Intermolecular Forces
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You should be comfortable drawing Lewis structures and identifying molecular geometry from the VSEPR model. Familiarity with basic electrostatics and the concept of dipole moments is assumed.
Why this matters
Intermolecular forces dictate the physical state, boiling point, and solubility of substances.
While covalent bonds hold atoms together within a molecule, the forces between molecules determine how those molecules interact in bulk. These interactions, collectively known as intermolecular forces (IMFs), are significantly weaker than intramolecular covalent bonds but are responsible for the macroscopic properties we observe, such as why water is a liquid at room temperature while methane is a gas. Understanding IMFs allows chemists to predict phase behaviour, design solvents, and explain biological recognition processes.
The Nature of Intermolecular Attraction
Imagine two neutral objects that are not touching. If they are perfectly uniform, they do not attract each other. However, if one object has a slight positive charge on one side and a negative charge on the other, or if the electrons in one object are slightly shifted, a weak attraction can arise. In chemistry, all molecules experience these attractions. The strength of the attraction depends on the size of the electron cloud, the shape of the molecule, and the presence of highly electronegative atoms. These forces are not chemical bonds in the traditional sense because they do not involve the sharing or transfer of electrons to form a new chemical entity; rather, they are physical interactions that hold molecules in proximity. The balance between the kinetic energy of the molecules and the potential energy of these attractions determines whether a substance is a solid, liquid, or gas.
London Dispersion Forces
London dispersion forces are the weakest type of intermolecular force but are present in all molecules, whether polar or non-polar. They arise from temporary fluctuations in the electron distribution around a nucleus. At any instant, the electron cloud may be slightly denser on one side of the atom or molecule, creating a temporary dipole. This temporary dipole induces a dipole in a neighbouring molecule, leading to a weak attractive force. The strength of dispersion forces increases with the size of the electron cloud (polarizability). Larger atoms and molecules have more electrons that are further from the nucleus, making them more easily distorted. Consequently, dispersion forces are the dominant IMF for non-polar molecules and become significant even in polar molecules as the molar mass increases.
Dipole–Dipole Interactions
Polar molecules possess a permanent dipole moment due to the unequal sharing of electrons in polar bonds and an asymmetric molecular geometry. The positive end of one dipole is attracted to the negative end of another. These dipole–dipole interactions are stronger than London dispersion forces for molecules of similar molar mass. The strength of this interaction depends on the magnitude of the dipole moment, which is a vector quantity. For a molecule to have a net dipole, the individual bond dipoles must not cancel out. For example, carbon dioxide is linear and non-polar, whereas water is bent and polar. The alignment of dipoles in a liquid or solid maximizes the attractive interactions, contributing to higher boiling points compared to non-polar analogues.
where \(\mu\) is the dipole moment, \(q\) is the partial charge, and \(d\) is the distance between the charges.
Comparative Strengths and Trends
The relative strength of intermolecular forces generally follows the order: London dispersion < dipole–dipole < hydrogen bonding. However, this hierarchy is not absolute; for very large molecules, dispersion forces can exceed dipole–dipole interactions. When comparing boiling points, one must consider both the type of IMF and the molar mass. A non-polar molecule with a high molar mass may have a higher boiling point than a small polar molecule because its dispersion forces are stronger. For instance, bromine (Br2) is a liquid at room temperature due to strong dispersion forces, while hydrogen fluoride (HF), despite its small size, boils at about 20 °C because of hydrogen bonding. The cumulative effect of these forces determines the energy required to separate molecules from the liquid phase into the gas phase.
A common misconception is that hydrogen bonding is a type of covalent bond. It is not. Hydrogen bonding is an electrostatic interaction between a hydrogen atom covalently bonded to a highly electronegative atom (N, O, or F) and a lone pair on another electronegative atom. It is an intermolecular force, not an intramolecular bond, and is significantly weaker than a typical O–H covalent bond. Confusing the two leads to incorrect predictions about bond dissociation energies and reaction mechanisms. C
Worked example
Problem. Compare the expected boiling points of the following three compounds: propane (C3H8, M = 44.1 g/mol), acetone (C3H6O, M = 58.1 g/mol), and 1-propanol (C3H8O, M = 60.1 g/mol). Rank them from lowest to highest boiling point and justify your ranking based on intermolecular forces.
- Identify the molecular structure and polarity of each compound. Propane is non-polar. Acetone is polar due to the C=O bond. 1-Propanol is polar and contains an O–H group capable of hydrogen bonding.
- Determine the dominant intermolecular forces. Propane exhibits only London dispersion forces. Acetone exhibits dipole–dipole interactions and London dispersion forces. 1-Propanol exhibits hydrogen bonding, dipole–dipole interactions, and London dispersion forces.
- Compare the molar masses. The molar masses are similar (44.1, 58.1, and 60.1 g/mol), so differences in dispersion forces are minor compared to the differences in specific IMF types.
- Rank the strengths of the IMFs. Hydrogen bonding is the strongest, followed by dipole–dipole, then dispersion. Therefore, 1-propanol will have the highest boiling point, followed by acetone, and finally propane.
Units check. The ranking is qualitative, but the justification relies on the relative strengths of IMFs, which are consistent with experimental boiling points (propane ~ -42°C, acetone ~ 56°C, 1-propanol ~ 97°C).
Three quick tests
- Which intermolecular force is present in all molecules?
Answer
London dispersion forces. - Why does water have a higher boiling point than hydrogen sulfide (H2S)?
Answer
Water molecules form strong hydrogen bonds with each other, whereas H2S molecules only experience weaker dipole–dipole and dispersion forces. - How does increasing the molar mass of a non-polar molecule affect its boiling point?
Answer
It increases the boiling point because larger molecules have more electrons, leading to stronger London dispersion forces.
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Every lecture and every result in this unit, with prerequisites and threads.
Lecture notes drafted with a local model and checked line by line by the editor; every number in the worked example was recomputed.