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Notes/Chemistry/Paper 1/Chemical Bonding
CAIEAS Level9701§3.1–3.7

Chemical Bonding

Electronegativity; ionic, metallic, covalent and coordinate bonding; dot-and-cross diagrams; VSEPR shapes and bond angles; σ and π bonds and hybridisation; bond polarity; and the forces between molecules that explain melting and boiling points.

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In Atomic Structure you learned how electrons are arranged in shells and sub-shells, and in Atoms, Molecules and Stoichiometry you wrote formulae for ionic compounds. This note explains what holds those atoms and ions together, and how strongly.

You start with electronegativity, which predicts the type of bond two atoms form. Then come ionic, metallic, covalent and coordinate bonds, and how to draw them as dot-and-cross diagrams. Next you predict the shapes of molecules, look at the two ways orbitals overlap to make a bond (σ\sigma and π\pi bonds), and decide whether a molecule is polar. Finally you use the forces between molecules to explain boiling points and the unusual properties of water.

Before you start you should be able to

  • Simple ionic bonding by electron transfer (e.g. Na⁺ and Cl⁻ forming NaCl) and simple covalent bonding by electron sharing (e.g. H₂, Cl₂, CH₄, HCl) (IGCSE 0620/0971 or O Level 5070/2059)

  • Basic dot-and-cross diagrams for simple molecules, showing only outer-shell electrons

  • That metals conduct electricity because of free-moving electrons

  • Electron configurations, s/p/d sub-shells and orbitals, and the terms nuclear charge and shielding (AS Atomic Structure)

  • Writing the formula of an ionic compound from the charges on its ions (AS Atoms, Molecules and Stoichiometry)

  • Reading group and period position from the Periodic Table

By the end of this page you can

  • Define electronegativity, explain the factors that affect it, and use Pauling values to predict whether a bond will be ionic or covalent

  • Define ionic and metallic bonding, describe ionic bonding in NaCl, MgO and CaF₂, draw a labelled diagram of metallic bonding, and explain conductivity and melting-point differences

  • Define covalent and coordinate (dative covalent) bonding, describe bonding in every named molecule, NH₄⁺ and Al₂Cl₆, and describe octet expansion in Period 3

  • Define bond energy and bond length, explain why bonds get weaker down Group 17, and use bond energies to compare the reactivity of covalent molecules

  • Draw dot-and-cross diagrams for ionic, covalent and coordinate bonding, including expanded-octet and odd-electron species

  • Use VSEPR theory to predict and explain the shape and bond angles of a molecule or ion, including ones not named in the syllabus

  • Describe σ and π bonds as orbital overlap, count them in a molecule, and use sp, sp² and sp³ hybridisation to describe bonding

  • Use electronegativity differences to explain bond polarity, and predict whether a molecule has an overall dipole from its shape

  • Describe hydrogen bonding and the two types of van der Waals' forces, draw a hydrogen bond, and use them to explain boiling-point trends and water's anomalous properties

01

Electronegativity

Syllabus requirement · §3.1

“

define electronegativity as the power of an atom to attract electrons to itself … explain the factors influencing the electronegativities of the elements in terms of nuclear charge, atomic radius and shielding by inner shells and sub-shells … state and explain the trends in electronegativity across a period and down a group … use the differences in Pauling electronegativity values to predict the formation of ionic and covalent bonds.

”

What electronegativity actually measures

Electronegativity is the power of an atom to attract electrons to itself. Learn that wording exactly: it is what the mark scheme asks for. In practice it describes an atom in a covalent bond, pulling on the pair of electrons it shares with another atom. The Pauling scale gives each element a number, from about 0.80.8 (potassium) up to 4.04.0 (fluorine, the most electronegative element). The values you need are given in the question when they matter.

The same three factors as ionisation energy

Electronegativity is decided by the same three factors that decided ionisation energy and atomic radius in the AS Atomic Structure note. It asks the same question: how hard does this nucleus pull on an electron? The only difference is that the electron is now shared.

  • Nuclear charge. More protons pull harder on the shared pair.
  • Atomic radius. A shorter distance between nucleus and shared pair means a stronger pull.
  • Shielding. More inner shells and sub-shells between the nucleus and the shared pair weaken the pull.
The two trends
  • Across a period, nuclear charge rises, the atoms get smaller, and shielding hardly changes (the same inner shells). So electronegativity increases.
  • Down a group, the atoms get larger and have more inner shells, so there is more shielding. This outweighs the rising nuclear charge, so electronegativity decreases.

Fluorine is at the top of Group 17 and near the end of Period 2, so it is the most electronegative element. (Noble gases are left out: they rarely form bonds.)

across Period 3: risesNa0.9Mg1.2Al1.5Si1.8P2.1S2.5Cl3.0more protons, same inner shells,smaller atom → stronger pulldown Group 17: fallsF4.0Cl3.0Br2.8I2.5more shells and shielding,larger atom → weaker pull

Pauling electronegativity rises across Period 3 (more protons, same inner shells, smaller atom) and falls down Group 17 (extra shells and shielding outweigh the rising nuclear charge) — the same factors that decided every trend in Atomic Structure.

From an electronegativity difference to a bond type

The difference in electronegativity between two atoms tells you what kind of bond they form.

  • A large difference: one atom pulls the electrons completely away from the other. Ions form, and the bonding is ionic.
  • No difference, or a small one: the electrons are shared equally, or nearly equally. The bonding is covalent.
  • A difference in between: the electrons are shared, but unequally. This is a polar covalent bond, where each atom carries a small charge (see “Bond polarity and dipole moments”).

You do not need to learn a cut-off value. Questions give you the Pauling values; work out the difference and say whether it is large or small.

Your turn

  1. 19701/22 F/M 2023 Q1(a)(i)1 mark

    Define electronegativity.

    Show solution
    1. 1

      Electronegativity is the power of an atom to attract electrons to itself.

      This is the mark-scheme wording. Answers that describe an ion, or say 'gains electrons', describe electron affinity or ion formation instead and do not score.

    Answer

    The power of an atom to attract electrons to itself.

  2. 2

    State and explain the trend in electronegativity across Period 3, from sodium to chlorine.

    Show solution
    1. 1

      State the trend: electronegativity increases from sodium to chlorine.

      A 'state and explain' question gives a mark for the trend on its own, so write it first and plainly.

    2. 2

      Nuclear charge increases across the period, because each element has one more proton than the one before.

      This is the cause of the trend, so it must be named.

    3. 3

      Every atom from Na to Cl has the same inner shells, so shielding stays about the same, and the atoms get smaller.

      Students often say 'shielding increases' across a period. It doesn't: the extra electrons go into the same outer shell.

    4. 4

      So the nucleus attracts a shared pair of electrons more strongly, and electronegativity increases.

      Always finish by linking the cause back to the attraction for the bonding electrons.

    Answer

    Electronegativity increases across the period, because nuclear charge increases, shielding stays about the same and the atoms get smaller, so the bonding pair is attracted more strongly.

  3. 3

    The Pauling electronegativity of sodium is 0.90.9 and of chlorine is 3.03.0. The Pauling electronegativity of carbon is 2.52.5 and of hydrogen is 2.12.1. Predict, with a reason, the type of bonding in sodium chloride and in methane.

    Stuck? Show hint

    Compare the size of the electronegativity difference in each case, not just whether one exists.

    Show solution
    1. 1

      Sodium chloride: difference =3.0−0.9=2.1= 3.0 - 0.9 = 2.1.

      Always subtract the smaller value from the larger one, and write the difference down.

    2. 2

      2.12.1 is a large difference, so chlorine takes sodium's outer electron completely: ionic bonding.

      A large difference means the electrons are transferred, not shared.

    3. 3

      Methane: difference =2.5−2.1=0.4= 2.5 - 2.1 = 0.4.

      Same method for the second compound.

    4. 4

      0.40.4 is a small difference, so the electrons are shared almost equally: covalent bonding.

      A small difference means the electrons are shared, with only slight polarity.

    Answer

    NaCl: ionic (large electronegativity difference, 2.1). CH₄: covalent (small difference, 0.4).

Paper 1 · MCQReal past-paper questions · Electronegativity and trendsPaper 2 · TheoryReal past-paper questions · Electronegativity and trends

The rest of this note

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Can you do all of these?

  • Define electronegativity and explain its trends using nuclear charge, radius and shielding

  • Predict bond type (ionic vs covalent) from a Pauling electronegativity difference

  • Define ionic bonding and describe it for NaCl, MgO and CaF2\text{CaF}_2

  • Define metallic bonding, draw a labelled diagram of it, and explain conductivity in metals and in molten ionic compounds

  • Explain differences in metallic melting points using ion charge, ion size and number of delocalised electrons

  • Define covalent and coordinate bonding, describe NH₄⁺ and Al₂Cl₆, and explain octet expansion in Period 3

  • Define bond energy and bond length, explain why halogen bonds weaken down the group, and use bond energies to compare reactivity

  • Draw dot-and-cross diagrams for ionic compounds, covalent molecules, coordinate bonds, expanded octets and odd-electron species

  • Predict a shape and bond angle by counting bonding pairs and lone pairs on the central atom, including for ions and unfamiliar molecules

  • Explain how lone pairs make bond angles smaller than the base-arrangement value

  • Describe σ and π bonds as head-on and sideways overlap, count them in a molecule, and assign sp/sp²/sp³ hybridisation from the number of electron regions

  • Decide whether a molecule has an overall dipole from its bond polarities and its shape's symmetry

  • Distinguish id–id, pd–pd and hydrogen bonding, and use them (with electron count and branching) to rank boiling points

  • Draw a hydrogen bond with δ+/δ− charges and a lone pair in line, and explain water's anomalous properties

Paper 1 · MCQ
Real past-paper questions, sorted by difficulty, with mark schemes
Paper 2 · Theory
Real past-paper questions, sorted by difficulty, with mark schemes