## 2.3.6. CFCs

(draft)

Authors: Steven Droge

Reviewer: John Parsons

Learning objectives:

You should be able to:

- realize what the ozone layer depletion was all about

- understand why certain replacement chemicals are still problematic

Keywords: Ozone layer, refrigerator, volatile chemicals, spray cans, radicals

Introduction

CFCs (chlorofluorocarbons) were very common air pollutants in the 20th century because they were the basic components of refrigerants and air conditioning, propellants (in spray can applications), and solvents, since the 1930s. They are still very common air pollutants, because they are very persistent chemicals, and emissions do still continue. In the first years as refrigerants, they replaced the much more toxic components ammonia (NH3), chloromethane (CH3Cl), and sulfur dioxide (SO2). Particularly the CFCs leaking from old refrigerating systems in landfills and waste disposal sites caused high emissions into the environment. Typically, these volatile CFC chemicals are based on the smallest carbon molecules methane (CH4), ethane (C2H6), or propane (C3H8). All hydrogen atoms in these CFC molecules are replaced by a mixture of chlorine and fluorine atoms.

Figure 1. Different common refrigerants and their boiling points. Freon 134 is chlorine free. . (Source: Steven Droge)

CFCs are less volatile than their hydrocarbon analogue, because the halogen atoms polarize the molecules, which causes stronger intermolecular attractions. Depending on the substitution with Cl or F, the boiling point can be tuned to the desired point for refrigerating cooling processes. The CFCs are also much less flammable than hydrocarbon analogues, making them much safer in all kinds of applications.

Naming of CFCs

CFCs were often known by the popular brand name Freon. Freon-12 (or R-12) for example stands for dichlorodifluoromethane (CCl2F2, boiling point -29.8 °C, while methane has -161 °C), as shown in Figure 1. The naming reflects the amount of fluor atoms as the most right number. The next value to the left is the number of hydrogen atoms plus 1, and the next value to the left is the number of carbon atoms less one (zeroes are not stated), and the remaining atoms are chlorine. Accordingly, Freon-113 could apply to 1,1,2-trichloro-1,2,2-trifluoroethane (C2Cl3F3, boiling point 47.7 °C, while ethane has -161 °C). The structure of any Freon-X number can also be derived from adding +90 to the value of X, so Freon-113 would give a value of 203. The first numerical is the number of C (2), the second numerical H (0), the third numerical F (3), and the remaining substitutions are by chlorine (C2X6 gives 3 chlorines).

The reason CFC depletes the ozone layer

The key issue with CFC emissions is the reaction under influence of light (“photodegradation”) that ultimately reduces ozone concentrations (“ozone depletion”) in the upper atmosphere (“stratosphere”). Ozone absorbs the high energy radiation of the solar UV-B spectrum (280–315nm), and the ozone layer therefore prevents this to reach the Earth's surface. The even more energetic solar UV-C spectrum (100-280nm) is actually causing the formation of ozone (O3) when reacting with oxygen (O2), as shown in Figure 2. Under the influence of intense light-energy in the upper atmosphere, CFC molecules can disintegrate into two highly reactive radicals (molecules with a free electron . ), for Freon-11:

It is the radical Cl. that catalyzes the conversion of ozone back into O2. The environmentally relevant role of the fluorine atoms in CFCs is that they make these chemicals very persistent after emission, because the C-F bond is one of the strongest covalent bonds known. With half-lives up to >100 years, high CFC levels can reach the upper atmosphere. James Lovelock was the first to detect the widespread presence of CFCs in air in the 1960s, while the damage caused by CFCs was discovered only in 1974. Another undesirable effect of CFC in the stratosphere is that they are a much more potent greenhouse gases than CO2.

Figure 2. The influence of UV on formation of chlorine radicals from CFCs, oxygen radicals from O2 and oxygen radicals from disintegration of O3. Ozone is not formed in the absence of UV (night time), but can still be reacting away by chlorine radicals. (Source: Steven Droge)

CFC replacements.

In 1978 the United States banned the use of CFCs such as Freon in aerosol cans. After several years of observations of the ozone layer depletions globally (Figure 3), particularly above Antarctica, the Montreal Protocol was signed in 1987 to drastically reduce CFC emissions worldwide. CFCs were banned by the late 1990s in most EU countries, and e.g. in South Korea by 2010. Due to the persistency of CFCs it may take until 2050-2070 before the ozone layer will return to 1980 levels (which were bad already).

The key damaging feature of CFCs in terms of ozone depletion is their persistency, so that emissions reach and build up in the stratosphere (starting from 20km above the equator, but only at 7km above the poles). CFC replacement molecules were initially found simply by adding more hydrogens in the CFC structures and somewhat less Cl (HCFCs), but fractions still contributed to Cl. radicals. Later alternatives lack the chlorine atoms and have even shorter lifetimes in the lower atmosphere, and simply cannot form the Cl radicals. These “hydrofluorocarbons” (HFCs) are currently common in automobile air conditioners, such as Freon-134 (do the math to see that there is no Cl, boiling point -26.1 °C).

Figure 3. The 2006 record size hole in the ozone layer above Antarctica (Source https://en.wikipedia.org/wiki/Ozone_depletion)

Still, HCFC as well as HFCs are still very potent greenhouse gasses, so the worldwide use of such chemicals remains problematic and gives rise to new legislations, regulations, and searches for alternatives. R-410A (which contains only fluorine) is becoming more widely used but is 1700 times more potent than CO2 as greenhouse gas, equal to Freon-22. Simple hydrocarbon mixtures such as propane/isobutane are already used extensively in mobile air conditioning systems, as they have the right thermodynamic properties for some uses and are relatively safe. Unfortunately, we did not have the technological skills, nor the awareness to apply this back in the 1930s.