Experiment 2 Intermolecular Forces There are three general types of intermolecular forces. All substances exhibit London Dispersion Forces (LDF), and they are generally the weakest of the three types. These London forces are due to the attractions between small, temporary dipoles that arise from the constant, random movement of the electrons in a substance. As molar mass increases, the size of the electron cloud increases as well. It becomes more easily distorted, and produces temporary dipoles of greater magnitude. This causes the attractions to be stronger, requiring more energy for both fusion and vaporization. For halogens, this results in increasing melting and boiling points, shown by the fact that at room temperature F2 and Cl2 are gaseous, Brz is liquid and I is solid. The extent to which the electron cloud can be distorted is called polarizability. Dipole-dipole forces exist between molecules that are polar. Since the dipoles are permanent, these attractions are generally stronger than London Dispersion Forces. This means that a polar molecule with similar molar mass as a nonpolar molecule will have higher melting points and boiling points. Not all molecules containing polar bonds are polar. The polar bonds must be unevenly dispersed in the molecule in order to produce a polar molecule. CO2 and CBr4, for example, have polar bonds but are not polar molecules. dipole attraction that is stronger than other dipole-dipole attractions. Hydrogen bonds form when a hydrogen atom is covalently bonded to a very electronegative atom. This causes its electron to be drawn away from its nucleus. The positive hydrogen is then attracted to the very electronegative atom in a neighboring molecule. In order to observe hydrogen bonding, the hydrogen atom must be covalently bonded to fluorine, oxygen or nitrogen. A hydrogen atom bonded to a carbon atom cannot create a hydrogen bond. It's important to note that, despite its name, a hydrogen bond is an intermolecular force, not a bond. The figure to the right illustrates H-bonding between water molecules. H-bonding is important in biochemistry; the structure of a biopolymer is largely determined by the formation of hydrogen bonds. H H O H- The relative strengths of the three types of intermolecular forces, and thus boiling points, are generally as follows: London Dispersion Forces < Dipole-Dipole Forces < H-Bonding However, this is not always true. Since molar mass is also a factor, a large non-polar molecule can have a higher boiling point than a compound that interacts with dipole-diploe forces, or even a substance with H-bonding. For example, octane, a component of gasoline, has a boiling point of 125C- much higher than acetone (dipole-dipole) and H20 (H-bonding). This is due to the polarizability of the large electron cloud. To make comparisons of the intermolecular forces of a substance, evaporation rate can be used instead of boiling point. Evaporation rate is the ratio of the change in temperature to the change in time as a substance evaporates. A faster rate of evaporation translates to a lower boiling point and, in turn, weaker intermolecular forces. Bergen Community College 15 General Chemistry II Laboratory