INTERMOLECULAR FORCES

All around us we see matter in different phases. The air we breathe is a gas, while the water you drink is a liquid and the chair you are sitting on is a solid. In this chapter we are going to look at one of the reasons that matter exists as solids and liquids.

In the previous chapter, we discussed the different forces that exist between atoms (interatomic 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.

Intermolecular forces allow us to determine which substances are likely to dissolve in which other substances and what the melting and boiling points of substances are. Without intermolecular forces holding molecules together we would not exist.

Note that we will use the term molecule throughout this chapter as the compounds we are looking at are all covalently bonded and do not exist as giant networks (recall from grade 10$\text{10}$ that there are three types of bonding: metallic, ionic and covalent). Sometimes you will see the term simple molecule. This is a covalent molecular structure.

Interatomic (between atoms) forces are also known as intramolecular (within molecules) forces. You can remember this by thinking of international which means between nations.

4.1 Intermolecular and interatomic forces (ESBMM)

Intermolecular forces

Intermolecular forces are forces that act between molecules.

You will also recall 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+${\delta }^{+}$), and the other end will have a slightly negative charge (δ−${\delta }^{-}$). The molecule is said to be a dipole.

A dipole molecule is a molecule that has two (di) poles. One end of the molecule is slightly positive and the other is slightly negative. We can depict this very simply as an oval with one positive side and one negative. In reality however, the molecules do not look like this, they look more like the images in Figure 4.1.

Figure 4.1: A different representation of dipole molecules. The red region is slightly negative, and the blue region is slightly positive.

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 4.1 to refresh your memory. Note that we have shown tetrahedral molecules with all the terminal atoms at 90$\text{90}$°$\text{°}$ to each other (i.e. flat or 2-dimensional), but the shape is really 3-dimensional.

Molecule	Chemical formula	Bond between atoms	Shape of molecule	Polarity of molecule
Hydrogen	H2${\text{H}}_2$	Non-polar covalent		Non-polar
Hydrogen chloride	HCl$\text{HCl}$	Polar covalent		Polar
Carbon tetrafluoride	CF4${\text{CF}}_4$	Polar covalent		Non-polar
Trifluoro-methane	CHF3${\text{CHF}}_3$	Polar covalent		Polar

Table 4.1: Polarity in molecules with different atomic bonds and molecular shapes.

Types of intermolecular forces (ESBMN)

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 intermolecular forces there are. This is important in explaining the properties of the substance.

1.  

   Ion-dipole forces

    

   As the name suggests, this type of intermolecular force exists between an ion and a dipole (polar) 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$\text{NaCl}$) dissolves in water. The positive sodium ion (Na+${\text{Na}}^{+}$) will be attracted to the slightly negative oxygen atoms in the water molecule, while the negative chloride ion (Cl−${\text{Cl}}^{-}$) is attracted to the slightly positive hydrogen atoms. These intermolecular forces weaken the ionic bonds between the sodium and chloride ions so that the sodium chloride dissolves in the water (Figure 4.2).

   

   Figure 4.2: Ion-dipole forces in a sodium chloride solution.

   This is a simplified diagram to highlight the regions of positive and negative charge. When sodium chloride dissolves in water it can more accurately be shown as:

   

2.  

   Ion-induced-dipole forces

    

   Similar to ion-dipole forces these forces exist between ions and non-polar molecules. The ion induces a dipole in the non-polar molecule leading to a weak, short lived force which holds the compounds together.

    

   These forces are found in haemoglobin (the molecule that carries oxygen around your body). Haemoglobin has Fe2+${\text{Fe}}^{2+}$ ions. Oxygen (O2${\text{O}}_2$) is attracted to these ions by ion-induced dipole forces.

3.  

   Dipole-dipole forces

    

   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 4.3). Examples of materials/substances that are held together by dipole-dipole forces are HCl$\text{HCl}$, SO2${\text{SO}}_2$ and CH3Cl${\text{CH}}_3\text{Cl}$.

   

   Figure 4.3: Two dipole molecules are held together by the attractive force between their oppositely charged poles.

   One special case of this is hydrogen bonding.

4.  

   Induced dipole forces

   These intermolecular forces are also sometimes called “London forces” or “momentary dipole” forces or “dispersion” forces.

   We know that while carbon dioxide is a non-polar molecule, we can still freeze it (and we can also freeze all other non-polar substances). This tells us that there must be some kind of attractive force in these kinds of molecules (molecules can only be solids or liquids if there are attractive forces pulling them together). This force is known as an induced dipole force.

    

   In non-polar molecules the electronic charge is usually evenly distributed but it is possible that at a particular moment in time, the electrons might not be evenly distributed (remember that the electrons are always moving in their orbitals). 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 forces are found in the halogens (e.g. F2${\text{F}}_2$ and I2${\text{I}}_2$) and in other non-polar molecules such as carbon dioxide and carbon tetrachloride.

    

   All covalent molecules have induced dipole forces. For non-polar covalent molecules these forces are the only intermolecular forces. For polar covalent molecules, dipole-dipole forces are found in addition to the induced dipole forces.

   When the noble gases condense, the intermolecular forces that hold the liquid together are induced dipole forces.

5.  

   Dipole-induced-dipole forces

    

   This type of force occurs when a molecule with a dipole induces a dipole in a non-polar molecule. It is similar to an ion-induced dipole force. An example of this type of force is chloroform (CHCl3${\text{CHCl}}_3$) in carbon tetrachloride (CCl4${\text{CCl}}_4$).

The following image shows the types of intermolecular forces and the kinds of compounds that lead to those forces.

Figure 4.4: The types of intermolecular forces. The boxes represent the type of compound while the lines represent the type of force.

The last three forces (dipole-dipole forces, dipole-induced dipole forces and induced dipole forces) are sometimes collectively known as van der Waals' forces. We will now look at a special case of dipole-dipole forces in more detail.

Hydrogen bonds

As the name implies, this type of intermolecular bond involves a hydrogen atom. When a molecule contains a hydrogen atom covalently bonded to a highly electronegative atom (O$\text{O}$, N$\text{N}$ or F$\text{F}$) this type of intermolecular force can occur. The highly electronegative atom on one molecule attracts the hydrogen atom on a nearby molecule.

Water molecules for example, are held together by hydrogen bonds between the hydrogen atom of one molecule and the oxygen atom of another (fig:hydrogen bonds). Hydrogen bonds are a relatively strong intermolecular force and are stronger than other dipole-dipole forces. It is important to note however, that hydrogen bonds are weaker than the covalent and ionic bonds that exist between atoms.

Do not confuse hydrogen bonds with actual chemical bonds. Hydrogen bonding is an example of a case where a scientist named something believing it to be one thing when in fact it was another. In this case the strength of the hydrogen bonds misled scientists into thinking this was actually a chemical bond, when it is really just an intermolecular force.

Figure 4.5: Two representations showing the hydrogen bonds between water molecules: space-filling model and structural formula.

The difference between intermolecular and interatomic forces (ESBMP)

It is important to realise that there is a difference between the types of interactions that occur in molecules and the types that occur between molecules. In the previous chapter we focused on the interactions between atoms. These are known as interatomic forces or chemical bonds. We also studied covalent molecules in more detail.

Remember that a covalent bond has an electronegativity difference of less than 2,1$\text{2,1}$. Covalent molecules have covalent bonds between their atoms. Van der Waals' forces only occur in covalent molecules. We can show the interatomic and intermolecular forces between covalent compounds diagrammatically or in words. Intermolecular forces occur between molecules and do not involve individual atoms. Interatomic forces are the forces that hold the the atoms in molecules together. Figure 4.5 shows this.

	Interatomic forces	Intermolecular forces
Atoms or molecules	Forces between atoms	Forces between molecules
Strength of forces	Strong forces	Relatively weak forces
Distance between atoms or molecules	Very short distances	Larger distances than bonds