Intermolecular Forces Chemistry Tutorial

Key Concepts

• Intermolecular forces refer to the forces that act between discrete molecules.¹
• Three types of intermolecular force can operate between covalent molecules:

  (i) Dispersion Forces: also known as

  ⚛ London Forces (named after Fritz London who first described these forces theoretically 1930)
  ⚛ Weak Intermolecular Forces
  ⚛ van der Waal's Forces² (namd after the person who contributed to our understanding of non-ideal gas behaviour).

  (ii) Dipole-dipole interactions

  (iii) Hydrogen bonds

• Relative strength of intermolecular forces:

  ⚛ Intermolecular forces (dispersion forces, dipole-dipole interactions and hydrogen bonds) are much weaker than intramolecular forces (covalent bonds, ionic bonds or metallic bonds)

  ⚛ Dispersion forces are the weakest intermolecular force (one hundredth-one thousandth the strength of a covalent bond), hydrogen bonds are the strongest intermolecular force (about one-tenth the strength of a covalent bond).

  ⚛ dispersion forces < dipole-dipole interactions < hydrogen bonds

• The strength of the intermolecular force acting between molecules determines how much energy is required to melt or boil a substance:

  ⚛ Higher melting point ≡ stronger intermolecular forces

  ⚛ Lower melting point ≡ weaker intermolecular forces

• Type of intermolecular force effects solubility in a solvent.

  Like dissolves like:

  ⚛ Polar covalent solute dissolves in polar covalent solvent.

  ⚛ Non-polar covalent solute dissolves in non-polar covalent solvent.

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Types of Intermolecular Forces

Three types of force can operate between covalent molecules:

1. Dispersion Forces: also known as

   ⚛ London Forces (named after Fritz London who first described these forces theoretically 1930)
   ⚛ Weak Intermolecular Forces
   ⚛ van der Waal's Forces² (namd after the person who contributed to our understanding of non-ideal gas behaviour).

2. Dipole-dipole interactions
3. Hydrogen bonds

1. Dispersion Forces (London Forces, Weak Intermolecular Forces, van der Waal's Forces):

• Dispersion forces are very weak forces of attraction that act between all molecules and result from:

  (i) momentary dipoles occurring due to uneven electron distributions in neighbouring molecules as they approach one another

  (ii) the weak residual attraction of the nuclei in one molecule for the electrons in a neighbouring molecule.

• The more electrons that are present in the molecule, the stronger the dispersion forces will be.
• Dispersion forces are the only type of intermolecular force operating between non-polar molecules, for example, dispersion forces operate between:

  ⚛ hydrogen (H₂) molecules in a volume of hydrogen gas
  ⚛ chlorine (Cl₂) molecules in a volume of chlorine gas
  ⚛ carbon dioxide (CO₂) molecules in a volume of carbon dioxide gas
  ⚛ dinitrogen tetroxide (N₂O₄) molecules in a volume of dinitrogen tetroxide gas
  ⚛ methane (CH₄) molecules in a volume of methane gas

2. Dipole-dipole Interactions

• Dipole-dipole interactions are stronger intermolecular forces than Dispersion forces
• Dipole-dipole interactions occur between molecules that have permanent net dipoles (polar molecules), for example, dipole-dipole interactions occur between:

  
  ⚛ SCl₂ molecules in a volume of SCl₂ gas
  ⚛ PCl₃ molecules in a volume of PCl₃ gas
  ⚛ CH₃Cl molecules in a volume of CH₃Cl gas

  
  If the permanent net dipole within the polar molecules results from a covalent bond between a hydrogen atom and either fluorine, oxygen or nitrogen, the resulting intermolecular force is referred to as a hydrogen bond (see below).
• The partial positive charge (δ+) on one molecule is electrostatically attracted to the partial negative charge (δ-) on a neighbouring molecule.
• Covalent bonds within a molecule are shown as solid lines.

  The dipole-dipole interaction between molecules is shown as a dotted line between the molecules, for example:

  				δ-Cl |
  H3Cδ+	-	Clδ-	..........	δ+C	H3

3. Hydrogen Bonds

• Hydrogen bonds occur between molecules that have a permanent net dipole resulting from hydrogen being covalently bonded to either fluorine, oxygen or nitrogen.

  For example, hydrogen bonds operate between:

  ⚛ water (H₂O) molecules in a volume of liquid water

  ⚛ ammonia (NH₃) molecules in a volume of ammonia gas

  ⚛ hydrogen fluoride (HF) molecules in a volume of hydrogen fluoride gas

  ⚛ hydrogen peroxide (H₂O₂) molecules in a volume of hydrogen peroxide

  ⚛ alkanols (alcohols) such as methanol (CH₃OH) molecules in a volume of liquid methanol

  ⚛ alkanoic acids (caboxylic acids) such as acetic acid (ethanoic acid, CH₃COOH) in a volume of liquid acetic acid

  ⚛ organic amines such as methanamine (methyl amine, CH₃NH₂) in a volume of methanamine

• Hydrogen bonds are a stronger intermolecular force than either Dispersion forces or dipole-dipole interactions since the hydrogen nucleus is extremely small and positively charged and fluorine, oxygen and nitrogen being very electronegative so that the electron on the hydrogen atom is strongly attracted to the fluorine, oxygen or nitrogen atom, leaving a highly localised positive charge on the hydrogen atom and highly negative localised charge on the fluorine, oxygen or nitrogen atom. This means the electrostatic attraction between these molecules will be greater than for the polar molecules that do not have hydrogen covalently bonded to either fluorine, oxygen or nitrogen.

Effect of Intermolecular forces on Melting Points and Boiling Points of Molecular Covalent Substances

Since melting or boiling result from a progressive weakening of the attractive forces between the covalent molecules, the stronger the intermolecular force is, the more energy is required to melt the solid or boil the liquid.

If only dispersion forces are present, then the more electrons the molecule has (and consequently the more mass it has) the stronger the dispersion forces will be, so the higher the melting and boiling points will be.

Consider the hydrides of Group 14 elements, all of which are non-polar molecules, so only dispersion forces act between the molecules.

The Group 14 hydrides are:
    CH₄ (molecular mass ≈ 16)
    SiH₄ (molecular mass ≈ 32)
    GeH₄ (molecular mass ≈ 77)
    SnH₄ (molecular mass ≈ 123)
    All of these molecules can all be considered non-polar covalent molecules.

As the mass of the molecules increases, so does the strength of the dispersion force acting between the molecules.

As the strength of the dispersion forces acting between the molecules increases, more energy is required to weaken the attraction between the molecules resulting in higher boiling points.

So, as the mass of the Group 14 hydride molecules increase, the boiling point of the substance increases as shown by the graph below:

If a covalent molecule has a permanent net dipole then the force of attraction between these molecules will be stronger than if only dispersion forces were present between the molecules.
As a consequence, this substance will have a higher melting point or boiling point than similar molecules that are non-polar in nature.

Consider the boiling points of the hydrides of Group 17 (halogen) elements.

The hydrides of Group 17 (halogen) elements are:
    HF (molecular mass ≈ 20)
    HCl (molecular mass ≈ 37)
    HBr (molecular mass ≈ 81)
    HI (molecular mass ≈ 128)

All of these molecules are polar, the hydrogen atom having a partial positive charge (H^δ+) and the halogen atom having a partial negative charge (F^δ-, Cl^δ-, Br^δ-, I^δ-).

As a consequence, the stronger dipole-interactions acting between the hydride molecules of Group 17 elements results in higher boiling points than for the hydrides of Group 14 elements as seen above.

The boiling point of the hydrides of Group 17 elements is shown in the graph below:

With the exception of HF, as the molecular mass increases, the boiling point of the hydrides increase.
HF is an exception because of the stronger force of attraction between HF molecules resulting from hydrogen bonds acting bewteen the HF molecules.
Weaker dipole-dipole interactions act between the molecules of HCl, HBr and HI.
So HF has a higher boiling point than the other molecules in this series.

There is a full discussion of the effect of intermolecular forces on the group 16 hydrides in the Trends in Group 16 Hydrides tutorial.

Effect of Intermolecular Forces on Solubility

In general like dissolves like:

• Non-polar solutes dissolve in non-polar solvents.

  For example, paraffin wax (C₃₀H₆₂) is a non-polar solute that will dissolve in non-polar solvents like oil, hexane (C₆H₁₄) or carbon tetrachloride (CCl₄).
  Paraffin wax will NOT dissolve in polar solvents such as water (H₂O) or ethanol (ethyl alcohol, C₂H₅OH).

• Polar solutes dissolve in polar solvents.³

  For example, polar solutes such as glucose (C₆H₁₂O₆) will dissolve in polar solvents such as water (H₂O) or ethanol (ethyl alcohol, C₂H₅OH) as the partially positively charged atom (δ+) of the solute molecule is attracted to the partially negatively charged atom (δ-) of the solvent molecule, and the partially negatively charged atom (δ-) of the solute molecule is attracted to the partially positively charged atom (δ+) of the solvent molecule.

  Glucose will NOT dissolve in non-polar solvents such as oil, hexane (C₆H₁₄) or carbon tetrachloride (CCl₄).

• Ionic solutes dissolve in polar solvents.³

  For example, ionic solutes such as sodium chloride (NaCl) will generally dissolve in polar solvents but not in non-polar solvents, since the positive ion (+) is attracted the partially negatively charged atom (δ-) in the polar solvent molecule, and the negative ion (-) of the solute is attracted to the partially positively charged atom (δ+) on the solvent molecule.