Ionic bonding and structure

IB Chemistry SL· Structure 3: Chemical Bonding, 3.1 Ionic bonding & structure· 20 min read

1. Formation of Ionic Bonds★★☆☆☆⏱ 5 min

📘 Definition

Ionic Bond

The strong electrostatic force of attraction between oppositely charged ions, formed by the complete transfer of one or more valence electrons from a metal atom to a non-metal atom

Example:

Electrostatic attraction between $N a^{+}$ and $C l^{-}$ in sodium chloride

Ionic bonding typically occurs when there is a large difference in electronegativity (usually > 1.8 Pauling units) between two elements. Metals lose valence electrons to form positively charged cations, achieving a stable noble gas electron configuration. Non-metals gain these electrons to form negatively charged anions, also achieving a full outer shell.

📐 Worked Example

Explain the formation of an ionic bond between magnesium (Z=12) and oxygen (Z=8)

1
Write the electron configuration of each neutral atom:
2
$Mg: 1s^2 2s^2 2p^6 3s^2 \qquad O: 1s^2 2s^2 2p^4$
3
Magnesium loses its 2 valence electrons to achieve a stable full outer shell:
4
$Mg \rightarrow Mg^{2+} + 2e^-$
5
Oxygen gains the 2 electrons to fill its outer shell:
6
$O + 2e^- \rightarrow O^{2-}$
7
Oppositely charged $M g^{2 +}$ and $O^{2 -}$ are attracted by strong electrostatic forces, forming an ionic bond in neutral magnesium oxide (MgO)

Exam tip:

IB examiners require you to explicitly mention electrostatic attraction between ions, not just transfer of electrons, to get full marks.

2. Dot-and-Cross Diagrams for Ionic Compounds★★☆☆☆⏱ 5 min

Dot-and-cross diagrams are used to visually represent electron transfer and ion charge in ionic compounds. Electrons from the metal are drawn as dots, and electrons from the non-metal as crosses (or vice versa) to track their origin. All ions are drawn in square brackets with the net charge written outside the bracket.

📐 Worked Example

Draw and describe the dot-and-cross diagram for calcium fluoride, $C a F_{2}$

1
Identify the ions formed: Calcium (Group 2) loses 2 electrons to form $C a^{2 +}$ . Each fluorine (Group 17) gains 1 electron to form $F^{-}$ , so 2 $F^{-}$ are needed to balance the 2+ charge of calcium.
2
Calcium transfers one valence electron to each fluorine atom, so all ions achieve a noble gas electron configuration.
3
Draw each ion in separate square brackets: $C a^{2 +}$ has no outer electrons (lost both), and each $F^{-}$ has a full outer shell of 8 electrons, with one electron of Ca origin on each $F^{-}$ .
4
Write the charge of each ion outside the bracket: $[C a]^{2 +} [F]^{-} [F]^{-}$

✓ Quick check

Test your understanding of ion ratios

What is the ratio of sodium ions to oxide ions in sodium oxide, $N a_{2} O$ ?
- 1:1
- 2:1
- 1:2
- 3:1
Reveal answer
2:1 —
Sodium is Group 1, forms $N a^{+}$ . Oxygen is Group 16, forms $O^{2 -}$ . Two $N a^{+}$ are required to balance the 2- charge of the oxide ion, giving a 2:1 ratio.

3. Giant Ionic Lattice Structure★★★☆☆⏱ 5 min

📘 Definition

Giant Ionic Lattice

A regular, continuous three-dimensional arrangement of alternating oppositely charged ions, held together by ionic bonds extending throughout the entire crystal

Example:

Sodium chloride forms a face-centred cubic lattice with coordination number 6 for both ions

Ionic compounds do not form discrete molecules. The electrostatic attraction between ions acts in all directions, so ions pack together to form a continuous giant lattice. The coordination number of an ion is the number of oppositely charged ions directly surrounding it in the lattice.

📐 Worked Example

Explain why the formula of sodium chloride is written as $N a C l$ , not $N a_{2} C l_{2}$

1
In the giant sodium chloride lattice, the ratio of $N a^{+}$ ions to $C l^{-}$ ions across the entire structure is 1:1.
2
Ionic compounds do not have discrete molecules, so we use the empirical formula: the simplest whole-number ratio of ions in the compound.
3
The simplest whole-number ratio 1:1 gives the formula $N a C l$ , which is accepted as the correct formula for sodium chloride.

Exam tip:

Always remember: ionic compounds use empirical formulas, not molecular formulas. This is a common multiple-choice question trap.

4. Physical Properties of Ionic Compounds★★★☆☆⏱ 5 min

All physical properties of ionic compounds are direct consequences of their giant ionic lattice structure and strong ionic bonding. The most frequently tested properties in IB exams are melting/boiling point, electrical conductivity, and solubility.

High melting/boiling points: Strong electrostatic attractions between oppositely charged ions require large amounts of heat energy to break.
Electrical conductivity: Only conducts electricity when molten or dissolved in water, not when solid. Ions are fixed in place in a solid lattice, but become mobile when molten/dissolved.
Solubility: Most ionic compounds are soluble in polar solvents like water (which hydrates individual ions), but insoluble in non-polar solvents.

📐 Worked Example

Explain why magnesium oxide (MgO) has a much higher melting point (2800°C) than sodium chloride (NaCl, 801°C)

1
Identify the ion charges: MgO contains $M g^{2 +}$ and $O^{2 -}$ , while NaCl contains $N a^{+}$ and $C l^{-}$ .
2
The strength of electrostatic attraction between ions is proportional to the product of their charges. The product of charges for MgO is (2 × 2) = 4, compared to (1 × 1) = 1 for NaCl.
3
This means ionic bonds in MgO are much stronger than in NaCl. Much more heat energy is required to break the stronger bonds, leading to a much higher melting point.

5. Common Pitfalls

Wrong move:

Describing ionic bonds as 'shared electrons' instead of transferred

Why:

Confuses ionic bonding with covalent bonding, leading to zero marks

Correct move:

Always state ionic bonds form by complete electron transfer, and are held by electrostatic attraction between oppositely charged ions

Wrong move:

Omitting square brackets and charges on ions in dot-and-cross diagrams

Why:

Examiners mark this as incomplete understanding of ion formation

Correct move:

Always enclose each ion in square brackets, and write the net charge outside the bracket for all ions

Wrong move:

Claiming ionic compounds conduct electricity when solid because they contain charged particles

Why:

Forgets that charge carriers must be free to move to conduct an electric current

Correct move:

State that ions are held in fixed positions in the solid lattice, so solid ionic compounds do not conduct; only conduct when molten or aqueous, when ions are mobile

Wrong move:

Claiming all ionic compounds are soluble in water

Why:

Overgeneralizes the solubility rule, ignoring common insoluble ionic compounds

Correct move:

State that many common ionic compounds are soluble in water, but insoluble examples include silver chloride and barium sulfate

6. Quick Reference Cheatsheet

Key Concept	Summary
Bond definition	Electrostatic attraction between oppositely charged ions
Electron movement	Complete transfer from metal to non-metal
Structure type	Giant 3D ionic lattice, no discrete molecules
Formula type	Empirical formula (simplest ion ratio)
Melting point	High, strong ionic bonds need much energy to break
Conductivity (solid)	Does not conduct, ions are fixed
Conductivity (molten/aqueous)	Conducts, ions are mobile

7. Frequently Asked

Why are no bonds 100% ionic?

All bonds have some covalent character, because the cation can polarize the electron cloud of the anion, leading to some electron sharing. Large differences in electronegativity just give predominantly ionic character.

Why do ionic compounds have empirical formulas?

Ionic compounds form giant continuous lattices, not discrete molecules. The formula only shows the simplest whole-number ratio of ions in the structure, which is the empirical formula.

When this came up on past exams

AI-estimated based on syllabus patterns — cross-check with official past papers for accuracy. Use only as revision-focus signals.

2025 · 1
Ionic melting point comparison
2024 · 2
Dot-and-cross diagram for ionic compound
2023 · 1
Conductivity of ionic compounds

What's Next

Ionic bonding is the first of the four main bonding types you will learn in this unit. Mastery of ionic structure and properties is essential for comparing it to other bonding types, and for answering exam questions on structure and bonding, which make up ~20% of your total IB SL Chemistry marks. The concepts of bond strength and lattice energy also connect directly to energetics topics later in the course. Next, you will learn about covalent bonding, the other primary intramolecular bond type, before moving on to intermolecular forces and metallic bonding.