# Ionic Bonding

> Metals give electrons to non-metals. Writing ionic formulas, naming ionic compounds, and why salt is hard and brittle.

From HSC Chemistry on WhizzNotes, free study notes for NSW students: https://whizznotes.com/chemistry/ionic-bonding/. Updated 3 November 2026.
Year 11 and Year 12 boxes go deeper than the Year 10 text around them.

### Story time: A dangerous pair makes dinner salt
Sodium is a soft, silvery metal that bursts into flame on water. Chlorine is a choking green gas that was used as a weapon in the First World War. Both are dangerous.

Put them together and they react with a bright yellow flash. What is left behind is a white powder you sprinkle on your chips: sodium chloride, table salt.

How can two dangerous elements make something you eat every day? The answer is that sodium **gives** an electron to chlorine. Both become ions with full outer shells, and the new substance behaves nothing like either element.

**In short:** An ionic bond forms when a metal gives electrons to a non-metal, making positive and negative ions that attract each other. The ions form a lattice, so ionic compounds have high melting points, are brittle, and conduct electricity only when molten or dissolved.

## What is an ionic bond?

An **ionic bond** forms between a **metal** and a **non-metal**.

1. The metal atom **loses** its valence electrons and becomes a **positive ion**.
2. The non-metal atom **gains** those electrons and becomes a **negative ion**.
3. The opposite charges **attract** strongly. That attraction is the ionic bond.

$$\mathrm{Na} \;(2,8,1) \;\rightarrow\; \mathrm{Na^+} \;(2,8) + e^- \qquad \mathrm{Cl} \;(2,8,7) + e^- \;\rightarrow\; \mathrm{Cl^-} \;(2,8,8)$$

*Diagram: Sodium atom transfers its single outer electron to a chlorine atom, forming a sodium ion and a chloride ion*
*Only the outer electrons are drawn. Sodium's one outer electron moves to the gap in chlorine's outer shell.*

## Charges of common ions

The charge of an ion comes from its group.

| Group | 1 | 2 | 13 | 15 | 16 | 17 |
|---|---|---|---|---|---|---|
| Valence electrons | 1 | 2 | 3 | 5 | 6 | 7 |
| Ion charge | $+1$ | $+2$ | $+3$ | $-3$ | $-2$ | $-1$ |
| Example | $\mathrm{Na^+}$ | $\mathrm{Mg^{2+}}$ | $\mathrm{Al^{3+}}$ | $\mathrm{N^{3-}}$ | $\mathrm{O^{2-}}$ | $\mathrm{Cl^-}$ |

## Writing ionic formulas

An ionic compound has **no overall charge**. The positive and negative charges must cancel.

### Key idea: The swap and drop method
1. Write the symbols with their charges: $\mathrm{Al^{3+}}$ and $\mathrm{O^{2-}}$.
2. **Swap** the numbers of the charges and **drop** them down as subscripts: Al gets 2, O gets 3.
3. Simplify if you can, and leave out any 1s.

Aluminium oxide is $\mathrm{Al_2O_3}$. Check: $2 \times (+3) + 3 \times (-2) = 0$.

### Worked example: More formulas
- Magnesium and chlorine: $\mathrm{Mg^{2+}}$, $\mathrm{Cl^-}$, so $\mathrm{MgCl_2}$.
- Calcium and oxygen: $\mathrm{Ca^{2+}}$, $\mathrm{O^{2-}}$. Swap gives $\mathrm{Ca_2O_2}$, simplify to $\mathrm{CaO}$.
- Potassium and sulfur: $\mathrm{K^+}$, $\mathrm{S^{2-}}$, so $\mathrm{K_2S}$.

## Naming ionic compounds

1. Name the **metal** first, unchanged.
2. Name the **non-metal** second, with the ending changed to **-ide**.

| Non-metal | Ion name |
|---|---|
| Chlorine | chloride |
| Oxygen | oxide |
| Sulfur | sulfide |
| Nitrogen | nitride |
| Fluorine | fluoride |

$\mathrm{MgO}$ is magnesium oxide. $\mathrm{Li_3N}$ is lithium nitride.

## Polyatomic ions

Some ions are groups of atoms bonded together that carry a charge as a whole.

| Ion | Formula | Ion | Formula |
|---|---|---|---|
| Hydroxide | $\mathrm{OH^-}$ | Carbonate | $\mathrm{CO_3^{2-}}$ |
| Nitrate | $\mathrm{NO_3^-}$ | Sulfate | $\mathrm{SO_4^{2-}}$ |
| Ammonium | $\mathrm{NH_4^+}$ | Phosphate | $\mathrm{PO_4^{3-}}$ |

Treat them as one unit. If you need more than one, use **brackets**: calcium hydroxide is $\mathrm{Ca(OH)_2}$, ammonium sulfate is $\mathrm{(NH_4)_2SO_4}$.

### Watch out: Do not write CaOH2
Without brackets, $\mathrm{CaOH_2}$ would mean one O and two H. You need two whole hydroxide ions: $\mathrm{Ca(OH)_2}$.

## Ionic lattices and their properties

Ionic compounds do not form separate molecules. Millions of ions pack into a regular 3D pattern called an **ionic lattice**, each ion surrounded by ions of the opposite charge. That is why salt forms neat cube-shaped crystals.

| Property | Reason |
|---|---|
| **High melting and boiling points** (salt melts at 801 °C) | Strong attraction between ions in every direction takes a lot of energy to break |
| **Hard but brittle** | A hit shifts a layer, like charges line up and repel, and the crystal shatters |
| **Do not conduct as solids** | Ions are locked in place |
| **Conduct when melted or dissolved** | Ions are free to move and carry charge |
| **Many dissolve in water** | Water molecules pull the ions apart |

### Exam tip: Conducting needs moving charges
For anything to conduct electricity, charged particles must be **free to move**. In ionic substances these are ions, so they only conduct when molten or in solution.

**Mind map**

```text
Ionic bonding
  Who
    Metal + non-metal
  How
    Metal loses e⁻: positive ion
    Non-metal gains e⁻: negative ion
    Opposites attract
  Formulas
    Charges cancel to zero
    Swap and drop
    Brackets for polyatomic ions
  Names
    Metal first
    Non-metal ends in -ide
  Lattice properties
    High melting point
    Brittle
    Conducts when molten or dissolved
```

### Year 11: What makes an ionic bond stronger
The strength of the attraction in a lattice depends on:

- **Charge**: bigger charges attract more. $\mathrm{MgO}$ ($\mathrm{Mg^{2+}}$ and $\mathrm{O^{2-}}$) melts at about 2850 °C, much higher than $\mathrm{NaCl}$ ($\mathrm{Na^+}$ and $\mathrm{Cl^-}$) at 801 °C.
- **Size**: smaller ions get closer together and attract more strongly.

This is why magnesium oxide is used to line furnaces.

**Transition metal ions** need Roman numerals in their names because they can have more than one charge: copper(II) chloride is $\mathrm{CuCl_2}$, copper(I) chloride is $\mathrm{CuCl}$.

### Year 12: Lattice energy and solubility
**Lattice energy** is the energy released when one mole of an ionic solid forms from its gaseous ions. It gets larger as the ion charges increase and the ion sizes decrease (Coulomb's law: force is proportional to $\frac{q_1 q_2}{r^2}$).

Whether a salt dissolves depends on a contest: the energy needed to break the lattice against the energy released when water molecules surround the ions (**hydration**). If the lattice is too strong, the salt is insoluble. This is why many salts with $2+$ and $2-$ ions, such as calcium carbonate, barely dissolve, while almost all sodium, potassium, ammonium and nitrate salts dissolve.

## Quick summary

- Ionic bonds form between a **metal and a non-metal**.
- Metals lose electrons (positive ions). Non-metals gain electrons (negative ions). Opposite charges attract.
- Formulas: charges must add to zero. Use swap and drop, then simplify.
- Names: metal first, non-metal ending in **-ide**. Polyatomic ions keep their own names.
- Ionic lattices have high melting points, are brittle, and conduct only when molten or dissolved.

## Quiz

Try these questions from the chapter quiz.

1. An ionic bond usually forms between:
   - A. Two non-metals
   - B. A metal and a non-metal
   - C. Two metals
   - D. Two noble gases
   - Answer: B. A metal and a non-metal. The metal gives electrons and the non-metal takes them.

2. In forming an ionic bond, a metal atom:
   - A. Loses electrons and becomes a positive ion
   - B. Gains electrons and becomes a negative ion
   - C. Shares electrons
   - D. Loses protons
   - Answer: A. Loses electrons and becomes a positive ion. Metals have few valence electrons, so they lose them.

3. What charge does a calcium ion (group 2) have?
   - A. +1
   - B. +2
   - C. −2
   - D. 0
   - Answer: B. +2. Group 2 atoms lose 2 electrons.

The whole chapter quiz is on the page, for signed-in students: https://whizznotes.com/chemistry/ionic-bonding/#quiz