5.1: Forces between Molecules

Learning Objectives

• Describe the types of intermolecular forces possible between atoms or molecules in condensed phases (dispersion forces, dipole-dipole attractions, and hydrogen bonding)
• Identify the types of intermolecular forces experienced by specific molecules based on their structures

Solids and liquids have volumes that (mostly) do not change when compressed. In other words, you can try to squeeze them but they don't get smaller. On the other hand, gases are compressible. Solids also have a rigid shape, but the shapes of liquids and gases change to match their containers. These properties are illustrated in the Figure below. In the solid and liquid the particles (atoms or molecules) are touching one another, but in the liquid the particles are able to spread out along the bottom of the container. In the gas the particles are not touching and the gas fills up its container in all directions.

Figure \(\PageIndex{1}\): Comparison of the arrangement of particles in solids, liquids, and gases. Brightyellowjeans, CC BY-SA 4.0, via Wikimedia Commons

The particles that make up solids and liquids are touching one another due to intermolecular forces (IMF). The covalent bonds between atoms in a molecule, such as the bonds that hold H and O together in a water molecule, are intramolecular forces. The prefix intra means within whereas the prefix inter means between different copies. Intramolecular forces are those within the molecule that keep the molecule together. Intermolecular forces are the attractions between molecules, which determine many of the physical properties of a substance. Figure \(\PageIndex{2}\) illustrates these different molecular forces. The strengths of these attractive forces vary widely, though usually the IMFs between small molecules are weak compared to the intramolecular forces that bond atoms together within a molecule. For example, to overcome the IMFs in one mole of liquid HCl and convert it into gaseous HCl requires only about 17 kilojoules of energy. However, to break the covalent bonds between the hydrogen and chlorine atoms in one mole of HCl requires about 25 times more energy—430 kilojoules.

Figure \(\PageIndex{2}\) Intramolecular forces keep a molecule intact. Intermolecular forces hold multiple molecules together and determine many of a substance’s properties.

Dipole–Dipole Forces

Dipole-dipole forces are the attractive forces that occur between polar molecules (see figure below). A molecule of hydrogen chloride has a partially positive hydrogen atom and a partially negative chlorine atom. A collection of many hydrogen chloride molecules will align themselves so that the oppositely charged regions of neighboring molecules are near each other.

Figure \(\PageIndex{3}\): Dipole-dipole forces in hydrogen chloride. Dipole-dipole forces result from the attraction between the positive end of one dipole and the negative end of a neighboring dipole. Dipole-dipole forces are similar to ionic bonds, but because they involve only partial charges, they are much weaker.

Dispersion Forces

Dispersion forces are the weakest of all intermolecular forces. They are often called London forces after Fritz London (1900 - 1954), who first proposed their existence in 1930. London dispersion forces are intermolecular forces that occur between all atoms and molecules due to the random motion of electrons.

For example, the electron cloud of a helium atom contains two electrons, and, when averaged over time, these electrons will distribute themselves evenly around the nucleus. However, at any given moment, the electron distribution may be uneven, resulting in an instantaneous dipole. This weak and temporary dipole can subsequently influence neighboring helium atoms through electrostatic attraction and repulsion. The formation of an induced dipole is illustrated below.

Figure \(\PageIndex{4}\) Random fluctuations in the electron density within the electron cloud of a helium atom results in a short-lived ("instantaneous") dipole. The attractive force between instantaneous dipoles and the resulting induced dipoles in neighboring molecules is called the London dispersion force.

Dispersion forces that develop between atoms in different molecules can attract the two molecules to each other. The forces are relatively weak, however, and become significant only when the molecules are very close. For similar substances, London dispersion forces get stronger with increasing atomic or molecular size.

Hydrogen Bonds

The attractive force between water molecules is an unusually strong type of dipole-dipole interaction. Water is a great example of a substance that can do hydrogen bonding. Water contains hydrogen atoms that are bound to a highly electronegative oxygen atom, making for very polar bonds. The partially positive hydrogen atom of one molecule is then attracted to the oxygen atom of a nearby water molecule (see figure below).

Figure \(\PageIndex{6}\) A hydrogen bond in water occurs between the hydrogen atom of one water molecule and the lone pair of electrons on the oxygen atom of a neighboring water molecule.

A hydrogen bond is an intermolecular attractive force in which a hydrogen atom that is covalently bonded to a small, highly electronegative atom (N, O, or F) is attracted to a lone pair of electrons on an atom in a neighboring molecule. Hydrogen bonds are very strong compared to other dipole-dipole interactions, but still much weaker than a covalent bond. A typical hydrogen bond is about \(5\%\) as strong as a covalent bond.

Hydrogen bonding occurs only in molecules where hydrogen is covalently bonded to one of three elements: fluorine, oxygen, or nitrogen (N, O, or F). These three elements are so electronegative that they withdraw the majority of the electron density from the covalent bond with hydrogen, leaving the \(\ce{H}\) atom very electron-deficient. Because the hydrogen atom does not have any electrons other than the ones in the covalent bond, its positively charged nucleus is almost completely exposed, allowing strong attractions to other nearby lone pairs of electrons.

A molecule must contain an H-F, H-O, or H-N bond in order to initiate a hydrogen bond, but the H-F, H-O, or H-N bond itself is NOT the hydrogen bond. 

The hydrogen bonding that occurs in water leads to some unusual, but very important properties. Most molecular compounds that have a mass similar to water are gases at room temperature. However, because of the strong hydrogen bonds, water molecules are able to stay condensed in the liquid state. The figure below shows how its bent shape and the presence of two very polar bonds per molecule allows each water molecule to hydrogen bond with several other molecules.

Figure \(\PageIndex{7}\)Multiple hydrogen bonds occur simultaneously in water because of its bent shape and the presence of two very polar bonds per molecule.

In the liquid state, the hydrogen bonds of water can break and reform as the molecules flow from one place to another. When water is cooled, the molecules begin to slow down. Eventually, when water is frozen to ice, the hydrogen bonds become more rigid and form a well-defined network (Figure \(\PageIndex{8}\)). The bent shape of the molecules leads to gaps in the hydrogen bonding network of ice. Ice has the very unusual property that its solid state is less dense than its liquid state. As a result, ice floats in liquid water. Virtually all other substances are denser in the solid state than in the liquid state. Hydrogen bonds also play a very important biological role in the physical structures of proteins and nucleic acids.

Figure \(\PageIndex{8}\) The Hydrogen-Bonded Structure of Ice.

Hydrogen bond formation requires both a hydrogen bond donor and a hydrogen bond acceptor.

Because ice is less dense than liquid water, rivers, lakes, and oceans freeze from the top down. In fact, the ice forms a protective surface layer that insulates the rest of the water, allowing fish and other organisms to survive in the lower levels of a frozen lake or sea. If ice were denser than the liquid, the ice formed at the surface in cold weather would sink as fast as it formed. Bodies of water would freeze from the bottom up, which would be lethal for most aquatic creatures. The expansion of water when freezing also explains why automobile or boat engines must be protected by “antifreeze” and why unprotected pipes in houses break if they are allowed to freeze.

Example \(\PageIndex{1}\)

Considering , C₂H₆, Xe, and , which can form hydrogen bonds with themselves? Draw the hydrogen-bonded structures.

Given: compounds

Asked for: formation of hydrogen bonds and structure

Strategy:

1. Identify the compounds with a hydrogen atom attached to O, N, or F. These are likely to be able to act as hydrogen bond donors.
2. Of the compounds that can act as hydrogen bond donors, identify those that also contain lone pairs of electrons, which allow them to be hydrogen bond acceptors. If a substance is both a hydrogen donor and a hydrogen bond acceptor, draw a structure showing the hydrogen bonding.

Solution:

A Of the species listed, xenon (Xe), ethane (C₂H₆), and trimethylamine [(CH₃)₃N] do not contain a hydrogen atom attached to O, N, or F; hence they cannot act as hydrogen bond donors.

B The one compound that can act as a hydrogen bond donor, methanol (CH₃OH), contains both a hydrogen atom attached to O (making it a hydrogen bond donor) and two lone pairs of electrons on O (making it a hydrogen bond acceptor); methanol can thus form hydrogen bonds by acting as either a hydrogen bond donor or a hydrogen bond acceptor. The hydrogen-bonded structure of methanol is as follows:

Exercise \(\PageIndex{1}\)

Considering , , NH₃, and CH₃F, which can form hydrogen bonds with themselves? Draw the hydrogen-bonded structures.

Answer

CH₃CO₂H and NH₃;

 

Example \(\PageIndex{2}\):

Identify the most significant intermolecular force in each substance.

1. C₃H₈
2. 
3. H₂S

Solution

a. Because C–H bonds are nonpolar the most significant intermolecular force for this substance would be dispersion forces.

b. This molecule has an H atom bonded to an O atom, so it will experience hydrogen bonding.

c. Although this molecule does not experience hydrogen bonding, the Lewis electron dot diagram and VSEPR indicate that it is bent, so it has a permanent dipole. The most significant force in this substance is dipole-dipole interaction.

Exercise \(\PageIndex{2}\)

Identify the most significant intermolecular force in each substance.

1. HF
2. HCl

Answer a

hydrogen bonding

Answer b

dipole-dipole interactions

Summary

• Molecules in liquids are held to other molecules by intermolecular interactions, which are weaker than the intramolecular interactions that hold the atoms together within molecules and polyatomic ions.
• Dipole–dipole interactions arise from the electrostatic interactions of the positive and negative ends of molecules with permanent dipole moments.
• London dispersion forces are due to the formation of instantaneous dipole moments in polar and nonpolar molecules as a result of short-lived fluctuations of electron charge distribution, which in turn cause the temporary formation of an induced dipole in adjacent molecules.
• Hydrogen bonds are especially strong dipole–dipole interactions between molecules that have hydrogen bonded to a highly electronegative atom, such as O, N, or F. The resulting partially positively charged H atom on one molecule (the hydrogen bond donor) can interact strongly with a lone pair of electrons of a partially negatively charged O, N, or F atom on adjacent molecules (the hydrogen bond acceptor).

Contributors and Attributions

• TextMap: Chemistry the Central Science (Brown et al.

• Paul Flowers (University of North Carolina - Pembroke), Klaus Theopold (University of Delaware) and Richard Langley (Stephen F. Austin State University) with contributing authors. Textbook content produced by OpenStax College is licensed under a Creative Commons Attribution License 4.0 license. Download for free at http://cnx.org/contents/85abf193-2bd...a7ac8df6@9.110).

• Marisa Alviar-Agnew (Sacramento City College)

• Libretext: Chemistry for Allied Health (Soult)