Section 1

Intermolecular Forces

A molecular diagram illustrating dipole-dipole interactions between polar molecules. Each molecule is shown with a red (δ⁺) and blue (δ⁻) end, indicating partial positive and negative charges. Green arrows between opposite charges represent attractive forces, while gray arrows between like charges represent repulsive forces. The network of arrows shows how molecules align based on electrostatic interactions.

Introduction

The individual atoms and molecules of a material are far too small to be perceived with our unaided senses. However, what we can see, handle, feel, and study are materials, which are a collection of atoms and molecules that comprise the material as a whole. Both sugar and salt are composed of countless individual molecules and ions, but their observed physical properties are the sum total of the individual interactions between individual particles. The intermolecular forces responsible for these interactions are electrical in nature and govern the strength of attraction between the individual particles of a material.

Introduction to Intermolecular Forces Transcript

Ion-Ion and Dipole-Dipole Interactions

The particles in salt (NaCl) are ions, and it is the ion–ion interactions of the oppositely charged sodium cations and chloride anions that hold salt together in its lattice arrangement. However, if the particles are molecules and not ions, there are multiple ways individual molecules can interact with each other. All of these intermolecular interactions involve some type of attraction between oppositely charged portions of the molecules that make up the material. If the molecule contains polar bonds, and if the overall arrangement of these bonds is such that the molecule as a whole is polar, the molecule has a dipole associated with it. Dipole means that one part of the molecule carries a partial positive charge, and another part carries a partial negative charge. When this is the case, oppositely charged parts of adjacent molecules form dipole–dipole interactions. The stronger or more pronounced the dipole associated with individual molecules, the stronger the interaction between the molecules, because the dipole–dipole forces are stronger.

A diagram comparing two types of intermolecular interactions. In part (a), elongated molecules are aligned such that the partially positive ends (δ⁺) of one molecule face the partially negative ends (δ⁻) of another, illustrating dipole-dipole interactions. In part (b), rounder molecules are aligned similarly, also showing attraction between opposite partial charges. Both examples highlight the directional nature of dipole-dipole forces between polar molecules.

Bond Polarity Transcript

Dipole-Dipole Interaction Transcript

Electronegativity Transcript

Molecular Polarity Transcript

Bonding and Polarity Review Transcript

Ion-Ion and Dipole-Dipole Interactions

Hydrogen Bonds—Exceptionally Strong Dipole–Dipole Interactions

A molecular model showing hydrogen bonding between water (H₂O) molecules. Each water molecule consists of one oxygen atom (red) and two hydrogen atoms (white). Partial charges are labeled: δ⁻ on the oxygen atoms and δ⁺ on the hydrogen atoms. Dotted lines represent hydrogen bonds formed between the hydrogen of one water molecule and the oxygen of another, illustrating the intermolecular attractions responsible for water's unique properties.

Between them is quite polar. This is the case when nitrogen (electronegativity 3.0), oxygen (electronegativity 3.5), and fluorine (electronegativity 4.0) bond with hydrogen (electronegativity 2.1). The force of electrical attraction depends not only on the amount of charge but also on the distance between the charges. The strong polarity associated with these bonds, coupled with the fact that the hydrogen atom is very small, results in an exceptionally strong dipole-dipole interaction known as a hydrogen bond. However, because hydrogen bonds are about two times stronger than a typical dipole–dipole interaction, they are identified and classified separately. It is the hydrogen bonds of the OH groups between water molecules that are responsible for the unique physical properties of water, and it is the hydrogen bonds that form between the A–T and G–C base pairs in DNA that are responsible for accurately copying and reading the messages of life.

Hydrogen Bonding Transcript

Hydrogen Bond Practice

Ion-Dipole Interactions

A molecule with a dipole, which has regions of partial overall negative and positive charge, can interact with an ion. These ion–dipole interactions are weaker than the ion- ion interactions associated with ionic compounds, but they are stronger than dipole– dipole interactions. Ion–dipole interactions are especially important in water solutions containing ions, (e.g., a solution of salt dissolved in water). In solutions, the individual Na+ ions interact with the slightly negatively charged oxygen atoms of the water molecules, and Cl– interacts with the slightly positively charged hydrogen atoms of the water molecules.

A diagram showing the hydration of a sodium ion (Na⁺) by water molecules. The central Na⁺ ion is surrounded by six water molecules. Each water molecule is oriented with its oxygen atom (red sphere) facing the sodium ion, and the hydrogen atoms (white spheres) pointing outward. This arrangement illustrates ion-dipole interactions, where the partially negative oxygen atoms are attracted to the positively charged sodium ion.

Ion-Dipole Interactions Transcript

London Dispersion Forces and Induced Dipole–Induced Dipole Interactions

Individual non-polar molecules can interact with each other through what are called London dispersion forces or induced dipole–induced dipole interactions. If the valence electrons in a non-polar molecule are dispersed around a molecule such that they spend more time on one side of the molecule than the other, a transitory slight negative charge builds up on one side of the molecule while a transitory positive charge builds on the opposite side of the molecule. When this is the case, the negatively charged side of the molecule can interact with a neighboring molecule, repelling the valence electrons of the neighbor and inducing a slight positive charge on the neighboring molecule. This results in a slight attraction between the molecules. These induced dipole–induced dipole interactions are much weaker than dipole–dipole interactions, but since there are typically many of these, especially in larger molecules, they are significant.

A diagram illustrating how London dispersion forces (temporary dipole interactions) arise between helium atoms. The first image shows a helium atom with an uneven distribution of electrons. The second image shows the same atom with a temporary (instantaneous) dipole, where one side is slightly negative (−) and the other slightly positive (+). The third image shows a neighboring helium atom developing an induced dipole in response, with opposite charges aligned. This demonstrates how momentary electron shifts can cause attraction between nonpolar atoms.



A diagram showing how London dispersion forces arise through temporary dipoles in nonpolar atoms and molecules. In part (a), two helium atoms are shown. Initially, there is no polarization. A temporary uneven electron distribution creates an instantaneous dipole in atom A (δ⁻ and δ⁺ ends), which induces a dipole in atom B. In part (b), the same process is shown with two hydrogen (H₂) molecules. An instantaneous dipole forms in molecule A, which then induces a dipole in molecule B. The diagram illustrates how transient dipoles lead to induced dipoles and weak intermolecular attractions.


Induced Dipole-Induced Dipole Interactions Transcript

Strength of Dispersion Forces Transcript

London Dispersion and Induced Dipole-Induced Dipole Interactions Practice


Summary of Intermolecular Forces

Collectively, all of the forces between the individual particles in a material are referred to as intermolecular forces, which collectively are the attractive forces between molecules and ions. It is the type and consequently, the relative strength of the intermolecular forces in a material that determine the physical properties and behavior of the material.

Reviewing Forces Transcript

Fundamental Knowledge and Skills - Intermolecular Forces

What You Need to Know :

You need to understand that the strength of intermolecular interactions is primarily due to the charges present on those molecules and the distance between molecules. You should be able to predict the relative strength of intermolecular interactions and, if given two molecules, you should also be able to predict which molecule will have the higher boiling point and water solubility. The physical properties of a material are the direct result of the strength of the intermolecular forces present in the material.

How to Learn It:

The strength of the interaction between two particles is due to the relative charges of the particles and the distance between them. Thus, the higher the charge, the stronger the interaction and the closer the particles are to each other, the stronger the interaction. Intermolecular interactions are classified by their relative strength. From strongest to weakest they are, Ion-Ion forces, Ion-Dipole, Dipole-Dipole, Hydrogen Bonding, Dipole-Induced Dipole, and Induced Dipole-Induced Dipole.

Ion-Ion intermolecular forces occur when two ions interact. This interaction is between two separate ions and is not to be confused with lattice energy, which is the sum total of the energy associated with interaction of an array of ions in a lattice.

Ion-Dipole intermolecular forces are between a charged ion and a polar molecule. A Dipole is a molecule that has two poles: portions of the molecule that are partially positively charged and portions of the molecule that are partially negatively charged due to the stringer attraction of the bonding valence electrons toward a more electronegative atom. These interactions are stronger than a dipole-dipole interaction because the charge associated with the ion is larger.

Dipole-Dipole intermolecular forces are between two molecules with permanent dipoles. The positive portion of one molecule that lacks electron density is attracted to the negative portion of another molecule that has extra electron density.

Hydrogen bonds are a special subset of Dipole-Dipole intermolecular forces. Hydrogen bonds are the interaction of a partially positively charged hydrogen bound to nitrogen, oxygen, or fluorine on one molecule to a partially negatively charged nitrogen, oxygen, or fluorine atom on a separate molecule. When hydrogen is bound to nitrogen, oxygen, or fluorine, there is a large electronegativity difference, and the bond is extremely polar. The vast majority of the electron density is centered on the more electronegative nitrogen, oxygen or fluorine atom. The polarity of a hydrogen atom bonded to these electronegative atoms, coupled with its small size, results in an especially strong dipole- dipole interaction. Hydrogen bonds are about twice as strong as non-hydrogen bond dipole-dipole interactions.

Dipole-Induced Dipole intermolecular forces are between one polar molecule and one non-polar molecule. The partial charge on the polar molecule forces the electrons on the non-polar molecule to shift, resulting in a weak temporary partial dipole itself, which, in turn, results in an attractive force between the molecules.

Induced Dipole-Induced Dipole intermolecular forces are between two non-polar molecules. There are many names for this phenomenon: Induced Dipole-Induced Dipole, London Dispersion Forces, van der Waals Forces. This has a mechanism that is very similar to a Dipole-Induced Dipole interaction. As the two non-polar molecules move closer to each other, the electrons in the molecules shift, forming very weak temporary dipoles. These dipoles act to attract the nonpolar molecules to each other, but they disappear as the molecules separate from each other. The strength of induced dipole-induced dipole forces increase with size and ‘overlappability’ of the molecules.

A diagram showing six types of intermolecular forces: ion-dipole (Na⁺ with water), hydrogen bonding (methanol with water), dipole-dipole (chloroform), ion-induced dipole (Cl⁻ with hexane), dipole-induced dipole (acetone with hexane), and dispersion forces (between two octane molecules).

A diagram showing six types of intermolecular forces using molecular models. From left to right, top to bottom:

  1. Ion-dipole interaction—Na⁺ ion surrounded by water molecules, with their oxygen atoms oriented toward the ion.

  2. Hydrogen bond—between methanol (CH₃OH) and water (H₂O), showing a dotted line from hydrogen to oxygen.

  3. Dipole-dipole interaction—between a polar CH₃OH molecule and chloroform (CHCl₃), with opposite partial charges aligned.

  4. Ion-induced dipole—Cl⁻ ion near a hexane (C₆H₁₄) molecule, inducing polarization.

  5. Dipole-induced dipole—acetone (CH₃COCH₃), a polar molecule, polarizing nonpolar hexane.

  6. Dispersion forces—between two nonpolar octane (C₈H₁₈) molecules showing temporary electron shift alignment.

Intermolecular forces impact many aspects of everyday life, including phase changes. A collection of molecules (such as a bucket of water) will form a solid, liquid, or gas depending on the temperature of those molecules and the intermolecular forces between them. Temperature is essentially a measure of the kinetic energy of a collection of molecules: higher temperatures have more kinetic energy and are thus moving more quickly. When temperature is low, solids form because the molecules are moving slowly and the intermolecular attractive forces overcome the motion of the molecules, arresting their motion and they form a crystal. When temperature is high, gases form because the molecules are moving quickly enough to overcome intermolecular attractive forces.

A diagram showing how molecular structure affects boiling point. In part (a), a series of alkanes—methane, ethane, propane, and n-butane—are shown with increasing molecular mass and boiling points: from −161.5 °C for methane to −0.5 °C for n-butane. In part (b), two molecules with the same molar mass (72 g/mol), 2,2-dimethylpropane (neopentane) and n-pentane, are compared. Despite identical mass, n-pentane has a much higher boiling point (36.1 °C vs. 9.5 °C) due to its elongated shape and greater surface area, which increases intermolecular dispersion forces.

The temperatures at which melting and boiling occur are directly related to the strength of the intermolecular forces. The stronger the intermolecular forces, the higher the melting point and the higher the boiling point. For example, most ionic salts (ion-ion intermolecular forces) have very high melting points. When considering the melting and boiling points of dipole-dipole interactions, look for hydrogen bonding capability. More hydrogen bonding increases the strength of intermolecular forces and raises melting and boiling points.

If intermolecular forces are weak, then their melting and boiling points are very low. Helium, which has the weakest induced dipole-induced dipole interactions of any compound, has a melting point of 1.3 Kelvin (-272.2ᵒC or -452ᵒF) and a boiling point of 4.2 Kelvin.

A diagram showing multiple water (H₂O) molecules forming hydrogen bonds. Each water molecule has one oxygen atom (red) and two hydrogen atoms (white). The molecules are connected by dotted lines representing hydrogen bonds, which occur between the hydrogen atom of one molecule and the oxygen atom of another. This network illustrates the strong intermolecular attractions that give water its unique physical properties.

Boiling points are also higher when the intermolecular forces between molecules are stronger. In particular, straight-chain non-polar molecules with less branches are able to interact with each other in more places in the liquid state, increasing the strength of the London Dispersion forces. Highly branched non-polar molecules, however, are unable to interact with each other as much as non-branched molecules. I like to think of these non-branched structures as being able to ‘tesselate’ with each other and thus interact more. You could even think of it as the molecules ‘cuddling’ with each other, increasing the strength of their ‘relationship’.

This is a concept that many students struggle with- that branched molecules tend to have higher melting points and lower boiling points than straight-chain molecules.

A diagram comparing the boiling points of different hydrocarbons based on molecular shape and surface area. On the left, methane (CH₄) molecules are small and spherical, with a boiling point of −161 °C. In the center, n-pentane (a straight-chain hydrocarbon) has a boiling point of 36 °C due to greater surface area and stronger dispersion forces. On the right, 2,2-dimethylpropane (neopentane), a more compact and branched molecule with the same formula as n-pentane, has a lower boiling point of 9.5 °C due to its reduced surface area and weaker intermolecular interactions.

The types of intermolecular forces matter too. Oil and water do not mix primarily because water is a polar molecule, whereas oil molecules are long nonpolar hydrocarbon chains. The strength of an interaction of a polar molecule and a nonpolar molecule is rather low compared to the interaction of polar molecules with polar molecules. Thus, the energetic ‘drive’ for like molecules to associate with like molecules is higher than the ‘drive’ for polar molecules to associate with nonpolar molecules. (This energetic drive is largely due to changes in entropy, which we will discuss more later).

Why It Matters:

Intermolecular forces impact your life on a daily basis. They are the primary factor in determining the observed physical behavior of everything. For example, London Dispersion Forces are presumed to be the primary explanation behind the stickiness of tape, and the ability of gecko feet to hold onto the ceiling and walls. The opposing strands of DNA are held together by hydrogen bonding between the nucleotide bases A-T and C-G. They are all important.


Additional Videos

Relative Strength of Intermolecular Forces Transcript

Molecular Interactions and Hydrogen Bonds Transcript

London Dispersion Forces Transcript

Intermolecular Forces Summary Practice

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