Intermolecular Forces

Now that we have learned out to identify what shape a molecule has and whether or not it is polar, let's revisit the question we posed earlier in the lesson.

We discussed the fact that salt (NaCl) has a melting point so much higher than water (H₂O). We explained this on the basis of salt being a network, held together by strong ionic bonds, whereas water's covalent bonds only operate within any given molecule, not between molecules.

But what is it that does hold water molecules together? What is it that stops water molecules from flying apart from each other and being a gas? And why is it that water has a much higher boiling point than other small molecular substances, such as carbon dioxide and methane? These questions and others require us to learn about the different types of intermolecular forces.

Nature of Intermolecular Forces

Intermolecular forces (IMFs) are forces of attraction that operate between two molecules (or two free atoms) that are in contact with each other. They are represented above as the faint dotted lines connecting the water molecules.

They are drawn faintly for a reason: they are much weaker than the ionic, covalent, and molecular bonds we learned about in the last two lessons. Breaking ionic or covalent bonds usually requires lots of energy, such as the heat of an open flame. Breaking IMFs is as easy as leaving an ice cube on the counter to melt.

However, just because IMFs are relatively weak does not mean they are unimportant. As we will see, IMFs come in several different types, some stronger than others; and the type of IMFs a substance has will help determine its melting and boiling points, as well as its ability to mix with other substances.

Types of Intermolecular Forces

There are three types of IMF you will need to know. By the end of this section, you should be able to do the following:

Look at a formula for a molecular compound and draw its Lewis structure
From the Lewis structure, determine its molecular shape and whether or not it is polar
From its polarity and structure, determine the types of IMFs it will experience

As you can see, this lesson is really a continuation of the previous two lessons. It is vital that you feel confident about those lessons, especially Lesson 8, in order to succeed here.

Dipole-Dipole Forces

This figure shows a simple polar molecule: two atoms sharing electrons unequally. The more electronegative atom on the right has a greater share of the electrons and thus a denser electron cloud. The resulting dipole can be shown using plus and minus signs, or by an arrow. The Cl end of the molecule is partially negative, and the H is partially positive.

Although we use plus and minus signs to show that the left side of this molecule is positively charged compared to the right side, the two atoms do not have full charges of +1 and -1 as they would if this were an ionic bond. The symbol δ indicates that the positive and negative charges are fractions of a full charge. The closer together the two atoms are in electronegativity, the smaller the fraction of charge each atom has and the less polar the bond.

When two polar molecules like HCl get near each other, the positive end of one attracts the negative end of the other. This attraction between the two dipoles is known as a dipole-dipole force. Compared to ionic or covalent bonds, dipole-dipole forces are very weak (because the partial charges are quite small).

Compared to other IMFs, they are moderate in their strength. We will see examples both of forces that are stronger and forces that are weaker.

Hydrogen "Bonds"

When the polarity of a molecule is caused by a bond between a hydrogen atom and either a nitrogen, an oxygen, or a fluorine atom, the resulting IMFs are unusually strong and so are given a special name: hydrogen bonds. Note that, even though they are called "bonds," they do not fall into the same category as ionic or covalent bonds. They are IMFs.

Hydrogen bonds are the strongest of dipole-dipole forces, so much so that we give them their own special category. F, O, and N are the three most electronegative atoms in the periodic table and this is part of the reason why the bonds between hydrogen and F, O, or N are the most polar of covalent bonds. The strong polarity of these bonds results in strong dipole-dipole forces between the molecules.

The only molecule that contains an H-F bond is HF. On the other hand, there are many molecules that contain H-O and/or H-N bonds are are therefore capable of intermolecular hydrogen bonding.

Important: in order for hydrogen bonding to occur, there must be a direct bond from H to O, N, or F. In dimethyl ether, shown at right, there is no O-H bond so the molecule has no hydrogen bonding IMFs. Ethanol, however, would.

Dispersion Forces

This image shows the average location of the electrons in a nonpolar molecule. The electrons are shared equally, neither end of the molecule is positive or negative, so no dipole-dipole bonding is possible. Nonpolar bonds like these are found when two atoms bonded to one another are the same element, or in the case of a C-H bond.

But electrons are in constant motion, and there are brief periods of time in which they are slightly to one side of the molecule or the other. During this time, the molecule has a very slight, temporary dipole. If another, similar, temporary dipole exists in a neighboring molecule, the weak mutual attraction causes the two dipoles to persist for slightly longer than they otherwise would, and the two molecules remain attracted for a short period of time.

These attractive forces are weak and very short-lived. But they do create an appreciable force between molecules, known as a dispersion force. Dispersion forces are the weakest type of IMF. While they are technically present in all substances, their effect is only important in nonpolar substances. This includes elemental substances like N₂ and P₄, as well as symmetric molecules like BF₃ and CH₄.

Practice

As usual, you should practice determining IMF types for various substances in your lab workbook. In the next two sections, we will conclude the lesson by applying our knowledge of IMF types to the basic physical properties of substances.

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