2.4.3 Covalent Bonding & Structure

• Covalent bonding occurs between two non-metals
• A covalent bond involves the electrostatic attraction between nuclei of two atoms and the bonding electrons of their outer shells
• No electrons are transferred but only shared in this type of bonding

The positive nucleus of each atom has an attraction for the bonding electrons shared in the covalent bond

• Non-metals are able to share pairs of electrons to form different types of covalent bonds
• Sharing electrons in the covalent bond allows each of the 2 atoms to achieve an electron configuration similar to a noble gas
  • This makes each atom more stable

Covalent Bonds & Shared Electrons Table

Bond energy

• The bond energy is the energy required to break one mole of a particular covalent bond in the gaseous states
  • Bond energy has units of kJ mol^-1

• The larger the bond energy, the stronger the covalent bond is
• Average bond enthalpy is also used as a measurement of the strength of a covalent bond
  • The average bond enthalpy term is the average amount of energy needed to break a specific type of bond, measured over a wide variety of different molecules

Bond length

• The bond length is the internuclear distance of two covalently bonded atoms
  • It is the distance from the nucleus of one atom to another atom which forms the covalent bond

• The greater the forces of attraction between electrons and nuclei, the more the atoms are pulled closer to each other
• This decreases the bond length of a molecule and increases the strength of the covalent bond
• Triple bonds are the shortest and strongest covalent bonds due to the large electron density between the nuclei of the two atoms
• This increase the forces of attraction between the electrons and nuclei of the atoms
• As a result of this, the atoms are pulled closer together causing a shorter bond length
• The increased forces of attraction also means that the covalent bond is stronger

Triple bonds are the shortest covalent bonds and therefore the strongest ones

Dot & cross diagrams

• Dot and cross diagrams are used to represent covalent bonding
• They show just the outer shell of the atoms involved
• 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 

Hydrogen, H₂

Covalent bonding in hydrogen

Chlorine, Cl₂

Covalent bonding in chlorine

Ammonia, NH₃

Covalent bonding in ammonia

Double covalent bonding

Oxygen, O₂

Covalent bonding in oxygen

Carbon dioxide, CO₂

Covalent bonding in carbon dioxide

Ethene, C₂H₄

Covalent bonding in ethene

Triple covalent bonding

Nitrogen, N₂

Covalent bonding in nitrogen

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

Sulfur dioxide, SO₂ – dot and cross diagram

Phosphorus pentachloride, PCl₅ – dot and cross diagram

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

• In simple covalent bonds, the two atoms involved share electrons
• Some molecules have a lone pair of electrons that can be donated to form a bond with an electron-deficient atom
  • An electron-deficient atom is an atom that has an unfilled outer orbital

• So both electrons are from the same atom
• This type of bonding is called dative covalent bonding or coordinate bonding
• An example with a dative bond is in an ammonium ion
  • The hydrogen ion, H⁺ is electron-deficient and has space for two electrons in its shell
  • The nitrogen atom in ammonia has a lone pair of electrons which it can donate to the hydrogen ion to form a dative covalent bond

 

Ammonia (NH₃) can donate a lone pair to an electron-deficient proton (H⁺) to form a charged ammonium ion (NH₄⁺)

• Aluminium chloride is also formed using dative covalent bonding
• At high temperatures aluminium chloride can exist as a monomer (AlCl₃)
  • The molecule is electron-deficient and needs two electrons to complete the aluminium atom’s outer shell

• At lower temperatures the two molecules of AlCl₃ join together to form a dimer (Al₂Cl₆)
  • The molecules combine because lone pairs of electrons on two of the chlorine atoms form two coordinate bonds with the aluminium atoms

Aluminium chloride is also formed with a dative covalent bond in which two of the chlorine atoms donate their lone pairs to each of the aluminium atoms to form a dimer