Factors Affecting Covalent Bond Strength

Chemistry: IB Diploma Higher Level, topic 2.2.3

The two main factors that affect the strength of a covalent bond are the bond length and bond order.

A covalent bond is formed when atoms share electron density.
The binding arises due to the electrostatic attraction between the nuclei of bonding atoms and the electron density in the bond.
The bond strength is how strongly the atoms in a bond are held together.
The strength is measured as the amount of energy required to break a mole of bonds in kJ mol^-1. This quantity is known as the bond enthalpy.

Bond length is defined as the average distance between the two nuclei of atoms bonded in a molecule.
The shorter the bond length the stronger the bond and the more energy required to break it (higher bond enthalpy).
Shorter bonds are between atoms with smaller atomic radii.

The relationship between bond length and bond strength is shown by considering the hydrogen halides in the table above.

Bond order is determined by the number of electron pairs in the bonding orbitals between the two atoms in the bond.
Higher bond orders result in shorter and stronger bonds.
An example of this is shown in the table above which compares how the increasing bond order of carbon-carbon bonds affects the bond length and bond energy.

Covalent Bond Polarity

Covalent bonds can be polar or apolar, depending on the difference in electronegativity of the bonded elements.

On average, in perfect covalent bonding, the electron density is located halfway between the two nuclei, and the covalent bond is considered to be apolar.
Apolar covalent bonds are formed between atoms with the same or very similar electronegativities.

The diagram above shows an example of an apolar covalent bond between two chlorine atoms to make a chlorine molecule.

Ionic bonds are created when electron density is considered to be completely transferred from one atom to another. The electrostatic attraction between cations and anions holds the ionic lattice together.
In between the extremes of apolar covalent and ionic bonding, there are polar covalent bonds.

When the electronegativity difference between two non-metals in a covalent bond is significant, the electron density in the bond is attracted more strongly to the more electronegative element.
A dipole is created due to the uneven distribution of electron density, resulting in a polar covalent bond.
The dipole in a molecule is represented by placing partial charges on the atoms.
The more electronegative element in the covalent bond has a partial negative charge (δ-) and the less electronegative element has a partial positive charge (δ+).

An example of a polar covalent molecule is HCl.
Chlorine has a greater electronegativity than hydrogen therefore the electron density in the bond is attracted more strongly to chlorine.
This creates a dipole with chlorine having a partial negative charge (δ-) and hydrogen a partial positive charge (δ+).

What is the correct unit for bond enthalpy?
Which of these statements is NOT correct?
Which of these electronegativity differences is likely to produce the largest dipole?