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# 16.2: Intermolecular Forces of Attraction

Created by: CK-12

## Lesson Objectives

The student will:

• explain the difference between intermolecular and intramolecular forces of attraction.
• name and describe the types of intramolecular forces that hold groups of molecules together.
• explain how dipole-dipole forces differ from hydrogen bonds and London dispersion forces.
• identify liquids whose intermolecular forces of attraction are due to London dispersion forces, polar attractions, and hydrogen bonding given the formulas, structural data, and phase change points.
• describe some of the unique properties of water that are due to hydrogen bonding.

## Vocabulary

• hydrogen bond
• London dispersion forces

## Introduction

Attractive forces on the molecular level are divided into two categories – the forces inside a molecule holding atoms together, and the forces between molecules holding molecules together. In the image below, the forces holding the oxygen and hydrogen atoms together are the intramolecular forces. Intramolecular forces are the forces inside the molecule and consist of ionic and covalent bonds. The forces that hold the $\mathrm{H}_2\mathrm{O}$ molecules together as either liquid water or ice are the intermolecular forces. Intermolecular forces are the forces between molecules and are responsible for holding molecules in the solid or liquid state when no covalent or ionic bonds are possible. You can relate these two types of forces to bus systems where intra-city buses move people around inside one city and inter-city buses move people from one city to another.

When a substance in the solid phase is heated sufficiently, the molecular motions increase to the point that they overcome the forces holding the solid together, so the solid melts to liquid form. If the liquid is then heated sufficiently, the molecular motion increases to the point that it completely overcomes the attractive forces, so the liquid will change to the gaseous phase. The reverse of this process occurs when the substance is cooled. The conversion of solid to liquid, liquid to gas, gas to liquid, and liquid to solid are called phase changes.

The intermolecular forces must be overcome during a phase change. Therefore, the stronger the intermolecular forces of attraction, the greater the molecular motion (temperature) required to overcome them. The solids and liquids with the strongest intermolecular forces of attraction will have the highest melting and boiling points, and vice versa. The phase change temperatures for the various types of solids and liquids cover a very wide range. Substances with very weak forces, like helium, will melt at only a couple of degrees above absolute zero, whereas solid substances like asbestos and diamond do not melt until the temperature is in excess of $3500^\circ\text{C}$.

## London Dispersion Forces

The weakest type of intermolecular force is called London dispersion forces. London dispersion forces occur between all atoms and molecules, but they are so weak, they are only considered when there is no other intermolecular forces. For example, London dispersion forces exist between water molecules, but water molecules also have a permanent polar attraction so much stronger than the London dispersion forces that the London dispersion force is insignificant and not mentioned.

The cause of London dispersion forces is not obvious. Although we usually assume that the electrons of an atom are uniformly distributed around the nucleus, this is not true at every instance. As the electrons move around the nucleus, at a given instance, more electrons may be on one side of the nucleus than the other. This momentary nonsymmetrical electron distribution can produce a temporary dipolar arrangement of charge. This temporary dipole can induce a similar dipole in a neighboring atom and produce a weak, short-lived attraction.

The cases where London dispersion forces would be considered as the only intermolecular force of attraction would be for the noble gases and non-polar molecules such as helium, neon, argon, krypton, xenon, hydrogen, oxygen, methane, carbon dioxide, and so forth. Since non-polar molecules do not have a permanent dipole and no further bonding capacity, their only means of attracting each other is through London dispersion forces. Some of the substances whose intermolecular forces of attraction are London dispersion forces are held in the liquid state so weakly, they have the lowest melting points of all substances (see Table below for examples).

Boiling Points of Some London Dispersion Forces Liquids
Substance Chemical Symbol Boiling Point, $^\circ\mathrm{C}$
Helium He $-269.7$
Neon $\mathrm{Ne}$ $-248.6$
Argon $\mathrm{Ar}$ $-189.4$
Krypton $\mathrm{Kr}$ $-157.3$
Xenon $\mathrm{Xe}$ $-111.9$
Hydrogen $\mathrm{H}_2$ $-253$
Oxygen $\mathrm{O}_2$ $-182$
Methane $\mathrm{CH}_4$ $-161$
Carbon Dioxide $\mathrm{CO}_2$ $-78$

The temporary dipoles that cause London dispersion forces are affected by the molar mass of the particle. The greater the molar mass of the particle, the greater the force of attraction caused by London dispersion forces. The molar masses of $\mathrm{H}_2$, $\mathrm{N}_2$, and $\mathrm{O}_2$ are $2, \ 28, \ \mathrm{and} \ 32 \ \mathrm{g/mol}$, respectively, and their boiling points increase in similar fashion; $-253^\circ\mathrm{C}$ for $\mathrm{H}_2$, $-196^\circ\mathrm{C}$ for $\mathrm{N}_2$, and $-183^\circ\mathrm{C}$ for $\mathrm{O}_2$. For molecules with a high molecular weight, the London dispersion forces become strong enough that the substance will be a liquid or solid even at room temperature. Carbon tetrachloride, molar mass $154 \ \mathrm{g/mol}$, and bromine, molar mass $160 \ \mathrm{g/mol}$, boil at $+77^\circ\mathrm{C}$ and $+59^\circ\mathrm{C}$, respectively. Many long carbon chain, non-polar substances such as gasoline and oil remain liquids at common temperatures.

## Dipole-Dipole Interactions

When covalent bonds form between identical atoms, such as in $\mathrm{H}_2$, $\mathrm{N}_2$, $\mathrm{O}_2$, and so on, the electrons shared in the bonds are shared equally. The two atoms have the same electronegativity and therefore the same pull on the shared electrons (as illustrated in the figure below).

The center of negative charge for the entire molecule will be in the exact center of the molecule. This will coincide with the center of positive charge for the molecule. When the center of negative charge and the center of positive charge coincide, there is no charge separation and no dipole.

In the case of a symmetrical molecule with polar bonds, like the one shown below, the symmetry of the electron displacements will also keep the center of negative charge in the center of the molecule, which coincides with the center of positive charge. As a result, no dipole will occur.

If the two atoms sharing the bonding pair of electrons are not of the same element, the atom with the greater electronegativity will pull the shared electrons closer to it. Because of the resulting uneven distribution of electrons, the center of negative charge will not coincide with the center of positive charge and a dipole is created on the molecule. When the centers of positive and negative charge do not coincide, a charge separation exists and a dipole is present. For example, in the $\mathrm{CO}_2$ molecule above, both carbon-oxygen bonds are polar and the bonding electrons are shifted toward the oxygen.

The end of the molecule with the more electronegative atom will have a partial negative charge, and the end of the molecule with the more electropositive atom will have a slight positive charge. The symbols $\delta ^+$ and $\delta ^-$, as illustrated in the figure below, are used because these are not full positive and negative charges.

This polarity is much less than the charge separation in an ionic bond. These charges are only fractions of the full “$+1$” and “$-1$” charges. How much polarity a bond will experience depends on the difference in the electronegativities of the atoms.

For molecules that have a permanent dipole, the attraction between oppositely charged ends of adjacent molecules are the dominant intermolecular force of attraction. The figure below represents a polar solid; a polar liquid would look similar except the molecules would be less organized. On average, these polar attractions are stronger than London dispersion forces, so polar molecules in general have higher boiling points than London dispersion liquids. There is significant overlap, however, between the boiling points of the stronger London dispersion molecules and the weaker polar molecules.

The organization of a substance composed of polar molecules depends on the competition between the strength of the polar attractions and the molecular motion of the molecules. At higher temperatures, the molecular motion of the molecules is strong enough to disrupt the polar attractions, but at low temperatures, the molecular motion is reduced so that the polar attractions can hold the molecules in a structured arrangement.

In liquid and gaseous forms, the molecules can also turn freely. This turning of polar molecules can be seen in a macroscopic situation. If we bring a charged object (rubber comb run through hair, balloon rubbed on wool sweater, etc) near a very thin stream of water running from a faucet, the stream will bend its path toward the charged object. It doesn't make any difference if the charged object is positively or negatively charged because the water molecules in the stream will turn their oppositely charged ends toward the charged object. In the sketch above, you can see that the path of a non-polar liquid is not deflected by the charged rod, but the path of the water stream is deflected by the charged rod.

## Hydrogen Bonds

There are several polar molecules whose polar bonds are so strong they merit separate attention. These are the polar attractions that occur in molecules where hydrogen is bonded to nitrogen, oxygen, or fluorine. The polar attractions in these molecules are nearly 10 times as strong as regular polar attractions. These extra strong polar attractions that occur with $\mathrm{H}-\mathrm{N}$, $\mathrm{H}-\mathrm{O}$, and $\mathrm{H}-\mathrm{F}$ bonds are called hydrogen bonds, which distinguishes them from regular polar attractions. Keep in mind, however, that they are still polar attractions, albeit very strong ones.

There is more than one explanation for why these three combinations form hydrogen bonds. There are only ten elements that have greater electronegativity than hydrogen, and only four that have a significantly greater electronegativity than hydrogen. Three of the elements that have significantly greater electronegativity than hydrogen are nitrogen, oxygen, and fluorine – the three elements that form hydrogen bonds in compounds with hydrogen. When hydrogen bonds with atoms whose electronegativities are less than or equal to the electronegativity of hydrogen, the other atom cannot pull the shared electrons away from hydrogen, as seen in the figure below.

When hydrogen chemically bonds with nitrogen, oxygen, or fluorine, the very high electronegativities of these atoms can pull the electrons far away from the hydrogen atom, thus removing the shielding electrons from the proton nucleus of hydrogen. As a result, when the polar attractions between hydrogen bonding compounds form, the negative end of a molecule can get very close to the proton on the positive end of another molecule because there are no electrons for shielding. The closeness of the charges causes the extra strong polar attractions in these compounds. The characteristics of a liquid that forms hydrogen bonds are significantly different from similar compounds that do not form hydrogen bonds.

A homologous series of compounds are compounds where the elements of a family are each bounded to the same element. For example, family 4A in the periodic table consists of carbon, silicon, germanium, and tin. If each of these is bounded to hydrogen, it would produce a homologous series, $\mathrm{CH}_4$, $\mathrm{SiH}_4$, $\mathrm{GeH}_4$, and $\mathrm{SnH}_4$. If we graph the boiling points of this homologous series, we would get the graph sketched on the left side in the figure below. A large majority of the graphs plotting the boiling points of a homologous series would look like the one below, where the boiling points increase as molar mass increases.

However, if we graph the boiling points of the homologous series of family 6A combined with hydrogen, we get quite a different graph, as seen on the right side of the sketch. The higher molar mass compounds in the series follow the normal pattern where the boiling points decrease as the molar masses decrease. When we get to water, the boiling point suddenly increases – it is more than $150^\circ$ higher. Hydrogen bonding explains why liquid water is held together far more tightly than expected. Graphs of the boiling points for the homologous series of hydrogen with 5A family members and hydrogen with 7A family members would be similar, with the boiling points of $\mathrm{NH}_3$ and $\mathrm{HF}$ being greatly different from what would be expected.

Significant hydrogen bonding causes water molecules to line up end-to-end. The fact that water forms hydrogen bonds has effects so large that it is impossible in this text to delineate them all. We will consider just a few of these effects here.

1. The normal boiling point of water is $100^\circ\mathrm{C}$. If water did not form hydrogen bonds and had a regular polar attraction between molecules, its boiling point would be somewhere around $-60^\circ\mathrm{C}$. The average temperature of the Earth's surface is $-15^\circ\mathrm{C}$, so if water did not form hydrogen bonds, the oceans and lakes would vaporize. Therefore, Earth would not be a watery planet and would not likely be a planet with life on it.

2. For almost all substances, the substance contracts as it cools – the molecules move around slower, and the intermolecular forces pull the molecules closer together. As a result, the solids are denser than the liquids, and the solids will sink in the liquids. Water, of course, is the exception. When water cools to its freezing point and solidifies, it expands. The molecular motion for water molecules at or above $4^\circ\mathrm{C}$ is sufficient to keep the water molecules in a molecular complex with large holes in the structure. When water is cooled below $4^\circ\mathrm{C}$, the molecular motion is inadequate to break up this complex structure, and the water molecules expand because of the holes in the structure. Therefore, solid water is less dense than liquid water, and ice floats on water.

One of the consequences of this effect is that when cold weather comes to areas in the northern and southern parts of the earth, the cold air freezes the surface water it comes into contact with. If water were like other substances, then the solid would sink to the bottom. The solid would continuously sink to the bottom as the cold air continues to freeze the new surface until the entire lake would be frozen from top to bottom. No water dwelling animals would be able to survive such an occurrence. In reality, however, when the cold air freezes the surface of a lake, the ice floats, stays on top, and insulates the rest of the water from the cold air. Only the surface freezes, allowing the animals that live in the water to survive the winter.

3. One factor that affects the weathering of rocks in some geographical areas is that rain water can cause rocks to fracture. When it rains, water can enter the cracks in rocks. During the winter, the water can freeze and expand, causing the rocks to fracture.

4. Some biologically active molecules, such as DNA, require a particular shape for their function. At points along its length, such a molecule can be linked to itself with different types of attractions – one of which are hydrogen bonds, as seen in the figure below.

5. As long as water is above $4^\circ\mathrm{C}$, it contracts and becomes denser upon cooling, just like other substances. At temperatures less than $4^\circ\mathrm{C}$, significant hydrogen bonding begins to form, causing water to expand and become less dense. The maximum density for water is at $4^\circ\mathrm{C}$.

Many animals that live in water and require oxygen use oxygen that is dissolved in the water. For deep lakes, diffusion is inadequate to move the oxygenated water to the bottom of the lake. Instead, water in lakes becomes oxygenated (has dissolved oxygen) by the action of wind and waves at the surface. For lakes in northern climates, the surface passes through the temperature $4^\circ\mathrm{C}$ twice a year, once as it cools in the fall and once as it warms in the spring. During these two times, the oxygenated water at the surface would sink to the bottom because it is denser. These periods are called spring and fall “turnovers.” The turnovers provide oxygenated water at the bottom of the lakes.

This video shows and narrates by hydrogen bonds form and shows variations in boiling point of a homologous series due to hydrogen bonding (2h): http://www.youtube.com/watch?v=LGwyBeuVjhU (1:40).

## Lesson Summary

• Molecules are held together in the liquid phase by intermolecular forces of attractions.
• London dispersion forces are a very weak intermolecular force of attraction caused by a temporary electrostatic attraction between the electrons of one molecule or atom and the nucleus of another.
• Polar attractions are a type of intermolecular force of attraction caused by the electrostatic attraction between permanent dipoles that exists on polar molecules.
• Hydrogen bonds are an exceptionally strong type of polar attraction that occurs between molecules that have $\mathrm{H}-\mathrm{F}$ bonds, $\mathrm{H}-\mathrm{O}$ bonds, or $\mathrm{H}-\mathrm{N}$ bonds.
• Hydrogen bonds are responsible for the unique characteristics of water.

The video below examines the various intermolecular and intramolecular forces of attraction.

## Review Questions

1. Identify the most important type of inter-particle force present in the following solids that is responsible for binding the particles into a solid.
1. $\mathrm{He}$
2. $\mathrm{NO}$
3. $\mathrm{HF}$
4. $\mathrm{CH}_4$
5. $\mathrm{CO}_2$
6. $\mathrm{CHCl}_3$
2. Predict which substance in the following pairs would have the stronger force of attraction between molecules and justify your answer.
1. $\mathrm{CO}_2$ or $\mathrm{OCS}$
2. $\mathrm{PF}_3$ or $\mathrm{PF}_5$
3. $\mathrm{H}_2\mathrm{O}$ or $\mathrm{H}_2\mathrm{S}$

Feb 23, 2012

Nov 26, 2014