Revise Intermolecular forces, electronegativity and bond properties for Cambridge 9701 (3.6)

Electronegativity is a measure of the power of an atom to attract the pair of electrons in a covalent bond. The Pauling scale is a common way to quantify this, with fluorine being the most electronegative element (value of 4.0). When two identical atoms bond (e.g., $Cl_2$), their electronegativities are equal, and the electrons are shared perfectly evenly. This is a non-polar covalent bond.

When two different atoms bond (e.g., $HCl$), there is usually a difference in electronegativity. In $HCl$, chlorine is more electronegative than hydrogen, so it pulls the bonding electrons closer to itself. This creates a polar covalent bond, with a partial negative charge ($\delta-$) on the chlorine atom and a partial positive charge ($\delta+$) on the hydrogen atom. This separation of charge is called a bond dipole.

A molecule containing polar bonds is not necessarily a polar molecule. The overall polarity of a molecule depends on both the polarity of its bonds and its three-dimensional shape. If the bond dipoles are arranged symmetrically, they can cancel each other out, resulting in a non-polar molecule. For example, carbon tetrachloride ($CCl_4$) has four polar $C-Cl$ bonds, but its tetrahedral shape means the dipoles cancel, making the molecule non-polar. The same is true for linear carbon dioxide ($CO_2$).

In contrast, a water molecule ($H_2O$) is polar. The two $O-H$ bonds are polar, and the V-shaped (bent) geometry means the dipoles do not cancel out. This results in an overall dipole moment, with the oxygen end being partially negative and the hydrogen end being partially positive.

There are three main types of intermolecular forces, which are all significantly weaker than covalent or ionic bonds. In order of increasing strength, they are:

These are the weakest IMFs and exist between all atoms and molecules. Electrons are in constant, random motion. At any given instant, the electron density in a molecule might be unevenly distributed, creating a temporary, instantaneous dipole. This fleeting dipole can then induce a dipole in a neighbouring molecule, leading to a weak electrostatic attraction. The strength of van der Waals forces increases with the number of electrons in a molecule, which generally correlates with its relative molecular mass ($M_r$).

These forces occur between molecules that are permanently polar (i.e., they have an overall dipole moment). The partially positive ($\delta+$) end of one molecule is electrostatically attracted to the partially negative ($\delta-$) end of a neighbouring molecule. These forces are stronger than van der Waals forces for molecules of comparable size and mass.

A hydrogen bond is a special, particularly strong type of permanent dipole-dipole interaction. It is not a true bond but is the strongest of the intermolecular forces. For hydrogen bonding to occur, a hydrogen atom must be covalently bonded to a very electronegative atom with at least one lone pair of electrons. In A-Level Chemistry, this is limited to Nitrogen (N), Oxygen (O), or Fluorine (F).

$X^{\delta-} - H^{\delta+} \cdots Y^{\delta-}$

Here, X and Y are N, O, or F. The dotted line represents the hydrogen bond between the partially positive hydrogen and the lone pair on the neighbouring electronegative atom. These bonds explain properties like the high boiling point of water and the structure of DNA.