Electronegativity – IGCSE Chemistry Definition and Key Facts

Electronegativity is a Supplement concept in the Cambridge 0620 syllabus. It provides the underlying explanation for why some bonds are covalent and others are ionic, and why some covalent molecules have polar bonds.

The 0620 definition

Electronegativity is a measure of the ability of an atom in a molecule to attract the bonding pair of electrons towards itself.

Fluorine has the highest electronegativity of any element. Francium has one of the lowest.

Trends in electronegativity

Across a period (left to right)

Electronegativity increases across a period. Atoms gain more protons across the period, increasing the nuclear charge. The atomic radius also decreases. Both effects mean the nucleus pulls more strongly on bonding electrons.

Example: In Period 2, lithium has low electronegativity; fluorine has the highest.

Down a group

Electronegativity decreases down a group. Although nuclear charge increases, the number of electron shells also increases. The bonding pair is further from the nucleus and is shielded by inner electrons. The attraction weakens.

Example: In Group VII, fluorine is more electronegative than chlorine, which is more electronegative than bromine.

Element	Group	Period	Relative electronegativity
F	VII	2	Very high (4.0)
Cl	VII	3	High (3.0)
O	VI	2	High (3.5)
N	V	2	Moderate (3.0)
C	IV	2	Moderate (2.5)
Na	I	3	Very low (0.9)

The Pauling scale values above are for reference — the 0620 exam does not require numerical values but does expect the trends.

Electronegativity and bond type

The difference in electronegativity between two atoms determines the type of bond formed.

Electronegativity difference	Bond type	Example
Zero or very small	Pure covalent bond	H–H, Cl–Cl
Small to moderate	Polar covalent bond	H–Cl, H–O
Large	Ionic bond	Na⁺ Cl⁻

In a polar covalent bond, the shared electrons are pulled closer to the more electronegative atom. This gives that atom a partial negative charge (delta minus, δ⁻) and the other atom a partial positive charge (δ⁺).

For example, in HCl the chlorine is more electronegative than hydrogen, so the bonding pair sits closer to chlorine: H^(δ+)–Cl^(δ⁻).

Why electronegativity matters at IGCSE level

1. Explains bond type — you can reason about whether NaCl is ionic (large difference) or H₂ is covalent (no difference)
2. Explains polar molecules — water is polar because oxygen is much more electronegative than hydrogen, creating δ+ on H and δ⁻ on O
3. Links to intermolecular forces — the polarity caused by electronegativity differences influences boiling points and solubility

Worked exam question

Explain why sodium chloride is an ionic compound but hydrogen chloride is a covalent compound. Use the idea of electronegativity in your answer. [3]

Sodium has a very low electronegativity and chlorine has a high electronegativity [1]. The large difference means sodium effectively transfers its electron to chlorine, forming ions [1]. In HCl, the electronegativity difference between hydrogen and chlorine is much smaller, so the electrons are shared rather than transferred [1].

Common exam mistakes

1. Confusing electronegativity with electron affinity or ionisation energy — electronegativity specifically refers to attraction within a bond.
2. Stating the trends backwards — electronegativity increases across a period and decreases down a group.
3. Saying fluorine has high electronegativity “because it wants to gain one electron” — the correct reasoning is about nuclear charge and atomic radius, not anthropomorphising the atom.