# Ionic bonding and structure

> IB Chemistry HL · Unit 2: Models of bonding and structure
> Source: https://www.owlsprep.com/study/ib-chemistry-hl-u2-ionic-bonding-and-structure/

This sub-topic covers ionic bond formation, ion charge prediction, giant ionic lattice structure, and how structure relates to the physical properties of ionic compounds. It is a core foundational bonding model for IB Chemistry HL.

**Prerequisites:** [Atomic structure and electron configuration](https://www.owlsprep.com/study/ib-chemistry-hl-u1-atomic-structure/); [Periodic trends in electronegativity](https://www.owlsprep.com/study/ib-chemistry-hl-u1-periodicity/)

## Learning objectives

- Describe the formation of ionic bonds between oppositely charged ions
- Predict ion charges from periodic table group positions
- Deduce the formula of ionic compounds from ion charges
- Explain how giant ionic lattice structure relates to physical properties

## Formation of Ionic Bonds

**Ionic bond** — Electrostatic attraction between oppositely charged ions formed by electron transfer between elements with large electronegativity differences (usually >1.7 on the Pauling scale)

*Example:* Bond between $\text{Na}^+$ and $\text{Cl}^-$ in sodium chloride

Ionic bonding typically occurs between a metal (low ionization energy, loses electrons easily) and a non-metal (high electron affinity, gains electrons easily). The metal forms a positive cation, the non-metal forms a negative anion, and electrostatic attraction holds the ions together.

**Worked example:** Show the formation of magnesium oxide from magnesium and oxygen atoms via electron transfer.

1. Write the ground-state electron configuration of each neutral atom:
2. $$Mg: 1s^2 2s^2 2p^6 3s^2 \\ O: 1s^2 2s^2 2p^4$$
3. Identify electrons lost/gained to form stable full outer shells:
4. $$Mg \rightarrow Mg^{2+} + 2e^- \\ O + 2e^- \rightarrow O^{2-}$$
5. Oppositely charged ions are attracted electrostatically to form neutral magnesium oxide, with a full ionic bond.

> **Exam tip:** Always explicitly state the attraction is electrostatic between oppositely charged ions for full marks; 'opposite charges attract' is not sufficient.

*Calculator:* forbidden

## Formula of Ionic Compounds

All neutral ionic compounds have zero net charge, so total positive charge from cations equals total negative charge from anions. We can use this rule to deduce the simplest whole-number (empirical) formula.

**Empirical formula (ionic)** — The simplest whole-number ratio of cations to anions in the giant ionic lattice

**Worked example:** Deduce the formula of aluminum sulfate, where aluminum forms $\text{Al}^{3+}$ and sulfate is the polyatomic ion $\text{SO}_4^{2-}$.

1. Find the lowest common multiple of the absolute charge values: $|+3| \times |-2| = 6$
2. Calculate the number of each ion needed for zero net charge: 2 $\text{Al}^{3+}$ give $2 \times (+3) = +6$, 3 $\text{SO}_4^{2-}$ give $3 \times (-2) = -6$. Total charge = 0.
3. Write the formula, using parentheses around polyatomic ions if there is more than one: $\text{Al}_2(\text{SO}_4)_3$

**Check your understanding**

Test your understanding of ion formulas

1. What is the formula of calcium phosphate, where calcium is $\text{Ca}^{2+}$ and phosphate is $\text{PO}_4^{3-}$?

   - $\text{CaPO}_4$
   - $\text{Ca}_3(\text{PO}_4)_2$
   - $\text{Ca}_2\text{PO}_4$
   - $\text{Ca}(\text{PO}_4)_2$

   *Answer:* $\text{Ca}_3(\text{PO}_4)_2$

   *Why:* 3 calcium ions give $3 \times (+2) = +6$, 2 phosphate ions give $2 \times (-3) = -6$, for a net charge of zero.

*Calculator:* forbidden

## Giant Ionic Lattice Structure

Ionic compounds do not form discrete small molecules. Instead, ions arrange into a regular, repeating 3D structure called a giant ionic lattice, where every ion is surrounded by oppositely charged ions. This arrangement maximizes attractive electrostatic forces and minimizes repulsion.

> **Coordination Number**
>
> The coordination number of an ion is the number of adjacent oppositely charged ions. In sodium chloride, both $\text{Na}^+$ and $\text{Cl}^-$ have a coordination number of 6.

**Worked example:** Explain why a cation has a smaller ionic radius than its parent neutral atom, while an anion has a larger radius than its parent.

1. For a cation: The metal loses all valence electrons to form the ion, so an entire electron shell is lost. The same nuclear charge acts on fewer electrons, pulling the remaining electron cloud closer to the nucleus, decreasing radius.
2. For an anion: The non-metal gains extra electrons to form the ion, so nuclear charge stays the same while electron count increases. Increased electron-electron repulsion expands the electron cloud, increasing radius.

*Calculator:* forbidden

## Physical Properties of Ionic Compounds

- **High melting/boiling points**: Strong electrostatic attraction between ions requires large amounts of energy to break the lattice.
- **Electrical conductivity**: Only conduct when molten or dissolved in water, not when solid. Solid lattices have fixed ions (no mobile charge carriers).
- **Brittleness**: Applying force shifts ion layers, bringing like charges next to each other; repulsion splits the crystal.
- **Solubility**: Most ionic compounds dissolve in polar solvents like water, which hydrate individual ions and break the lattice.

**Worked example:** Explain why solid sodium chloride does not conduct electricity, but molten sodium chloride does.

1. In solid sodium chloride, $\text{Na}^+$ and $\text{Cl}^-$ ions are held in fixed positions in the giant ionic lattice by strong electrostatic attraction. There are no free mobile charge carriers to carry an electric current.
2. When sodium chloride melts, the ionic bonds are broken and ions become free to move throughout the liquid. Mobile charged ions can carry an electric current, so molten NaCl conducts.

> **Exam tip:** Always link the observed physical property directly to ionic lattice structure to earn full marks in IB exams.

*Calculator:* forbidden

## Common pitfalls

- **Wrong:** Stating ionic bonding is just "attraction between ions" without mentioning electrostatic attraction
  - Why it fails: IB exams require explicit reference to the electrostatic nature of the bond for full marks
  - Correct: Always describe ionic bonding as the electrostatic attraction between oppositely charged ions
- **Wrong:** Writing calcium hydroxide as $\text{CaOH}_2$ instead of $\text{Ca(OH)}_2$
  - Why it fails: Omitting parentheses around polyatomic ions leads to ambiguity about the number of ions
  - Correct: Use parentheses around polyatomic ions when more than one is present in the formula
- **Wrong:** Claiming ionic compounds conduct electricity via free electrons
  - Why it fails: This confuses ionic bonding with metallic bonding; ionic compounds have no delocalized free electrons
  - Correct: State conductivity comes from mobile ions when molten or dissolved in water
- **Wrong:** Assuming all ionic compounds are soluble in water
  - Why it fails: Many common ionic compounds (e.g. silver chloride, barium sulfate) are insoluble
  - Correct: State most ionic compounds are soluble in polar solvents, but solubility varies based on lattice and hydration energy
- **Wrong:** Thinking ionic compounds only contain two elements
  - Why it fails: Polyatomic ions (e.g. nitrate, ammonium) form ionic compounds with 3+ elements
  - Correct: Recognize ionic compounds can contain multiple elements when polyatomic ions are present

## Cheatsheet

| Concept | Key Exam Fact |
| --- | --- |
| Ionic bond | Electrostatic attraction between oppositely charged ions |
| General formation | Metal (loses e⁻ → cation) + Non-metal (gains e⁻ → anion) |
| Common ion charges | Group 1: +1, Group 2: +2, Group 15: -3, Group 16: -2, Group 17: -1 |
| Structure | Giant ionic lattice, no discrete molecules |
| Melting point | High: lots of energy to break strong ionic attraction |
| Solid conductivity | No: no mobile charge carriers |
| Molten/aqueous conductivity | Yes: mobile charged ions carry current |

## What's next

Ionic bonding is the first core bonding model you will learn in IB Chemistry HL, and it forms the foundation for all subsequent topics in this unit and the rest of the syllabus. You will build on this knowledge to compare ionic properties with covalent and metallic bonding, and extend it to HL topics like lattice enthalpy and Born-Haber cycles. Mastery of ionic compound formulas is also essential for stoichiometric calculations, writing net ionic equations, and balancing redox reactions later in the course.

- [Covalent bonding and Lewis structures](https://www.owlsprep.com/study/ib-chemistry-hl-u2-covalent-bonding-and-lewis-structures/)
- [VSEPR theory and basic hybridization](https://www.owlsprep.com/study/ib-chemistry-hl-u2-vsepr-theory-and-basic-hybridization/)
- [Intermolecular forces](https://www.owlsprep.com/study/ib-chemistry-hl-u2-intermolecular-forces/)

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