Although the outer portion of the atmosphere, beyond the stratosphere, contains only a small fraction of the atmospheric mass, it plays an important role in determining the conditions of life at Earth's surface. This upper layer forms the outer defense against the hail of radiation and high-energy particles that continually bombard the planet. As this occurs, the molecules and atoms of the upper atmosphere undergo chemical changes.
The Sun emits radiant energy over a wide range of wavelengths. The shorter-wavelength, higher-energy radiations in the ultraviolet range of the spectrum are sufficiently energetic to cause chemical changes. Recall that electromagnetic radiation can be pictured as a stream of photons. The energy of each photon is given by the relationship E = h , where h is Planck's constant and is the frequency of the radiation. For a chemical change to occur when radiation falls on Earth's atmosphere, two conditions must be met. First, there must be photons with energy sufficient to accomplish whatever chemical process is being considered. Second, molecules must absorb these photons. When these requirements are met, the energy of the photons is converted into some other form of energy within the molecule.
The rupture of a chemical bond resulting from absorption of a photon by a molecule is called photodissociation. Photodissociation does not form ions. The bond cleavage leaves half the bonding electrons with each of the two atoms forming two neutral particles.
One of the most important processes occurring in the upper atmosphere above about 120 km elevation is the photodissociation of the oxygen molecule:
The minimum energy required to cause this change is determined by the dissociation energy of O2, 495 kJ/mol. In Sample Exercise 18.2 we calculate the longest wavelength photon having sufficient energy to dissociate the O2 molecule.
What is the maximum wavelength of light (in nanometers) that has enough energy per photon to dissociate the O2 molecule?
SOLUTION The dissociation energy of O2 is 495 kJ/mol. Using this value, we can calculate the amount of energy needed to break the bond in a single O2 molecule:
We next use the Planck relation, E = h , to calculate the frequency, , of a photon that has this amount of energy:
Finally, we use the relationship between the frequency and wavelength of light (Section 6.1) to calculate the wavelength of the light:
Thus, ultraviolet light of wavelength 242 nm has sufficient energy per photon to photodissociate an O2 molecule. Because photon energy increases as wavelength decreases, any photon of wavelength shorter than 242 nm will have sufficient energy to dissociate O2.
The bond energy in N2 is 941 kJ/mol (Table 8.4). What is the longest wavelength photon that has sufficient energy to dissociate N2? Answer: 127 nm
The second condition that must be met before dissociation actually occurs is that the photon must be absorbed by O2. Fortunately for us, O2 absorbs much of the high-energy, short-wavelength radiation from the solar spectrum before it reaches the lower atmosphere. As it does, atomic oxygen, O, is formed. At higher elevations the dissociation of O2 is very extensive. At 400 km, only 1 percent of the oxygen is in the form of O2; the other 99 percent is in the form of atomic oxygen. At 130 km O2 and O are just about equally abundant. Below this elevation O2 is more abundant than O.
The bond-dissociation energy of N2 is very high (Table 8.4). As shown in Practice Exercise 18.2, only photons of very short wavelength possess sufficient energy to cause dissociation of this molecule. Furthermore, N2 does not readily absorb photons, even when they do possess sufficient energy. The overall result is that very little atomic nitrogen is formed in the upper atmosphere by dissociation of N2.
In 1901 Guglielmo Marconi performed a sensational experiment. He received a radio signal in St. John's, Newfoundland, that had been transmitted from Land's End, England, some 2900 km away. Because radio waves were thought to travel in straight lines, it had been assumed that radio communication over large distances on Earth would be impossible. Marconi's successful experiment suggested that Earth's atmosphere in some way substantially affects radio-wave propagation. His discovery led to intensive study of the upper atmosphere. In about 1924 the existence of electrons in the upper atmosphere was established by experimental studies.
For each electron present in the upper atmosphere, there must be a corresponding positively charged ion. The electrons in the upper atmosphere result mainly from the photoionization of molecules, caused by solar radiation. In the photoionization process a molecule absorbs radiation, and that absorbed energy causes the loss of an electron. Photoionization requires that a photon be absorbed by the molecule and also that this photon have enough energy to remove an electron.
Some of the more important ionization processes occurring in the upper atmosphere above about 90 km are shown in Table 18.2, together with the ionization energies and , the maximum wavelength of a photon capable of causing ionization. Photons with energies sufficient to cause ionization have wavelengths in the high-energy region of the ultraviolet. These wavelengths are completely filtered out of the radiation reaching Earth as a result of their absorption by the upper atmosphere.