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This lesson covers the three main types of strong chemical bonding — ionic, covalent, and metallic — as well as intermolecular forces, the shapes of molecules, electronegativity, bond polarity, sigma and pi bonds, and modern forms of carbon. Understanding bonding and structure allows you to explain the physical properties of a wide range of substances and is examined in detail across all A-Level Chemistry specifications.
Key Definition: Ionic bonding is the strong electrostatic attraction between oppositely charged ions in an ionic lattice.
Ionic bonding involves the transfer of electrons from a metal atom to a non-metal atom, forming oppositely charged ions that are held together by strong electrostatic attractions. The resulting structure is a giant ionic lattice.
Key features of ionic compounds:
The lattice energy of an ionic compound is a measure of the strength of the ionic bonding. It depends on the charge and radius of the ions — small, highly charged ions form stronger lattices (e.g. MgO has a much higher melting point than NaCl because Mg²⁺ and O²⁻ have higher charges and smaller radii than Na⁺ and Cl⁻).
A dot-and-cross diagram for ionic bonding shows the electron transfer between atoms. For sodium chloride (NaCl), the diagram would show: a sodium atom with electron configuration 2,8,1 transferring its single outer electron (shown as a cross) to a chlorine atom with configuration 2,8,7. The resulting Na⁺ ion has the configuration 2,8, and the Cl⁻ ion has the configuration 2,8,8. Each ion is drawn in square brackets with its charge written outside. Only the outer shell electrons need to be shown — the sodium's donated electron appears as a cross among the dots on chlorine's outer shell.
For magnesium oxide (MgO), two electrons are transferred from Mg to O. The diagram shows Mg losing two outer electrons and O gaining two electrons, giving Mg²⁺ (2,8) and O²⁻ (2,8).
Exam Tip: In dot-and-cross diagrams for ionic compounds, always show the square brackets around each ion and include the charge. Use dots for electrons from one atom and crosses for electrons from the other to show clearly which electrons have been transferred.
Key Definition: A covalent bond is the strong electrostatic attraction between a shared pair of electrons and the nuclei of the atoms involved in the bond.
Covalent bonding involves the sharing of one or more pairs of electrons between atoms. A single covalent bond consists of one shared pair of electrons; a double bond consists of two shared pairs; a triple bond consists of three shared pairs.
For covalent molecules, dot-and-cross diagrams show the shared and lone pairs. For example, in water (H₂O): the central oxygen atom has two bonding pairs (each shared with a hydrogen atom) and two lone pairs. The diagram uses dots for oxygen's electrons and crosses for hydrogen's electrons. In the shared region between O and H, one dot and one cross overlap.
For nitrogen (N₂): a triple bond is shown with three shared pairs (three dots from one N and three crosses from the other), plus one lone pair on each nitrogen.
For carbon dioxide (CO₂): the central carbon shares two pairs with each oxygen (two double bonds). Each oxygen also has two lone pairs. Carbon has no lone pairs.
A dative covalent bond (also called a coordinate bond) is a covalent bond in which both electrons in the shared pair come from the same atom. This is common in:
Once formed, a dative bond is indistinguishable from an ordinary covalent bond.
Key Definition: A sigma bond is formed by the head-on (end-on) overlap of atomic orbitals along the internuclear axis. A pi bond is formed by the sideways overlap of p orbitals above and below the internuclear axis.
A single bond consists of one σ bond. A double bond consists of one σ bond and one π bond. A triple bond consists of one σ bond and two π bonds.
The σ bond has its electron density concentrated along the axis between the two nuclei and allows free rotation. The π bond has its electron density above and below the plane of the nuclei and prevents free rotation — this is why cis–trans isomerism (E/Z isomerism) occurs in alkenes.
In ethene (C₂H₄), each carbon is sp² hybridised. Three sp² hybrid orbitals form three σ bonds (two C–H and one C–C), and the remaining unhybridised p orbital on each carbon overlaps sideways to form the π bond. The molecule is planar with bond angles of approximately 120°.
In ethyne (C₂H₂), each carbon is sp hybridised. Two sp hybrid orbitals form two σ bonds (one C–H and one C–C), and the two remaining p orbitals on each carbon overlap sideways to form two π bonds. The molecule is linear with bond angles of 180°.
Key Definition: Electronegativity is the ability of an atom to attract the bonding pair of electrons in a covalent bond towards itself. It is measured on the Pauling scale.
| Trend | Direction | Reason |
|---|---|---|
| Across a period | Increases (left to right) | Increasing nuclear charge with similar shielding pulls bonding electrons more strongly |
| Down a group | Decreases | Increasing atomic radius and shielding reduce the attraction for bonding electrons |
Fluorine is the most electronegative element (4.0 on the Pauling scale), followed by oxygen (3.4), nitrogen and chlorine (3.0), and carbon (2.6).
When two atoms with different electronegativities form a covalent bond, the bonding electrons are drawn more towards the more electronegative atom. This creates a polar bond with partial charges: δ+ on the less electronegative atom and δ− on the more electronegative atom.
The greater the difference in electronegativity, the more polar the bond. If the electronegativity difference is very large (typically > 1.7), the bond is considered ionic rather than covalent.
A molecule can contain polar bonds but still be non-polar overall if the bond dipoles cancel out due to molecular symmetry.
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