3.5 Intermolecular Forces

Learning Objectives

  • Understand the types of intermolecular forces: dipole-dipole interactions, hydrogen bonding and London dispersion forces, and their relative strengths.
  • Learn how to identify what attraction forces are likely to present within a given molecule depending on its structure.

The bonds we have just discussed (covalent, ionic and metallic) are all examples of intramolecular forces: the forces inside a molecule that keep it together. However, molecules commonly interact with each other. For instance, the formation of solids is due to the compounds being within close proximity to each other. What is holding these molecules together?

Dipole-Dipole Interactions and Hydrogen Bonding

Like the ends of a magnet, opposite charges attract, while similar charges repel. As a result, attraction forces between molecules are naturally present in those with charge. Recall that hydrogen chloride, a polar covalent molecule, experiences partial charges on its constituents. In this case, chlorine is more electronegative and is partially negative, given the closeness of the electron to the atom. On the other hand, hydrogen is partially positive because the electron is further away (see Figure 3.5.1).

A lewis dot diagram of hydrogen chloride with single covalent bond between the elements hydrogen and chlorine. Hydrogen has a partial positive charge; and chlorine a partial negative charge along with three pairs of valence electrons.
Figure 3.5.1: Hydrogen chloride (or more commonly hydrochloric acid) with partial charges.

When a molecule of hydrogen chloride is surrounded by another molecule of hydrogen chloride, it will begin to align according to attraction forces. Opposite ends will neighbour each other, as seen in Figure 3.5.2.

A lewis dot diagrams of two hydrogen chloride molecules with a partial positive charge on the hydrogen atoms and the partial negative charge on the chlorine atoms. In the middle of the molecules is a series of closely spaced vertical lines located between the electronegative chlorine of the first molecule and the electropositive hyrodgen of the second molecule, representing a magnetic attraction force between the two atoms.
Figure 3.5.2: Hydrogen chloride molecules will begin to align positive and negative partial charges with each other, similar behaviour to that experienced by a magnet.

This is known as a dipole-dipole interaction – where molecules experience attraction and repulsion forces due to a permanent charge. Hydrogen bonds are a special kind of dipole-dipole interaction, as hydrogen exhibits unusually strong intermolecular forces when paired with a highly electronegative atom such as chlorine or oxygen (see Figure 3.5.3).

Water (ice) molecules aligned in a 3D lattice. The smaller white hydrogen atoms point towards nearby larger red oxygen atoms, with dotted lines representing attraction forces between them.
Figure 3.5.3: When water solidifies, hydrogen bonding between the molecules forces the molecules to line up in a way that creates empty space between the molecules, increasing the overall volume of the solid. This is why ice is less dense than liquid water. Image attribution: Chem&121: Introduction to Chemistry Copyright © 2023 by Lake Washington Institute of Technology is licensed under a Creative Commons Attribution-NonCommercial-ShareAlike 4.0 International License.
Diagram of two adjacent polar molecules (illustrated as colourful ovals) showing the distribution of charges with a delta plus on the hydrogen end and a delta minus on the chlorine end, connected by a dashed line representing the polar attraction.
Figure 3.5.4: Dipole-Dipole Interactions. Oppositely charged ends of polar molecules attract each other. Image attribution: Chem&121: Introduction to Chemistry Copyright © 2023 by Lake Washington Institute of Technology is licensed under a Creative Commons Attribution-NonCommercial-ShareAlike 4.0 International License.

Intermolecular forces allow us to understand why molecules have a tendency to stay together (see Figure 3.5.3). It also explains the melting and boiling points of different compounds. Molecules that are closely attracted to one another require higher kinetic energy and, therefore, higher temperature to separate and undergo a phase change. The greater the strength of the attraction and the more interactions that are allowed to occur, the greater the energy requirement. This is why ionic compounds (particularly salts) generally have high melting points. For reference, sodium chloride has a melting point of 802°C and a boiling point of 1465°C!

London Dispersion Forces

Under this theory, we should expect that for non-polar covalent molecules that don’t hold permanent dipoles, melting and boiling points should be the same due to a lack of partial charge. However, as Table 3.5.1 demonstrates, boiling points do change with respect to the length of the molecule.

Table 3.5.1: Alcohol Table of Boiling Points (cited from W.M. Haynes CRC Handbook of Chemistry 97th Edu.).
Chemical Name Molecular Formula Boiling Point (°C)
Methane [latex]\ce{CH_4}[/latex] -161.5
Ethane [latex]\ce{C_2H_6}[/latex] -88.6
Propane [latex]\ce{C_3H_8}[/latex] -42.1
Butane [latex]\ce{C_4H_10}[/latex] -0.5
Pentane [latex]\ce{C_5H_12}[/latex] 36.1
Hexane [latex]\ce{C_6H_14}[/latex] 68.7
Heptane [latex]\ce{C_7H_16}[/latex] 98.4
Octane [latex]\ce{C_8H_18}[/latex] 125.6
Nonane [latex]\ce{C_9H_20}[/latex] 150.8
Decane [latex]\ce{C_10H_22}[/latex] 174.1

All molecules, regardless of dipoles, are subject to dispersion forces (also known as London forces). These occur due to electrons continuously shifting within molecules. These electrons may occasionally generate partial charges, which induce electrons in other atoms to do the same, generating a slightly attractive force between molecules. As the surface area of a molecule increases, so does the occurrence of dispersion forces. As a result, longer-chain alkanes (listed in Table 3.5.1) possess higher boiling points.

Those are the three kinds of intermolecular forces: hydrogen bonding, dipole-dipole interactions and London dispersion forces. You can see them listed in Table 3.5.2 and their relative strengths. In addition to dictating boiling and melting points within molecules, they also play an important part in the solubility of molecules within water.

Table 3.5.2: Types of Intermolecular Forces.
Intermolecular Force Type of Molecule Strength of Attraction
Hydrogen Bonding Molecules containing H-X[1] Very Strong
Dipole-Dipole Interactions Polar Molecules Strong
(London) Dispersion Forces All molecules. Weak

Key Takeaways

  • Intramolecular forces are types of bonding (ionic, covalent and metallic), while intermolecular forces are the attraction and repulsion experienced between molecules.
  • Dipole-dipole interactions are experienced between molecules with partial charges.
  • Hydrogen bonding are a form of dipole-dipole interactions, but only feature on molecules with hydrogen next to a highly electronegative element.
  • London dispersion forces are felt by all molecules.


Practice Questions



  1. X = Electronegative Atom (O, N, Cl, Br, F, etc.)


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