Covalent bonding is central to understanding molecular structure and chemical reactivity. These notes introduce covalent and dative bonding, show how electron-pair sharing forms molecules, and explain how average bond enthalpy provides a comparative measure of bond strength across different covalent species.

Covalent Bonding

Covalent bonding arises when atoms share pairs of electrons to achieve greater stability. This type of bonding dominates in molecules composed of non-metals and is strongly related to electronegativity and orbital overlap.

Covalent bond is introduced here.

Covalent bonds result in the formation of discrete molecules, each with a specific arrangement of nuclei and shared electrons. Electron sharing can be equal or unequal, depending on electronegativity differences, although the OCR specification emphasises the structure and nature of the electron-pair interaction rather than polarity.

Dot-and-cross diagrams

Dot-and-cross diagrams are pictorial representations that show how atoms share electrons to form covalent bonds. They use different symbols to distinguish the electrons originating from each atom. Although simplified, they help visualise:

• Shared electron pairs (bonding pairs)

• Lone (non-bonding) electron pairs

• The number of bonds formed by each atom

• The octet rule where applicable

These diagrams reinforce that each covalent bond consists of two electrons, contributed by one or both bonding atoms.

Dot-and-cross diagram of hydrogen chloride (HCl), showing a shared pair of electrons between hydrogen and chlorine and three lone pairs on chlorine. This visual reinforces how a single covalent bond and lone pairs are represented in electron-pair diagrams. The diagram uses dots rather than different symbols for each atom, but the underlying idea of shared and non-bonding pairs is identical to OCR’s required dot-and-cross approach. Source

Dative (Coordinate) Covalent Bonding

In some molecules and ions, one atom donates both electrons to the shared pair. This still forms a covalent bond but differs in its origin.

Dative covalent bond is introduced here.

A common example is the formation of ammonium ions from ammonia and hydrogen ions. Dative bonds are treated identically to normal covalent bonds once formed, but their formation highlights the flexibility of electron pair donation.

Singly donated pairs often use arrows in diagrams (from donor to acceptor), although the resulting bond behaves as an ordinary covalent bond in all structural and energetic contexts.

Diagram illustrating a coordinate (dative) covalent bond, with an arrow showing a lone pair from the donor atom forming a bond to the acceptor atom. This highlights that both electrons in the bond come from the same atom, even though the resulting bond is indistinguishable from other covalent bonds. The image may also mention Lewis acids and bases, which is useful context but not required knowledge for OCR A-Level Chemistry. Source

Bond Strength and Average Bond Enthalpy

To compare covalent bonds quantitatively, chemists use average bond enthalpies. These values reflect the energy required to break one mole of a specific bond type in the gaseous state, averaged across multiple compounds because bond strengths vary with molecular environment.

Before this can be explored, the term is introduced.

Average bond enthalpies provide insight into molecular stability and reactivity. Higher values indicate stronger bonds requiring more energy to break. This concept supports predictions of reaction energetics and helps explain trends in molecular structures and chemical behaviour.

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Because actual bond enthalpies differ depending on molecular surroundings—for example, C–H bonds in methane differ subtly from those in ethane—averaged values enable consistent comparison across compounds. OCR requires students to understand the meaning of these values and their use as a measure of covalent bond strength, rather than to calculate enthalpy changes in full thermochemical cycles.

Potential-energy curve for the H–H bond showing how energy changes with internuclear distance. The minimum of the curve marks the equilibrium bond length, while the depth of the well corresponds to the bond energy required to separate the atoms. The diagram may also label regions of net attraction and repulsion, which extend slightly beyond the OCR specification but help explain why stronger bonds are shorter and have higher bond enthalpies. Source

Factors Influencing Covalent Bond Strength

Bond strength is influenced by several conceptual factors that relate back to electronic structure and orbital interaction. These factors help explain why some covalent bonds have higher bond enthalpies than others.

Bond length

Shorter bonds tend to be stronger. As atomic radii decrease, the shared electrons are held more closely to both nuclei, enhancing electrostatic attraction within the bond. This is why bonds involving smaller atoms (e.g., H–H, F–F) often possess high bond enthalpies.

Bond order

Bond order refers to the number of shared electron pairs between atoms:

• Single bonds involve one shared pair

• Double bonds share two pairs

• Triple bonds share three pairs

Increasing bond order generally leads to:

• Greater electron density between nuclei

• Shorter bond length

• Higher bond enthalpy

Thus, a C≡C triple bond is significantly stronger than a C=C double bond.

Electronegativity

Differences in electronegativity can influence bond strength by affecting how tightly electrons are held within the shared region. Polar covalent bonds often exhibit enhanced ionic character, which can strengthen or weaken the bond depending on the size and charge distribution of the atoms involved.

Multiple Covalent Bonds and Their Energetics

In molecules featuring multiple bonds, bond enthalpies cannot simply be multiplied to find total energies because each bond in a multiple bond differs slightly in strength. However, OCR focuses on understanding:

• Single, double, and triple covalent bonds

• How increasing shared electron pairs affects bond strength

• Why bond enthalpies used in thermochemical calculations are averages

Students should be able to interpret values from data tables and relate them to structural features of molecules.

Covalent Structures and Strong Bonds

Covalent bonding underpins both simple molecular substances and giant covalent structures. In giant lattice structures, such as those found in carbon allotropes (diamond, graphite), the network of strong covalent bonds extends throughout the structure, giving rise to:

• High melting and boiling points

• Insolubility in most solvents

• Hardness or, in some cases, electrical conductivity

While this subsubtopic does not require detailed study of lattice structures, recognising that strong covalent bonds give rise to these bulk physical properties reinforces the conceptual link between bond strength and material behaviour.

Summary of Key Points Aligned to OCR Requirements

• A covalent bond involves shared pairs of electrons.

• A dative covalent bond involves both electrons donated by one atom.

• Dot-and-cross diagrams illustrate how electron pairs are shared.

• Average bond enthalpy measures the energy needed to break a covalent bond and reflects bond strength.

• Factors such as bond length, bond order, and electronegativity influence bond strength.