Covalent Bonds & Strength
- 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
Examiner Tip
The values for bond length and bond enthalpy are often different depending on which source they are taken from
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’
Examiner Tip
Covalent bonding takes place between nonmetal atoms.
Remember: Use the periodic table to decide how many electrons are in the outer shell of a nonmetal atom.