Chemical Bonding: Electronegativity, Ionic and Metallic Bonds

This guide covers subtopics 3.1, Electronegativity and bonding, 3.2, Ionic bonding, and 3.3, Metallic bonding, from Topic 3 of Cambridge International AS & A Level Chemistry 9701, 2025–2027 series. All three are AS Level content.

Before studying this

At IGCSE or O Level, you met ionic bonding as electron transfer forming oppositely charged ions, and metallic bonding as a lattice of positive ions in a “sea” of delocalised electrons — including the giant lattice structures and the properties (melting point, conductivity, malleability) that follow from them. Ionic, Covalent and Metallic Bonding covers all of that.

AS Level keeps the same electrostatic models but adds two things: a formal, quantitative basis for predicting whether a given pair of atoms will bond ionically or covalently (electronegativity and the Pauling scale), and a proper explanation of why electronegativity itself varies across the Periodic Table, in terms of nuclear charge, atomic radius and shielding — the same three ideas from Atomic Structure: Particles, Radius and Isotopes. This page assumes you’ve read that one.

Syllabus coverage

CAMBRIDGE INTERNATIONAL AS & A LEVEL CHEMISTRY 9701 — AS Level, Topic 3

3.1 Electronegativity and bonding — defining electronegativity as the power of an atom to attract electrons to itself; explaining the factors influencing electronegativity in terms of nuclear charge, atomic radius and shielding by inner shells and sub-shells; stating and explaining the trends in electronegativity across a period and down a group; using differences in Pauling electronegativity values to predict the formation of ionic and covalent bonds (Pauling values given where necessary; covalent character in some ionic compounds not assessed).

3.2 Ionic bonding — defining ionic bonding as the electrostatic attraction between oppositely charged ions (positively charged cations and negatively charged anions); describing ionic bonding including the examples of sodium chloride, magnesium oxide and calcium fluoride.

3.3 Metallic bonding — defining metallic bonding as the electrostatic attraction between positive metal ions and delocalised electrons.

Electronegativity

Electronegativity is the power of an atom in a covalent bond to attract the shared pair of electrons towards itself. It depends on the same three factors that govern atomic radius: a greater nuclear charge pulls bonding electrons in more strongly; a smaller atomic radius puts the nucleus closer to the bonding pair; and greater shielding from inner shells and sub-shells weakens the nucleus’s pull. These factors move together, which is why electronegativity’s trend mirrors the reverse of the atomic radius trend.

Across a period, electronegativity increases — nuclear charge rises while shielding stays roughly constant, so each atom pulls a bonding pair in more strongly. Down a group, electronegativity decreases — the added shells increase shielding and atomic radius by more than the extra nuclear charge compensates for, weakening the pull on a bonding pair despite the larger nucleus. Fluorine, at the top right of the Periodic Table, is the most electronegative element.

Predicting bond type from electronegativity difference: a large difference in Pauling electronegativity between two bonding atoms means one atom can pull the shared pair almost entirely to itself — effectively forming ions — so the bond is ionic. A small or zero difference means neither atom dominates, and the pair is genuinely shared — the bond is covalent. There is no sharp cutoff value to memorise; you’re expected to compare given Pauling values and reason about which end of the spectrum a bond sits at.

Ionic bonding

Ionic bonding is the electrostatic attraction between oppositely charged ions — positively charged cations and negatively charged anions — formed when one atom transfers one or more electrons to another. The specification names three examples worth being able to describe directly:

• Sodium chloride, NaCl — sodium (2,8,1) transfers one electron to chlorine (2,8,7), forming Na⁺ (2,8) and Cl⁻ (2,8,8).
• Magnesium oxide, MgO — magnesium (2,8,2) transfers two electrons to oxygen (2,6), forming Mg²⁺ (2,8) and O²⁻ (2,8).
• Calcium fluoride, CaF₂ — calcium (2,8,8,2) transfers one electron to each of two fluorine atoms (2,7), forming Ca²⁺ (2,8,8) and two F⁻ (2,8).

In each case, the resulting ions pack into a giant ionic lattice held together by electrostatic attraction extending in all directions — the same lattice structure you met at IGCSE, now understood as a direct consequence of the electronegativity difference between the two elements being large enough to favour complete electron transfer over sharing.

Metallic bonding

Metallic bonding is the electrostatic attraction between positive metal ions and delocalised electrons. Metal atoms release their outer-shell electrons into a shared “sea” that is free to move throughout the entire structure, leaving behind a regular lattice of positive metal ions. The bond is the electrostatic attraction between that lattice of cations and the delocalised electrons surrounding them — not a bond between any two specific atoms, but a bonding force extending through the whole structure.

This is the same electrostatic model you already met at IGCSE, and it is still what explains the properties you learned there: electrical conductivity (the delocalised electrons carry charge freely), malleability (the layers of ions can slide past each other without breaking any specific bond, since the electron sea simply redistributes around them), and generally high melting points (a great many electrostatic attractions must be overcome at once). AS Level adds nothing new to the properties — the value added here is the precise, formal definition in electrostatic terms.

Common mistakes

• Defining electronegativity as “how reactive an atom is.” It specifically describes the power to attract a shared bonding pair of electrons — a property of an atom within a covalent bond, not a general reactivity scale.
• Explaining the electronegativity trend across a period with “shielding increases.” As with atomic radius, shielding is roughly constant across a period — the trend is driven by increasing nuclear charge.
• Treating “ionic” and “covalent” as a strict binary decided by a fixed electronegativity-difference cutoff. The syllabus asks you to use the difference to predict which bond type is favoured, not to apply a memorised threshold number.
• Describing metallic bonding as “bonds between metal atoms.” The attraction is between the lattice of positive ions and the delocalised electrons — individual metal atoms are not bonded to specific neighbours.

Quick revision checklist

• Electronegativity: definition, the three factors (nuclear charge, atomic radius, shielding), and the period/group trends
• Using electronegativity difference to predict ionic vs covalent bonding
• Ionic bonding definition, and the NaCl, MgO and CaF₂ examples with their electron transfers
• Metallic bonding definition: electrostatic attraction between metal cations and delocalised electrons

Related resources

• Ionic, Covalent and Metallic Bonding — the IGCSE/O Level lattice structures and properties this resource builds on
• Atomic Structure: Particles, Radius and Isotopes — nuclear charge, atomic radius and shielding explained
• Chemical Bonding: Shapes and Intermolecular Forces — covalent bonding, molecular shape and intermolecular forces (subtopics 3.4–3.7)
• Cambridge AS & A Level Chemistry hub

Written against Cambridge International AS & A Level Chemistry 9701, 2025–2027 series. Always check the current syllabus for your examination year.