CHM 1020--Chemistry for Liberal Studies--Spring 1999

Chemistry 1020--Lecture 6--Notes

Wavelength x frequency = velocity, or  ln=c

C is the velocity of light, which in a vacuum is 3.00 x 108 ms-1.

Visible light spans from l of 400 nm to 700 nm.

Spectrum of rainbow: ROY G. BIV (700nm® 400nm)

Slightly longer wavelengths are the infrared; slightly shorter is the ultraviolet.

Known spectrum extends from l of 10-14 m to 104 m as follows:

Cosmic rays—gamma rays—X-rays—UV—visible—infrared—microwaves—TV—radio. (Learn relative order, not necessarily range of l for each.)

Note Figure 2.3 in the book, showing the spectrum of radiation from the sun impinging on the upper atmosphere.

Frequencies vary from 1022 Hertz to 104 Hertz. (AM radio frequencies are in kilohertz; FM frequencies are in megahertz)

Go to this web site to explore the electromagnetic spectrum--showing the properties of radiation at various parts of the spectrum.  The site also has a calculator for interconverting wavelength, frequency, and energy (see later in the lecture).

Then do some conversions:

Wavelength of WFSQ signal (88.9 MHz)

Frequency of blue light at 680 nm.

Quantum theory: sort of an atomic theory for energy. Idea that energy comes in definite packages called quanta, and that atoms exist only in discrete energy states. A quantum of electromagnetic radiation is considered to be a particle called a photon.

Particle-wave duality seems to our senses to be a contradiction in terms. Yet not only does electromagnetic radiation seem to behave that way, matter at the very tiny level of electrons and below also behaves that way.

Through a combination of experiments by Max Planck and Albert Einstein, we can put a quantitative figure on this particle, and relate the particle and energy properties of light. The basic relationship is:

E = hn (or the counterpart E = hc/l ).

h = Planck’s constant, which has the value 6.63 x 10-34 Js

The higher the frequency, or shorter the wavelength, the greater the energy.

### Calculate some energies of various forms of radiation:

Microwave (l = 0.95 cm)

Infrared (l = 10 m m) E = 1.989x10-20 J = 2.0x10-20 J
Visible (l = 500 nm) E = 3.978x10-19 J = 3.98x10-19 J
UV ( l = 200 nm) E = 9.945x10-19 J = 9.95x10-19 J (~600 kJ/mol)
X-ray (l = 0.1 nm) E = 1.989x10-15J = 2x10-15 J (~1x106 kJ/mol)

The answers come out in Joules/photon, and the numbers don’t seem very big. But when you consider the energy required to do something to a molecule, we see that microwave and infrared radiation only have enough energy to enhance the motions (vibrations) of molecules. Infrared radiation is radiant heat, and associated with motions in molecules. When one gets to ultraviolet light, then the energy per quantum begins to be enough to lead to breakage of chemical bonds in molecules. That is why from UV to cosmic radiation, the photons are considered hazardous.

### Some energy relationships—see Chapter 4 for more details:

1 Joule = 0.239 calories (1 calorie = 4.184 Joules)
1 calorie = heat to raise 1 g of water by 1 degree Celsius
1 calorie = 10-3 Calories (food calories)
1 Joule = energy to accelerate 1 Kg object to 1 ms-1
or about amount of energy for one beat of heart
or amount to raise 1 Kg book by 10 cm (~4 in)
1 Joule = 1 watt-second.
burning 60 watt bulb for 1 minute = 3600 Joules

### Back to ozone:

Energy to break double bond of oxygen equivalent to about energy in 242 nm of radiation. Screens out radiation 240 nm and shorter. But UV radiation between 242 and 340 would pass through, and mutation rates are maximum about 280 nm which is the adsorption maximum for DNA. Note figures 2.4 and 2.5 that ozone, which has a weaker bond (only about 1.5 bonds due to resonance structure) will absorb the radiation in this region.

### How is ozone made?

Production:

O2 + photon ® 2 O
O2 + O ® O3

Removal:

O3 + photon ® O2 + O
and slow: O + O3 ® 2 O2

This is called the Chapman cycle. Should produce a steady state level of ozone, which will vary with altitude. Note figure 2.7 which shows most of the ozone is found in the stratosphere.

Can measure in laboratory the rates of various steps and compare with observed ozone. Find ozone concentration lower than predicted by these equations alone. That means there are alternative mechanism of destruction.

These alternative mechanisms involve free radicals such as H× and OH× , which can be formed by UV action on water vapor. (A free radical is an unstable species with an unpaired electron.)

Another free radical is NO× , which is formed naturally to some extent but can be formed in combustion engines. The concern over the supersonic transport airplanes, and the decision by the US not to participate in their production, was based on the concern of release of NO in the troposphere.

But that still should lead to a steady state of a lower concentration. Yet we have been noticing that ozone concentrations have been decreasing the last several years. (Look at data in figures 2.8 and 2.9).

The source of this decrease is from an unlikely place. Through the magic of chemistry, several toxic gases (ammonia, sulfur dioxide) were replaced in refrigeration systems by chlorofluorocarbons. Two examples are:

CFC-12: (Freon 12) CF2Cl2 (dichlorodifluoromethane)
CFC-11: (Freon 11) CFCl3 (trichlorofluoromethane)

These are stable, non toxic molecules with the correct boiling points to serve as gases that are easily liquified and can be used in air conditioners and refrigerators. (Must be able with pressure changes to cycle between liquid and solid state). They don’t break down easily, and therefore survive in the atmosphere and spread to the troposphere.

In the troposphere, they can be dissociated by UV, for example:

CCl2F2 + photon ---> • CClF2 + •Cl

The •Cl can act as a catalyst to degrade ozone:

• Cl + O3 --> • ClO + O2

• ClO + O  --> • Cl + O2

The sum of these equations is

O3 + O ---> 2 O2

Note that • Cl is regenerated. It was the discovery of and study of these processes that led to the Noble prize mentioned in the beginning of the chapter. A clinching piece of evidence is shown in Figure 2.11, which shows a comparison of ozone and • ClO concentrations at various latitudes over Antarctica, where the ozone "hole" was discovered.

The concern is serious. Not just increase in mutation rate, skin cancer, and other direct damage to humans, but destruction of phytoplankton and disruption of the food chain are potential side effects of increased UV penetration.

Production of chlorofluorocarbons have been banned. People working on air conditioners are now required to recapture the freon and not release it to the atmosphere. And the search is on for a replacement for these refrigerant gases.

This is an expensive "fix". Is it worth it? What do you think?

For more information on ozone, check out the link on the course page to some NASA ozone sites.