Effect
   Slight Increase in Atomic Radius

4. Ionisation energy decreases between Nitrogen and Oxygen due to:
   Spin-Pair Repulsion