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Intermolecular Forces

In the previous chapter, we discussed the different forces that exist between atoms (intramolecular forces). When atoms are joined to one another they form molecules, and these molecules in turn have forces that bind them together. These forces are known as intermolecular forces, and we are going to look at them in more detail in this next section.

Definition 1: Intermolecular forces

Intermolecular forces are forces that act between stable molecules.

You will also remember from the previous chapter, that we can describe molecules as being either polar or non-polar. A polar molecule is one in which there is a difference in electronegativity between the atoms in the molecule, such that the shared electron pair spends more time close to the atom that attracts it more strongly. The result is that one end of the molecule will have a slightly positive charge (δ+δ+), and the other end will have a slightly negative charge (δ+δ+). The molecule is said to be a dipole. However, it is important to remember that just because the bonds within a molecule are polar, the molecule itself may not necessarily be polar. The shape of the molecule may also affect its polarity. A few examples are shown in Table 1 to refresh your memory!

Table 1: Polarity in molecules with different atomic bonds and molecular shapes
Molecule Chemical formula Bond between atoms Shape of molecule Polarity of molecule
Hydrogen H22 Covalent
Figure 1
Figure 1 (CG11C2_001.png)
Non-polar
Hydrogen chloride HCl Polar covalent
Figure 2
Figure 2 (CG11C2_002.png)
Polar
Carbon tetrafluoromethane CF44 Polar covalent
Figure 3
Figure 3 (CG11C2_003.png)
Non-polar

Types of Intermolecular Forces

It is important to be able to recognise whether the molecules in a substance are polar or non-polar because this will determine what type of inermolecular forces there are. This is important in explaining the properties of the substance.

  1. Van der Waals forces These intermolecular forces are named after a Dutch physicist called Johannes van der Waals (1837 -1923), who recognised that there were weak attractive and repulsive forces between the molecules of a gas, and that these forces caused gases to deviate from 'ideal gas' behaviour. Van der Waals forces are weak intermolecular forces, and can be divided into three types:
    1. Dipole-dipole forces Figure 4 shows a simplified dipole molecule, with one end slightly positive and the other slightly negative.
      Figure 4: A simplified diagram of a dipole molecule
      Figure 4 (CG11C2_004.png)
      When one dipole molecule comes into contact with another dipole molecule, the positive pole of the one molecule will be attracted to the negative pole of the other, and the molecules will be held together in this way (Figure 5). Examples of materials/substances that are held together by dipole-dipole forces are HCl, FeS, KBr, SO22 and NO22.
      Figure 5: Two dipole molecules are held together by the attractive force between their oppositely charged poles
      Figure 5 (CG11C2_005.png)
    2. Ion-dipole forces As the name suggests, this type of intermolecular force exists between an ion and a dipole molecule. You will remember that an ion is a charged atom, and this will be attracted to one of the charged ends of the polar molecule. A positive ion will be attracted to the negative pole of the polar molecule, while a negative ion will be attracted to the positive pole of the polar molecule. This can be seen when sodium chloride (NaCl) dissolves in water. The positive sodium ion (Na++) will be attracted to the slightly negative oxygen atoms in the water molecule, while the negative chloride ion (Cl--) is attracted to the slightly positive hydrogen atom. These intermolecular forces weaken the ionic bonds between the sodium and chloride ions so that the sodium chloride dissolves in the water (Figure 6).
      Figure 6: Ion-dipole forces in a sodium chloride solution
      Figure 6 (CG11C2_006.png)
    3. London forces These intermolecular forces are also sometimes called 'dipole- induced dipole' or 'momentary dipole' forces. Not all molecules are polar, and yet we know that there are also intermolecular forces between non-polar molecules such as carbon dioxide. In non-polar molecules the electronic charge is evenly distributed but it is possible that at a particular moment in time, the electrons might not be evenly distributed. The molecule will have a temporary dipole. In other words, each end of the molecules has a slight charge, either positive or negative. When this happens, molecules that are next to each other attract each other very weakly. These London forces are found in the halogens (e.g. F22 and I22), the noble gases (e.g. Ne and Ar) and in other non-polar molecules such as carbon dioxide and methane.
  2. Hydrogen bonds As the name implies, this type of intermolecular bond involves a hydrogen atom. The hydrogen must be attached to another atom that is strongly electronegative, such as oxygen, nitrogen or fluorine. Water molecules for example, are held together by hydrogen bonds between the hydrogen atom of one molecule and the oxygen atom of another (Figure 7). Hydrogen bonds are stronger than van der Waals forces. It is important to note however, that both van der Waals forces and hydrogen bonds are weaker than the covalent and ionic bonds that exist between atoms.
    Figure 7: Two representations showing the hydrogen bonds between water molecules: space-filling model and structural formula.
    Figure 7 (CG11C2_007.png)

Types of intermolecular forces

  1. Complete the following table by placing a tick to show which type of intermolecular force occurs in each substance:
    Table 2
    FormulaDipole-dipoleMomentary dipoleIon-dipolehydrogen bond
    HCl    
    CO22    
    I22    
    H22O    
    KI(aq)    
    NH33    
  2. In which of the substances above are the intermolecular forces...
    1. strongest
    2. weakest

Understanding intermolecular forces

The types of intermolecular forces that occur in a substance will affect its properties, such as its phase, melting point and boiling point. You should remember, if you think back to the kinetic theory of matter, that the phase of a substance is determined by how strong the forces are between its particles. The weaker the forces, the more likely the substance is to exist as a gas because the particles are far apart. If the forces are very strong, the particles are held closely together in a solid structure. Remember also that the temperature of a material affects the energy of its particles. The more energy the particles have, the more likely they are to be able to overcome the forces that are holding them together. This can cause a change in phase.

Definition 2: Boiling point

The temperature at which a material will change from being a liquid to being a gas.

Definition 3: Melting point

The temperature at which a material will change from being a solid to being a liquid.

Now look at the data in Table 3.

Table 3: Melting point and boiling point of a number of chemical substances
Formula Formula mass Melting point (00C) Boiling point (00C) at 1 atm
He 4 -270 -269
Ne 20 -249 -246
Ar 40 -189 -186
F22 38 -220 -188
Cl22 71 -101 -35
Br22 160 -7 58
NH33 17 -78 -33
H22O 18 0 100
HF 20 -83 20

The melting point and boiling point of a substance, give us information about the phase of the substance at room temperature, and the strength of the intermolecular forces. The examples below will help to explain this.

Example 1: Fluorine (F22)

Phase at room temperature

Fluorine (F22) has a melting point of -22000C and a boiling point of -18800C. This means that for any temperature that is greater than -18800C, fluorine will be a gas. At temperatures below -22000C, fluorine would be a solid, and at any temperature inbetween these two, fluorine will be a liquid. So, at room temperature, fluorine exists as a gas.

Strength of intermolecular forces

What does this information tell us about the intermolecular forces in fluorine? In fluorine, these forces must be very weak for it to exist as a gas at room temperature. Only at temperatures below -18800C will the molecules have a low enough energy that they will come close enough to each other for forces of attraction to act between the molecules. The intermolecular forces in fluorine are very weak van der Waals forces because the molecules are non-polar.

Example 2: Hydrogen fluoride (HF)

Phase at room temperature

For temperatures below -8300C, hydrogen fluoride is a solid. Between -8300C and 2000C, it exists as a liquid, and if the temperature is increased above 2000C, it will become a gas.

Strength of intermolecular forces

What does this tell us about the intermolecular forces in hydrogen fluoride? The forces are stronger than those in fluorine, because more energy is needed for fluorine to change into the gaseous phase. In other words, more energy is needed for the intermolecular forces to be overcome so that the molecules can move further apart. Intermolecular forces will exist between the hydrogen atom of one molecule and the fluorine atom of another. These are hydrogen bonds, which are stronger than van der Waals forces.

What do you notice about water? Luckily for us, water behaves quite differently from the rest of the halides. Imagine if water were like ammonia (NH33), which is a gas above a temperature of -3300C! There would be no liquid water on the planet, and that would mean that no life would be able to survive here. The hydrogen bonds in water are particularly strong and this gives water unique qualities when compared to other molecules with hydrogen bonds. This will be discussed more in chapter (Reference). You should also note that the strength of the intermolecular forces increases with an increase in formula mass. This can be seen by the increasing melting and boiling points of substances as formula mass increases.

Applying your knowledge of intermolecular forces

Refer to the data in Table 3 and then use your knowledge of different types of intermolecular forces to explain the following statements:

  • The boiling point of F22 is much lower than the boiling point of NH33
  • At room temperature, many elements exist naturally as gases
  • The boiling point of HF is higher than the boiling point of Cl22
  • The boiling point of water is much higher than HF, even though they both contain hydrogen bonds

Intermolecular forces in liquids

Intermolecular forces affect a number of properties in liquids:

  • Surface tension You may have noticed how some insects are able to walk across a water surface, and how leaves float in water. This is because of surface tension. In water, each molecule is held to the surrounding molecules by strong hydrogen bonds. Molecules in the centre of the liquid are completely surrounded by other molecules, so these forces are exerted in all directions. However, molecules at the surface do not have any water molecules above them to pull them upwards. Because they are only pulled sideways and downwards, the water molecules at the surface are held more closely together. This is called surface tension.
    Figure 8: Surface tension in a liquid
    Figure 8 (CG11C2_008.png)
  • Evaporation In a liquid, each particle has kinetic energy, but some particles will have more energy than others. We therefore refer to the average kinetic energy of the molecules when we describe the liquid. When the liquid is heated, those particles which have the highest energy will be able to overcome the intermolecular forces holding them in the liquid phase, and will become a gas. This is called evaporation. Evaporation occurs when a liquid changes to a gas. The stronger the intermolecular forces in a liquid, the higher the temperature of the molecules will have to be for it to become a gas. You should note that a liquid doesn't necessarily have to reach boiling point before evaporation can occur. Evaporation takes place all the time. You will see this if you leave a glass of water outside in the sun. Slowly the water level will drop over a period of time. What happens then to the molecules of water that remain in the liquid? Remember that it was the molecules with the highest energy that left the liquid. This means that the average kinetic energy of the remaining molecules will decrease, and so will the temperature of the liquid. A similar process takes place when a person sweats during exercise. When you exercise, your body temperature increases and you begin to release moisture (sweat) through the pores in your skin. The sweat quickly evaporates and causes the temperature of your skin to drop. This helps to keep your body temperature at a level that is suitable for it to function properly.
  • Capillarity

Investigation : Capillarity

Half fill a beaker with water and hold a hollow glass tube in the centre as shown below. Mark the level of the water in the glass tube, and look carefully at the shape of the air-water interface in the tube. What do you notice?

Figure 9
Figure 9 (CG11C2_009.png)

At the air-water interface, you will notice a meniscus, where the water appears to dip in the centre. In the glass tube, the attractive forces between the glass and the water are stronger than the intermolecular forces between the water molecules. This causes the water to be held more closely to the glass, and a meniscus forms. The forces between the glass and the water also mean that the water can be 'pulled up' higher when it is in the tube than when it is in the beaker. Capillarity is the surface tension that occurs in liquids that are inside tubes.

Note: Interesting Fact :

Transpiration in plants - Did you know that plants also 'sweat'? In plants, this is called transpiration, and a plant will lose water through spaces in the leaf surface called stomata. Although this water loss is important in the survival of a plant, if a plant loses too much water, it will die. Plants that live in very hot, dry places such as deserts, must be specially adapted to reduce the amount of water that transpires (evaporates) from their leaf surface. Desert plants have some amazing adaptations to deal with this problem! Some have hairs on their leaves, which reflect sunlight so that the temperature is not as high as it would be, while others have a thin waxy layer covering their leaves, which reduces water loss. Some plants are even able to close their stomata during the day when temperatures (and therefore transpiration) are highest.

Tip:

In the same way that intermolecular forces affect the properties of liquids, they also affect the properties of solids. For example, the stronger the intermolecular forces between the particles that make up the solid, the harder the solid is likely to be, and the higher its melting point is likely to be.

Figure 10

Summary

  • Intermolecular forces are the forces that act between stable molecules.
  • The type of intermolecular force in a substance, will depend on the nature of the molecules.
  • Polar molecules have an unequal distribution of charge, meaning that one part of the molecule is slightly positive and the other part is slightly negative. Non-polar molecules have an equal distribution of charge.
  • There are three types of Van der Waal's forces. These are dipole-dipole, ion-dipole and London forces (momentary dipole).
  • Dipole-dipole forces exist between two polar molecules, for example between two molecules of hydrogen chloride.
  • Ion-dipole forces exist between ions and dipole molecules. The ion is attracted to the part of the molecule that has an opposite charge to its own. An example of this is when an ionic solid such as sodium chloride dissolves in water.
  • Momentary dipole forces occur between two non-polar molecules, where at some point there is an uequal distribution of charge in the molecule. For example, there are London forces between two molecules of carbon dioxide.
  • Hydrogen bonds occur between hydrogen atoms and other atoms that have a high electronegativity such as oxygen, nitrogen and fluorine. The hydrogen atom in one molecule will be attracted to the nitrogen atom in another molecule, for example. There are hydrogen bonds between water molecules and between ammonia molecules.
  • Intermolecular forces affect the properties of substances. For example, the stronger the intermolecular forces, the higher the melting point of that substance, and the more likely that substance is to exist as a solid or liquid. Its boiling point will also be higher.
  • In liquids, properties such as surface tension, capillarity and evaporation are the result of intermolecular forces.

Summary Exercise

  1. Give one word or term for each of the following descriptions:
    1. The tendency of an atom in a molecule to attract bonding electrons.
    2. A molecule that has an unequal distribution of charge.
    3. A charged atom.
  2. For each of the following questions, choose the one correct answer from the list provided.
    1. The following table gives the melting points of various hydrides:
      Table 4
      HydrideMelting point (00C)
      HI-50
      NH33-78
      H22S-83
      CH44-184
      In which of these hydrides does hydrogen bonding occur?
      1. HI only
      2. NH33 only
      3. HI and NH33 only
      4. HI, NH33 and H22S
      (IEB Paper 2, 2003)
    2. Refer to the list of substances below: HCl, Cl22, H22O, NH33, N22, HF Select the true statement from the list below:
      1. NH33 is a non-polar molecule
      2. The melting point of NH33 will be higher than for Cl22
      3. Ion-dipole forces exist between molecules of HF
      4. At room temperature N22 is usually a liquid
  3. The respective boiling points for four chemical substances are given below: Hydrogen sulphide -6000C Ammonia -3300C Hydrogen fluoride 2000C Water 10000C
    1. Which one of the substances exhibits the strongest forces of attraction between its molecules in the liquid state?
    2. Give the name of the force responsible for the relatively high boiling points of ammonia and water and explain how this force originates.
    3. The shapes of the molecules of hydrogen sulfide and water are similar, yet their boiling points differ. Explain.
    (IEB Paper 2, 2002)

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