Chlorofluorocarbons

Scott Gunderson

Chlorofluorocarbons (CFCs) are a family of compounds containing chlorine, fluorine, and carbon.[2] CFCs first got their start in the early 1930’s when new innovations such as the refrigerator and air conditioning were beginning to be introduced into homes and businesses throughout the world. CFCs, at the time, were thought to be safe chemical compounds that were nonflammable, had a low toxicity, and had no effect on the surrounding environment. At a 1930 meeting on CFCs the inventor of CFCs inhaled CFC vapor and then blew out a candle to illustrate these safe properties. What the inventor was unable to illustrate were the hidden environmental effects CFCs were to have on the environment until 1974 when F.Sherwood Rowland and Mario Molina published a paper in the journal, Nature, proposing a connection between CFCs and the destruction of the ozone layer.[4]

CFCs were widely used in four areas: aerosol propellants, solvents in the cleaning of electronic components and metals, refrigerants, and the blowing agents in foam plastic manufacturing. In the aerosol area CFCs such as tricholorfluoromethane (CCl₃F, CFC-11) and dichlorodifluoromethane (CCl₂F₂) were normally packed with materials like paint, insecticide, or cosmetic preparations in pressurized canisters. Upon depressurization the propellant vaporizes and expels the materials inside the canister in the form of the aerosol spray.[2] CFCs such as trichlorotrifluoroethane (CCl₂FCClF₂) can dissolve grease and were widely used as solvents in the cleaning of electronic components and metals. In the refrigerants area Freon, a series of CFCs, and CFC-12 were used. These two compounds absorb the heat of vaporization on evaporation resulting in the cooling of the surroundings. These two compounds were commonly used as refrigerants in refrigerators and air conditioners. CFCs were also used in the production of many foamed plastics. The CFCs remain enclosed within the cells of the plastic and increase the thermal insulation of the air-filled foams.[3] The production of these materials using CFCs began to decrease when the connection between CFCs and the destruction of the ozone layer became public in 1974.

Under natural equilibrium, the rate of producing ozone is somewhat equivalent to the rate of destroying ozone, and thus a fairly equivalent concentration of ozone is maintained. The natural destruction of ozone is displayed in figure 1.

Figure1 [2]

Ozone in the stratosphere undergoes photodissociation by absorbing ultraviolet radiation:

O₃ Þ O₂ + O Step I

The free oxygen atom further reacts with another molecule of ozone resulting in destruction of ozone:

O + O₃ Þ O₂ + O₂ Step II

This natural occurring process has been taking place for thousands of years because of its balance with ozone creation and its lack of a catalyst, which would not be consumed in the reaction, thus allowing it to be used over and over again to fuel the reaction.[1] The influence of CFCs has caused the chlorine catalyst to enter this natural process. When CFCs are exposed to the sun in the troposphere, the ultraviolet gives enough energy to the CFC that a single chlorine atom breaks free.[1] The chlorine atom is what reacts with and destroys ozone. In this process, chlorine acts as a catalyst and is recycled each time it is used, thus one chlorine can destroy a large number of ozone molecules. This six-step process is illustrated in figure 2 below.

Figure 2 [1]

I) Ultraviolet radiation strikes a CFC molecule:

ÞÞÞÞÞÞÞÞCFCl₃

_{(UV radiation)}

II) Ultraviolet radiation causes a chlorine atom to break away:

CFCl₂ ÞÞÞÞÞ Cl

III) The chlorine atom collides with an ozone molecule:

Cl ÞÞÞÞÞ O₃

IV) Chlorine atom steals an oxygen atom to form chlorine monoxide and leave a molecule of ordinary oxygen:

ClO O₂

V) A free atom of oxygen collides with the chlorine monoxide:

O ÞÞÞÞÞ ClO

VI) The two oxygen atoms form a molecule of oxygen. The chlorine atom is released and free to destroy more ozone:

O_{2 Cl}

Current figures estimate that there is currently enough chlorine in the atmosphere to continue to affect the ozone layer for 30-100 more years.[1] This estimate reflects the fact government regulations, starting in 1981, have begun to phase-out CFCs and forced the introduction of replacement CFCs by former CFC producing companies.

In response to the scientific consensus that CFCs would eventually deplete the ozone layer, the United Nations Environmental Program began negotiations in 1981 aimed at protecting the ozone layer.[5] In September 1987, the Montreal Protocol on Substances that Deplete the Ozone Layer was signed by 20 nations and took force on January 1, 1989.[5] This treaty called for the limitation of specified CFCs to 50% of 1986 levels by 1998, the freezing of production of specified Halons, convening the signatories yearly reevaluate, and update the Protocol articles in light of new developments.[5] Since this first meeting the parties have met a total of three more times. At a meeting in 1990 the parties passed the 1990 Clean Air Act. This act regulated the production and use of CFCs, specified stepped-up timetables for the total phase out of ozone depleting chemicals by the year 2000, and stated that total HCFC (semi-harmful CFC substitute) production be banned effective January 1,2030. With these regulations in place many companies began to tackle the task of trying to find possible replacements for CFCs.

The ideal CFC substitute should have identical or better performance properties then the CFC it replaces. The substitute must not harm the ozone layer and must have a short atmospheric lifetime to ensure a low greenhouse warming potential. The substitute should also be nontoxic, nonflammable, and chemically and physically stable under normal conditions.[5] The chemical industry has found substitutes that match many but not all of these criteria.

The general strategy has been to incorporate at least one hydrogen atom in the proposed CFC substitutes structure which provides a means for its destruction via hydrogen atom abstraction by tropospheric hydroxyl radicals.[5] The haloalkylradicals formed are then rapidly degraded to acids and CO₂ which are both removed from the atmosphere by natural processes.[5] Since fluorine does not participate in the ozone destruction cycle, substitutes composed of only hydrogen, fluorine, and carbon would be ideal, yet HFCs for every application have not been found yet. This has resulted in the introduction of some HCFC and PFC compounds that would be a tradeoff for the much more harmful CFC compounds. These HCFC and PFC compounds break down more quickly in the atmosphere. Though they have a lower percentage of chlorine and hence a lower ozone depletion potential than CFCs, they could damage the ozone if overused.[2] Table 1 below displays alternatives that have been identified.

Table 1 [5]

*( In the numbering system for CFCs, the digit on the right denotes the number of fluorine atoms in the compound, the second digit from the right indicates the number of hydrogen atoms plus 1, and the third digit from the right indicates the number of carbon atoms minus 1)

Other possible replacements of CFCs experimented with have been hydrocarbons such as butane and propane. These compounds are rather cheap and contain no chlorine, but they are flammable and poisonous. Another option is water and steam but they have only been found to be useful in replacing CFCs as solvents in cleaning.

The manufacture of CFC alternatives has also produced a challenge because these alternatives are far more complex then the original CFCs. The very design feature which makes the alternatives tropospherically labile, the hydrogen atom subsistent, also significantly complicates their manufacture because of potential byproduct formation or catalyst inactivation.[5] HFC-134 will likely be the single largest volume CFC alternative produced, yet no single-step process has been found to produce it easily with no byproducts. Figure 3 illustrates the two-step process that has been commercialized.

Figure 3 [5]

Trichloroethylene reacts with HF in the vapor or liquid phase to form HCFC-133a:

CCl₂=CHCl + 3HF Þ CF₃CH₂Cl + 2HCl

HCFC-133a is then separated and reacts again with HF to form HFC-134a:

CF₃CH₂Cl + EXCESS HF Þ CF₃CH₂F + HCL

Manufacturing facilities for CFC alternatives are just beginning to start full production. Some of these companies include: 3M, Suva, Du Pont, and AlliedSignal Corporation. The size of the markets for the alternatives is estimated to be quite large (several thousand tons/year), but it will not be as large as the prior markets for CFCs themselves.[5] This being cause by the higher cost of the alternatives, which are typically 3-5 times that of CFCs.

Over the past twenty years the usage and production of CFCs have decreased because of environmental awareness and government regulations. Dr. Molina’s published paper in 1974 started the scientific movement toward advanced research of ozone depletion caused by CFCs. This in return jump-started the research and development of CFC substitutes.

References

[1]Http://www.bhs.berkeley.k12.ca.us/departm…cience/APWEB/MEC

AGODZILLA/html/cfcs.html

[2]Http://crucial.ied.edu.hk/pollute/ozone.html

[3] De Grave, Isidor. (1995). Ullaman’s Encyclopedia of

IndustrialChemistry: Vol. A11 (5^th ed., pp461-462). New York: VCH.

[4]Nemecek, Sahsa. (1997,November). Profile: Mario Molina Rescuing

theOzone Layer. Scientific American. Pp. 40-43

[5] Smart, Burce. (1994). Encyclopedia of Chemical Technology: Vol

11(4^th ed., pp 510-515). New York: John Wiley and

Sons.