Covalent vs Ionic Bonding
An ionic bond is the electrostatic attraction between oppositely charged ions formed when electrons are transferred (usually metal → non-metal). A covalent bond is a shared pair of electrons between two atoms (usually non-metal + non-metal). The difference in how electrons are handled explains their very different properties.
In reality bonding is a spectrum — a large electronegativity difference gives ionic character, a small one gives covalent.
Side-by-side comparison
The single difference — transfer vs sharing of electrons — drives every property below.
Property | Ionic | Covalent |
|---|---|---|
What happens to electrons | Transferred from one atom to another | Shared between two atoms |
Typical elements | Metal + non-metal | Non-metal + non-metal |
Particles formed | Oppositely charged ions (e.g. Na⁺, Cl⁻) | Molecules (e.g. H₂O) or giant covalent structures |
Melting / boiling point | High — strong lattice of ionic attractions | Usually low (simple molecules); very high for giant covalent (e.g. diamond) |
Electrical conductivity | Conducts when molten or dissolved (ions free to move) | Does not conduct (no free charges)* |
Solubility in water | Often soluble | Often insoluble (many dissolve in non-polar solvents) |
*Exception: graphite is covalent yet conducts, because each carbon has a delocalised electron.
It's a continuum, not two boxes
The bond type depends on the electronegativity difference (Δχ) between the atoms:
electrons shared fairly evenly → covalent (e.g. Cl–Cl).
unevenly shared → polar covalent (e.g. H–Cl) — covalent with partial charges.
electrons effectively transferred → ionic (e.g. Na⁺Cl⁻). Even "ionic" compounds have some covalent character.
Worked example — classify the bonding
State whether the bonding in each is mainly ionic or covalent, and justify it: (a) MgO, (b) Cl₂, (c) HCl.
(a) MgO: magnesium is a metal, oxygen a non-metal → electrons are transferred (Mg → Mg²⁺, O → O²⁻).
The large electronegativity difference confirms ionic bonding, and MgO's very high melting point fits a strong ionic lattice.
(b) Cl₂: two identical non-metal atoms, so Δχ = 0 → electrons shared equally → non-polar covalent.
(c) HCl: two non-metals but different electronegativities → shared but unevenly → polar covalent (δ⁺H–Clδ⁻), not ionic.
(a) ionic, (b) covalent (non-polar), (c) covalent (polar).
Common mistakes
Saying ionic compounds "conduct electricity" without a condition — they only conduct when molten or dissolved, not as a solid (ions are locked in the lattice).
Assuming all covalent substances have low melting points — giant covalent structures like diamond and SiO₂ have very high melting points.
Treating ionic and covalent as completely separate — bonding is a continuum; polar covalent bonds sit in between.
Frequently asked questions
What is the main difference between ionic and covalent bonds?
In an ionic bond, electrons are transferred from one atom to another to form oppositely charged ions that attract each other. In a covalent bond, a pair of electrons is shared between two atoms. Ionic bonding is typical between a metal and a non-metal; covalent bonding between two non-metals.
How can you tell if a bond is ionic or covalent?
Look at the elements and their electronegativity difference. Metal + non-metal with a large difference (roughly above 1.7–2.0) is mainly ionic; two non-metals with a small difference is covalent. A medium difference gives a polar covalent bond, which is covalent with partial charges.
Why do ionic compounds conduct electricity but covalent ones usually don't?
Ionic compounds contain charged ions. When molten or dissolved in water, these ions are free to move and carry a current. Most covalent substances have no free charged particles, so they don't conduct — the main exception is graphite, which has delocalised electrons.
This concept in your exam
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