Ionic Bonding and Ionic Compounds

The Process

1. A metal atom has a few electrons in its outer shell
2. A non-metal atom needs a few more electrons to fill its outer shell
3. The metal atom loses its outer electrons → forms a positive ion (cation)
4. The non-metal atom gains these electrons → forms a negative ion (anion)
5. The opposite charges attract, forming an ionic bond

Why Does This Happen?

Atoms are most stable with a full outer electron shell (like noble gases). Ionic bonding allows both atoms to achieve this stable configuration:

• Metal atoms achieve stability by losing electrons
• Non-metal atoms achieve stability by gaining electrons

Sodium (Na) — Group 1, electron configuration: 2, 8, 1 Chlorine (Cl) — Group 7, electron configuration: 2, 8, 7

1. Sodium loses 1 electron → $\text{Na}^+$ (configuration: 2, 8)
2. Chlorine gains 1 electron → $\text{Cl}^-$ (configuration: 2, 8, 8)

$\text{Na} \rightarrow \text{Na}^+ + e^-$ $\text{Cl} + e^- \rightarrow \text{Cl}^-$

The $\text{Na}^+$ and $\text{Cl}^-$ ions are held together by the electrostatic attraction between opposite charges. This is the ionic bond.

Dot-and-Cross Diagram for NaCl

In a dot-and-cross diagram:

• Show the outer electrons of the metal atom transferring to the non-metal
• Draw square brackets around each ion
• Write the charge on each ion ($+$ or $-$)
• The transferred electrons should be shown as a different symbol (dot vs cross)

Magnesium (Mg) — Group 2, electron configuration: 2, 8, 2 Oxygen (O) — Group 6, electron configuration: 2, 6

1. Magnesium loses 2 electrons → $\text{Mg}^{2+}$ (configuration: 2, 8)
2. Oxygen gains 2 electrons → $\text{O}^{2-}$ (configuration: 2, 8)

$\text{Mg} \rightarrow \text{Mg}^{2+} + 2e^-$ $\text{O} + 2e^- \rightarrow \text{O}^{2-}$

Magnesium loses 2 electrons, but each chlorine can only accept 1. So two chlorine atoms are needed:

$\text{Mg} \rightarrow \text{Mg}^{2+} + 2e^-$ $2\text{Cl} + 2e^- \rightarrow 2\text{Cl}^-$

Formula: $\text{MgCl}_2$ — one $\text{Mg}^{2+}$ ion for every two $\text{Cl}^-$ ions.

The ratio ensures the overall charge is zero (the compound is electrically neutral).

The charge of an ion depends on which group the element is in:

Ionic compounds do not exist as individual molecules. Instead, billions of positive and negative ions are arranged in a regular, repeating 3D pattern called a giant ionic lattice.

In a giant ionic lattice:

• Each positive ion is surrounded by negative ions
• Each negative ion is surrounded by positive ions
• The electrostatic forces act in all directions
• The structure is held together by strong ionic bonds throughout

For NaCl:

• Each Na⁺ is surrounded by 6 Cl⁻ ions
• Each Cl⁻ is surrounded by 6 Na⁺ ions
• This gives a cubic crystal structure

Electrical Conductivity Explained

For a substance to conduct electricity, it needs charged particles that can move.

• Solid ionic compound: Ions are locked in position in the lattice → cannot conduct
• Molten ionic compound: The lattice breaks down, ions are free to move → conducts
• Dissolved in water: Ions separate and are free to move in solution → conducts

• In dot-and-cross diagrams, show the transferred electron clearly and use brackets with charges
• When explaining high melting points, use the phrase "strong electrostatic forces of attraction between oppositely charged ions"
• For conductivity questions, always mention whether ions are free to move or in fixed positions
• To work out formulae, ensure the total positive charge = total negative charge

• Ionic bonding involves the transfer of electrons from metals to non-metals
• Metal atoms form positive ions (cations); non-metal atoms form negative ions (anions)
• Ions form a giant ionic lattice held by strong electrostatic forces
• Properties: high melting points, conduct when molten/dissolved, brittle, often soluble in water
• Formula of ionic compounds: total positive charge = total negative charge