intermolecular forces review
Intermolecular forces – review The states of matter adopted by chemical substances and the physical behaviour of the various states of matter depend ultimately upon the strength and nature of forces between atoms and molecules in the substance. Melting points and boiling points are key measures of the strength of intermolecular forces. The stronger the intermolecular interactions in a substance between itself and between others substances, the more energy is required to cause the substance to turn into a liquid or a gas. For instance, iodine in its standard state is a solid. At the molecular scale, its molecules are tightly packed, interacting very strongly with one another. The iodine has to be heated to a higher temperature to cause it to turn into a liquid and then a gas. Note that iodine is located near the bottom of the periodic table; I2 is a large molecule. In contrast, bromine is liquid under standard temperature and pressure. At the molecular level, the molecules are not fixed in place or rigid. Though they are still closely packed, they are able to freely rotate and slide past one another fairly rapidly. They spread out to the shape of their container for this reason, but they do not go into the gas phase and thus do not fill up the container entirely. Note that bromine is one period above iodine in the periodic table. One period further up lies chlorine, which under normal conditions is a gas. Its molecules are far apart from one another and fly around at really high speeds (thousands of kilometres an hour), not interacting very much with one another. The intermolecular forces in gases are very weak, but they do have an impact on the ideal gas law. The reason for some substances being gases, others liquids and others solids is that a substance exists in a condensed phase (i.e. a solid or a liquid) when its molecules have too little average kinetic energy to overcome intermolecular forces of attraction. As can be seen from the graph, the intermolecular attractive energy of the above molecules is I2>Br2>Cl2. Note that for iodine and bromine, the average kinetic energy at room temperature (that is, the average kinetic energy of molecules in the gas phase) is lower than their intermolecular attractive energy. This energy is not overcome and thus they remain in their condensed phase. In contrast, chlorine and fluorine both have an intermolecular attractive energy which is significantly lower than the average kinetic energy; thus at room temperature they are gases. The larger a molecule gets, the lower its velocity in the gas phase for a fixed kinetic energy.
Types of intermolecular forces The trends we have seen above are all the result of forces between the molecules. Molecules, for the purpose of intermolecular forces, are defined as small groups of atoms with strong bonds between atoms in the molecule, but weak bonding at the surfaces between molecules. There are three basic types of intermolecular forces (also known as van der Waals forces): • Dipole-dipole forces • These occur when a permanent dipole moment on a molecule leads to a partial negative charge on one part of it and a partial positive charge on the other. • The partially negatively charged atom is attracted to the partially positively charged atom of another molecule. • Dipole moments will align and this causes forces to form between molecules. • The stronger dipole-dipole interactions are between molecules of a substance, the more energy is required to turn that substance into a liquid or a gas. • Dispersion forces (also known as London forces) • Hydrogen bonding Dipole-dipole forces Two molecules with permanent dipole moments will attract each other by the approach of their oppositely charged ends. Consider acetone. It contains three carbon atoms, here represented by black spheres. The central carbon has three regions of electron density, which makes it a trigonal planar structure with sp2 hybridisation. A remaining unhybridised p orbital forms a π bond with the unhybridised p orbital of the oxygen (red sphere), causing a double bond to form. The oxygen is a very electronegative atom, so it sucks electron density away from the carbon and hydrogen atoms and towards itself. This causes a negative partial charge (∂–) to form on one side of the molecule, whereas the carbons and hydrogens attain a partial positive charge (∂+). The dipole moments align with one another and the interactions relatively are strong. This is why acetone is a liquid under standard conditions rather than a gas. However, the interactions are not as strong as in water, which is why acetone has an odour (i.e. some molecules escape into the gas phase). Consider also carbon tetrafluoride (CF4). It has no permanent dipole and a boiling point of -125˚C. The boiling point of CHF3, on the other hand, is -84.4˚C. This is because it does have a permanent dipole moment.
Types of intermolecular forces The trends we have seen above are all the result of forces between the molecules. Molecules, for the purpose of intermolecular forces, are defined as small groups of atoms with strong bonds between atoms in the molecule, but weak bonding at the surfaces between molecules. There are three basic types of intermolecular forces (also known as van der Waals forces): • Dipole-dipole forces • These occur when a permanent dipole moment on a molecule leads to a partial negative charge on one part of it and a partial positive charge on the other. • The partially negatively charged atom is attracted to the partially positively charged atom of another molecule. • Dipole moments will align and this causes forces to form between molecules. • The stronger dipole-dipole interactions are between molecules of a substance, the more energy is required to turn that substance into a liquid or a gas. • Dispersion forces (also known as London forces) • Hydrogen bonding Dipole-dipole forces Two molecules with permanent dipole moments will attract each other by the approach of their oppositely charged ends. Consider acetone. It contains three carbon atoms, here represented by black spheres. The central carbon has three regions of electron density, which makes it a trigonal planar structure with sp2 hybridisation. A remaining unhybridised p orbital forms a π bond with the unhybridised p orbital of the oxygen (red sphere), causing a double bond to form. The oxygen is a very electronegative atom, so it sucks electron density away from the carbon and hydrogen atoms and towards itself. This causes a negative partial charge (∂–) to form on one side of the molecule, whereas the carbons and hydrogens attain a partial positive charge (∂+). The dipole moments align with one another and the interactions relatively are strong. This is why acetone is a liquid under standard conditions rather than a gas. However, the interactions are not as strong as in water, which is why acetone has an odour (i.e. some molecules escape into the gas phase). Consider also carbon tetrafluoride (CF4). It has no permanent dipole and a boiling point of -125˚C. The boiling point of CHF3, on the other hand, is -84.4˚C. This is because it does have a permanent dipole moment.
