Covalent Bonding & Structure (OCR AS Chemistry A): Revision Note

Covalent bonds & strength

• Covalent bonding occurs between two non-metals

• A covalent bond is the electrostatic attraction between the nuclei of bonding atoms and a shared pair of outer-shell electrons

• This attraction allows atoms to achieve a stable noble gas configuration without transferring electrons

• Different types of covalent bonds can form, depending on the number of electron pairs shared

Covalent bonds & shared electrons table

Bond energy

• Bond energy is the energy required to break one mole of covalent bonds in the gaseous state

  • Bond energy has units of kJ mol^-1

• The larger the bond energy, the stronger the covalent bond

• Average bond enthalpy is also used as a measurement of bond strength

  • It is the average energy needed to break a specific bond type

  • It is measured across many different compounds

• Multiple covalent bonds (e.g. double or triple) generally have higher bond enthalpies than single bonds

  • This is due to the greater electron density between the nuclei

• In general, stronger bonds are shorter bonds:

  • Triple bonds are shorter and stronger than double bonds

  • Double bonds are shorter and stronger than single bonds

  • This is due to increased electron density pulling the bonded atoms closer together

Dot & cross diagrams

• Dot and cross diagrams are used to represent covalent bonding

  • They show only the outer shell electrons of each atom

• To differentiate between the two atoms involved, dots for electrons of one atom and crosses for electrons of the other atom are used

• Electrons are shown in pairs on dot-and-cross diagrams

Single covalent bonding 

Chlorine, Cl₂

Hydrogen chloride, HCl

Ammonia, NH₃

Double covalent bonding

Oxygen, O₂

Carbon dioxide, CO₂

Ethene, C₂H₄

Triple covalent bonding

Nitrogen, N₂

Exceptions to the octet rule

• In some instances, the central atom of a covalently bonded molecule can accommodate more or less than 8 electrons in its outer shell

Expanded octet

• Being able to accommodate more than 8 electrons in the outer shell is known as ‘expanding the octet rule’

• Some examples of this occurring can be seen with period 3 elements

Electron deficient

• Accommodating less than 8 electrons in the outer shell means than the central atom is ‘electron deficient’

Dative covalent / coordinate bonding

• In regular covalent bonds, each atom contributes one electron to the shared pair

• In dative covalent bonding, both electrons in the bond come from the same atom

  • This can also be called coordinate bonding

• This occurs when one atom has a lone pair, and the other atom is electron-deficient

  • This means that the other atom has an incomplete outer shell

• A dative covalent bond is shown using an arrow (→) pointing from the atom donating the lone pair to the atom accepting it

Example: Formation of the ammonium ion (NH₄⁺)

• The nitrogen atom in ammonia (NH₃) has a lone pair

• A hydrogen ion (H⁺) has no electrons, so it can accept a lone pair to form a dative bond

• This forms the ammonium ion (NH₄⁺)

Example: Dimer formation in aluminium chloride (Al₂Cl₆)

• At high temperatures, aluminium chloride exists as AlCl₃

  • This compound is electron-deficient

• At lower temperatures, two AlCl₃ molecules join to form Al₂Cl₆

• So, two chlorine atoms each donate a lone pair to an aluminium atom

  • This forms two dative covalent bonds