Absolutely Small - Michael D. Fayer [137]
We solved the quantum particle in a box problem and found that the energy levels depend on the size of the box. A larger box (larger molecule) has energy levels that are more closely spaced than a smaller box. The result, which applies to real molecules as well as the particle in a box, is that large molecules tend to absorb light in the red part of the spectrum. Red light is low energy light, and for large molecules there are relatively small energy differences between the energy levels. Smaller molecules absorb light in the blue because the molecular energy level differences are larger, and blue light is higher energy than red light. Really small molecules, like benzene (see Chapter 18), absorb in the ultraviolet part of the spectrum. Thus, they do not absorb visible light. Crystals of small molecules like naphthalene (mothballs) look white because none of the visible wavelengths can be absorbed. The energy levels are too widely spaced to absorb visible light. All of the visible light bounces off of the crystals, and they look white. This is the reason that salt crystals in a salt shaker are white. It is also the reason sugar crystals are white. Both salt and sugar have widely spaced energy levels that absorb in the ultraviolet, and reflect all visible colors of light.
QUANTUM PHENOMENA HOLD ATOMS TOGETHER TO MAKE MOLECULES AND DETERMINE THEIR SHAPES
We now know what holds atoms together to make molecules, what gives molecules their shapes, and why details of the shapes are important. We have seen that the electron waves of atoms combine to make molecular orbitals. Sharing the electrons among atoms in molecular orbitals can result in chemical bonds that hold atoms together to form molecules. In Chapters 12 through 14, we looked at molecular orbitals in some detail. We learned that molecular orbitals come in two flavors, bonding and antibonding. Appropriately placing electrons in the simple molecular orbital energy level diagrams can provide a great deal of information. In the hydrogen molecule (Chapter 12), the two electrons from the two hydrogen atoms go into the lowest energy molecular orbital, the bonding MO. The result is a covalent bond in which the atoms share a pair of electrons. However, the same considerations enabled us to see why a diatomic helium molecule does not exist. Each helium atom contributes two electrons to the hypothetical helium diatomic molecule. The first two go into the bonding MO but because of the Pauli Exclusion Principle, the next two electrons must go into the antibonding MO. The result is no net bonding, and the He2 molecule does not exist. The covalent chemical bond is intrinsically a quantum phenomenon with no classical mechanics explanation.
For atoms larger than hydrogen, combining different s and p atomic orbitals produced hybrid orbitals with various shapes. The combination of different-shaped hybrid atomic orbitals to form molecular orbitals is responsible for the types of bonds that are formed (single, double, triple) and the shapes of molecules. We paid particular attention to organic molecules, which are molecules formed mainly from carbon, hydrogen, oxygen, and several other atoms. Organic molecules are important because they are the basis for life, as well as for materials such as plastics. We found that the types of bonds are very important. A molecule can readily rotate around a carbon-carbon single bond, changing the molecule’s shape, but a molecule cannot rotate around a carbon-carbon double bond. The inability of organic molecules