- Online Text
- 1. Introduction
- 2. Valence Electron Patterns and Lewis Structures
- 3. Ionic Bonds
- 4. Covalent Bonds and the Octet Rule
- 5. Polarity and Basic Lewis Structures
- 6. Advanced Lewis Structures
- 7. VSEPR Theory
- 8. Hybrid Orbitals
- 9. Intermolecular Forces
- 10. Physical Properties of Molecules
- 11. Conclusion
- 12. Further Reading
- Unit Guide (PDF)
Section 9: Intermolecular Forces
Ionic and covalent bonds hold atoms together in molecules. There are also forces of attraction that exist between molecules themselves. These intermolecular forces, while much weaker than ionic and covalent bonds, have important effects on the way molecules, once formed, interact with one another. We'll look at three of these forces, listed here from strongest to weakest:
- Hydrogen bonding
- Dipole-dipole forces
- London dispersion forces
The strongest form of intermolecular force is called "hydrogen bonding." There are two things that are required for a hydrogen bond to form. First, there must be a hydrogen atom that is attached to one of the most electronegative atoms: nitrogen, oxygen, or fluorine. This makes the hydrogen very electron deficient, as the electronegative atom is pulling electron density away from the hydrogen. The second thing that is required is a lone pair on another very electronegative atom, a fluorine, oxygen, or nitrogen. The electron deficient hydrogen shares electrons weakly with this lone pair, almost like the sharing in a covalent bond. The difference is that a hydrogen bond is about a hundred times weaker than a normal covalent bond. Because organic molecules have so many molecules with nitrogens and oxygens in them, hydrogen bonding is a common motif in biochemistry.
FIgure 5-18. The Structure of Ice
This is a representation of the water molecules in an ice crystal. In the picture, red balls are oxygen atoms and white balls are hydrogen atoms. Hydrogen bonds are represented by the dotted lines. Note how each water molecule has both hydrogen atoms making hydrogen bonds, and both lone pairs on the oxygen are hydrogen bonding to nearby waters.
© Wikimedia Commons, Public Domain.
Water is a very special compound in terms of hydrogen bonding. It has two hydrogens on an oxygen atom and two lone pairs on an oxygen atom. This means that, on average, every water molecule in a glass of water can make two hydrogen bonds. So, even though water is a three-atom molecule with a molecular weight of only 18 atomic mass units, it is a liquid that boils at the high temperature of 100°C. All other molecules of that size and weight are gases, because they don't have those strong additional intermolecular forces. This affects many of the properties of water. This is why there is surface tension that keeps water just over the rim of a glass or capillary action that draws water up thin tubes, like in the xylem tubes in plants that bring water up the stems. Water molecules work very hard to make sure they can make as many hydrogen bonds as possible, and this becomes very apparent when water begins to freeze into ice. In the solid state, in order to maximize each water molecule making its hydrogen bonds, it adopts the structure seen in Figure 5-18. This structure is very open, and has large hexagonal spaces in it. This open hexagonal pattern caused by the hydrogen bonding explains why snowflakes are six-sided and why ice is less dense than liquid water and it floats.
While many of water's properties are directly related to its ability to make hydrogen bonds, any molecule with hydrogens bound to nitrogen or oxygen, with lone pairs on a nitrogen or oxygen atom, can make hydrogen bonds. Proteins in the body are often held to their shape by the hydrogen bonding between the different amino acids. But, on an even grander scale, the two strands of DNA are held together by hydrogen bonds (Figure 5-19). While a hydrogen bond isn't very strong, with thousands of base pairs in DNA—each making two or three hydrogen bonds to each other—those intermolecular forces add up quickly, and therefore the DNA double helix is very stable.
Figure 5-19. Hydrogen Bonding in DNA
In the double helix of DNA, there are strands of molecules called "nucleoside phosphates" that are covalently bound together. However, each strand is held to the other strand by only weak hydrogen bonds. This diagram shows how the guanine and cytosine bases are perfectly designed by nature to hydrogen bond to each other. The hydrogen bonds are represented by the yellow dotted lines. Note how the molecules perfectly placed so that each hydrogen on an N or O lines up with a lone pair on another N or O.
© Science Media Group.
Earlier in this unit, we saw that molecules with an unsymmetrical cloud of electrons are called "polar molecules." A molecule that is polar can also be called a "dipole." The more uneven the cloud of electrons is, the more polar a molecule is, and the larger its dipole is. All of the molecules that can hydrogen bond are polar; this is because of all the lone pairs, such as the ones on water that make one side of the water molecule much more negative than the other. However, not all molecules that are polar can hydrogen bond.
The next strongest intermolecular force is that which occurs between polar molecules, called "dipole-dipole interactions" or "dipole forces." This type of force is also generally called a "van der Waals force." Van der Waals forces is a collective term for all non-covalent attractions between molecules, which will also include the London forces discussed at the end of this section. The strength of these dipole forces is related to how polar the molecules are or how big their dipoles are. Essentially, molecules are like little miniature magnets with a more negative end (usually with more lone pairs) and a more positive end. These molecules line themselves up so that their positive ends are near the negative ends of the next molecules. Then, a very weak coulombic force attracts them together. These are the types of intermolecular forces that hold molecules like chloroform (CHCl3) and hydrogen chloride (HCl) to each other.
London Dispersion Forces
London dispersion forces are the weakest intermolecular force. These forces are also under the category of van der Waals forces and are sometimes called "London forces" or "dispersion forces." These are the only type of forces available to nonpolar molecules, which have perfectly symmetrical magnetic clouds. Since these clouds aren't imbalanced, they don't have a very negative end and a very positive end like the polar molecules, which experience the dipole forces. However, a big cloud of electrons, just due to probability, is likely to have more electrons on one side compared to the other at any particular point in time. This creates a very small or temporary dipole in the molecule, which can then cause a nearby molecule to shift its cloud of electrons to balance it, making an induced dipole next to it. This pair of induced dipoles are very weakly attracted together, and they cause the molecules to stick together. These forces of attraction are called London forces after the physicist Fritz London, who first characterized them in the 1930s. For non-polar molecules, like carbon dioxide (CO2) in the form of dry ice or the xenon molecules in liquid xenon, the forces holding these atoms and molecules together are the London forces.
London forces are extremely weak, but if a molecule is large, they can add up to a noticeable effect. For example, the stronger the force of attraction between molecules, the more energy has to be used to separate them. When one heats a liquid, at a certain temperature the molecules separate and become a gas. We call the temperature at which this happens the boiling point; molecules with stronger intermolecular forces have higher boiling points. Larger molecules tend to have higher boiling points than smaller molecules because the London forces between them make the molecules stick together more. Table 5-1 shows a group of non-polar molecules of increasing size. As molecular size increases, the boiling points increase as a result of increasing London forces.
|Molecule (from smallest to largest)||Boiling Point|
In summary, while only certain molecules that have the proper sets of atoms present can hydrogen bond, ALL molecules have London dispersion forces. Also, all polar molecules, whether they can hydrogen bond or not, have dipole-dipole forces. In most molecular compounds, their liquid and solid forms are held together by combinations of all three of these intermolecular forces.