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Ionic bonding is one of the three main types of chemical bonding you need to understand for A-Level Chemistry. It occurs between metals and non-metals and involves the complete transfer of electrons from one atom to another, producing oppositely charged ions that are held together by strong electrostatic attraction.
Atoms are electrically neutral — they contain equal numbers of protons and electrons. When a metal atom reacts with a non-metal atom, the metal loses one or more electrons from its outer shell, forming a positive ion (cation). The non-metal gains those electrons, forming a negative ion (anion). Both ions achieve a stable noble gas electron configuration.
For example, sodium (2,8,1) loses one electron to become Na⁺ (2,8), and chlorine (2,8,7) gains that electron to become Cl⁻ (2,8,8). Each ion now has a full outer shell.
The driving force behind this electron transfer is the difference in ionisation energy and electron affinity between the two atoms. Metals have low ionisation energies (they lose electrons easily), while non-metals have high electron affinities (they gain electrons readily).
Calcium (1s² 2s² 2p⁶ 3s² 3p⁶ 4s²) has 2 electrons in its outermost shell. It loses both to form Ca²⁺ (achieving the argon configuration). Oxygen (1s² 2s² 2p⁴) needs 2 electrons to fill its outer shell, forming O²⁻ (achieving the neon configuration). The compound formed is CaO, with ions in a 1:1 ratio because the charges balance.
Aluminium (2,8,3) loses 3 electrons to form Al³⁺. Fluorine (2,7) gains 1 electron to form F⁻. Three fluoride ions are needed per aluminium ion to balance charges, giving the formula AlF₃.
Exam tip: Always check that the total positive charge equals the total negative charge in the formula. If Al is 3+ and F is 1−, you need three F⁻ ions: Al³⁺(F⁻)₃ → AlF₃.
The ionic bond itself is the electrostatic attraction between oppositely charged ions. This is a crucial definition to learn. Note that an ionic bond is not between two specific atoms — it acts in all directions. Each positive ion attracts every surrounding negative ion, and vice versa.
This non-directional nature of ionic bonding means that ionic compounds do not exist as discrete molecules. Instead, they form giant ionic lattices — extended three-dimensional structures containing billions of ions arranged in a regular, repeating pattern.
The sodium chloride lattice is the most common example you will encounter. In this structure:
The ions are packed in a face-centred cubic arrangement. If you imagine a cube, Na⁺ and Cl⁻ ions alternate at every corner and along every edge. This maximises the attractive forces between oppositely charged ions while minimising repulsion between like charges.
Magnesium oxide has the same crystal structure as NaCl but with some important differences. Magnesium loses two electrons to form Mg²⁺, and oxygen gains two electrons to form O²⁻. The higher charges on the ions mean that the electrostatic attraction is much stronger.
This is reflected in the melting points: NaCl melts at 801°C, while MgO melts at 2852°C. Both have the same structure, but the doubled charges in MgO produce dramatically stronger ionic bonds.
Lattice energy is the enthalpy change when one mole of an ionic compound is formed from its gaseous ions under standard conditions. It is always exothermic (negative) because forming a lattice releases energy as the ions come together from an infinite separation.
For NaCl: Na⁺(g) + Cl⁻(g) → NaCl(s) ΔH_lattice = −787 kJ mol⁻¹
The more negative (more exothermic) the lattice energy, the more stable the ionic lattice.
Two key factors determine the magnitude of lattice energy:
1. Ionic charge: Higher charges produce stronger electrostatic attraction. The force between ions is proportional to the product of the charges (q⁺ × q⁻). Doubling one charge doubles the attraction; doubling both quadruples it. This is why MgO (2+ and 2−) has a much more exothermic lattice energy than NaCl (1+ and 1−).
2. Ionic radius: Smaller ions allow the charges to get closer together, increasing the electrostatic attraction. Lattice energy is inversely proportional to the sum of the ionic radii. LiF, where both ions are very small, has a more exothermic lattice energy than KBr, where both ions are larger.
These two factors together give the concept of charge density — the ratio of charge to size. High charge density (small, highly charged ions) leads to stronger lattice energies.
| Compound | Ions | Lattice energy / kJ mol⁻¹ | Melting point / °C |
|---|---|---|---|
| LiF | Li⁺, F⁻ | −1037 | 845 |
| NaCl | Na⁺, Cl⁻ | −787 | 801 |
| KBr | K⁺, Br⁻ | −682 | 734 |
| MgO | Mg²⁺, O²⁻ | −3850 | 2852 |
| CaO | Ca²⁺, O²⁻ | −3401 | 2614 |
| Al₂O₃ | Al³⁺, O²⁻ | −15916 | 2072 |
Notice how the lattice energy broadly correlates with melting point: more exothermic lattice energy → higher melting point.
The strong electrostatic forces in ionic lattices explain several characteristic properties:
High melting and boiling points: A large amount of energy is required to overcome the strong attractions between ions throughout the lattice. The melting points generally increase with lattice energy.
Brittleness: When a force is applied to an ionic lattice, layers of ions may shift. If like-charged ions are brought next to each other, the resulting repulsion causes the crystal to shatter.
Electrical conductivity: Solid ionic compounds do not conduct electricity because the ions are held in fixed positions and cannot move. When melted or dissolved in water, the ions become free to move and can carry charge, so molten and aqueous ionic compounds are good electrical conductors.
Solubility in water: Many ionic compounds dissolve in water because the polar water molecules can surround and stabilise the individual ions (hydration). The ion-dipole attractions between the water molecules and the ions must be strong enough to compensate for the lattice energy that must be overcome.
| Property | Ionic | Simple covalent | Giant covalent | Metallic |
|---|---|---|---|---|
| Melting point | High | Low | Very high | Variable (high) |
| Hardness | Hard, brittle | Soft | Very hard | Malleable |
| Solid conductivity | No | No | No (except graphite) | Yes |
| Molten conductivity | Yes | No | N/A | Yes |
| Solubility in water | Often soluble | Depends on polarity | Insoluble | Insoluble |
The tendency to form ionic bonds increases with greater difference in electronegativity between the two elements. A rough guideline is that an electronegativity difference greater than about 1.7 typically results in an ionic bond, though this boundary is not sharp — bonding exists on a continuum from purely covalent to predominantly ionic.
Elements from Group 1 and Group 2 bonded with elements from Group 6 and Group 7 typically form ionic compounds. The further apart the elements are on the periodic table (in terms of electronegativity), the more ionic the bond.
Predict the order of melting points: NaF, NaCl, NaBr, NaI.
All four compounds have the NaCl-type structure with 1+ and 1− ions. The cation (Na⁺) is the same in all cases. The anion changes: F⁻ < Cl⁻ < Br⁻ < I⁻ in size. Smaller anions allow closer approach, giving stronger electrostatic attraction and more exothermic lattice energies.
Predicted order: NaI < NaBr < NaCl < NaF (increasing melting point).
Actual values confirm this: NaI (661°C) < NaBr (747°C) < NaCl (801°C) < NaF (993°C).
Common exam mistake: Students sometimes say "NaF has stronger ionic bonds because fluorine is more electronegative." While fluorine is indeed more electronegative, the correct reasoning is about ionic radius — F⁻ is the smallest halide ion, allowing the closest approach and strongest electrostatic attraction. Electronegativity explains why the bond is ionic, but ionic radius explains the relative strength.