Covalent Bonding & Structure

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Edexcel IAL Chemistry · Topic 1.7

Covalent Bonding & Structure

Non-metal atoms get more stable by sharing electrons rather than giving them away — and whether that sharing produces a tiny simple molecule or a giant 3D lattice is what decides almost every physical property the substance has.

Summary — The Whole Chapter at a Glance

A covalent bond is the electrostatic attraction between two positive nuclei and a shared pair of electrons between them — nothing is transferred, only shared.
Bonding electrons aren't fixed in place — they're best pictured as a moving charge cloud between the two nuclei.
If the two bonded atoms have equal electronegativity, the bond is non-polar. If they differ, the bond is polar, with δ+ and δ− partial charges appearing.
Atoms usually aim for a full outer shell (the octet rule) — but this can be broken: expanded octets (Period 3+, e.g. PCl₅, SF₆) or electron-deficient atoms (e.g. BF₃, BeCl₂).
Molecules can share more than one pair of electrons — giving double bonds (O₂, CO₂, C₂H₄) and triple bonds (N₂) — these are shorter and stronger than single bonds.
In a dative (coordinate) covalent bond, both shared electrons come from just one of the two atoms (e.g. NH₄⁺, Al₂Cl₆).
Simple molecular structures (H₂O, CO₂, Cl₂...) have low melting/boiling points because only weak intermolecular forces need to be broken — the strong covalent bonds inside each molecule stay intact.
Giant covalent structures (diamond, graphite, silicon(IV) oxide) have extremely high melting/boiling points because you're breaking a continuous network of strong covalent bonds.
Diamond and graphite are both pure carbon but behave completely differently — diamond is hard and doesn't conduct; graphite is soft, slippery, and does conduct — all because of how the atoms are arranged.
Electronegativity (Pauling scale) predicts whether a bond will be covalent, polar covalent, or ionic based on the electronegativity difference between the two atoms.
VSEPR theory predicts 3D molecular shape and bond angle from the number of bonding pairs and lone pairs around the central atom — lone pairs repel more than bonding pairs.

1. What Actually Is a Covalent Bond?

Think of it like this: two atoms both "want" a full outer shell of electrons (like a noble gas). Instead of one atom stealing electrons from the other (that's ionic bonding), covalent bonding is more like a compromise — both atoms bring one electron each to a shared "electron pot" in the middle, and both atoms count those shared electrons as their own. Everyone's happy, nobody actually loses anything.

Formally: a covalent bond is the electrostatic attraction between the nuclei of two atoms and the shared pair of electrons in their outer shells. It forms because two atomic orbitals overlap to create one molecular orbital — and the electrons in that shared space are more stable being attracted to two positive nuclei instead of just one.

H• + H× ──────► H (•×) H 1s¹ 1s¹ H–H Two hydrogen atoms, each with 1 electron, share that pair. Both atoms now "see" 2 electrons — a full shell (like helium).

charge cloud

Key rule: Covalent bonding happens between two non-metals. If you see two non-metal elements combining, expect covalent bonding, not ionic.

Practice Question

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Also in the full note

2. Bond Polarity — Why Some Bonds Are "Lopsided"
3. The Octet Rule — and When It Breaks
4. Single, Double & Triple Bonds
5. Dative (Coordinate) Covalent Bonding
6. Simple Covalent (Molecular) Structures
7. Giant Covalent (Macromolecular) Structures
8. Electronegativity
9. Polar Bonds vs Polar Molecules — A Crucial Distinction

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