Haloalkanes and the environment

In a nutshell

Some haloalkanes, such as $CFCs$ , form chlorine radicals that destroy the ozone layer. $CFCs$ have been banned but some remain in the atmosphere. Alternatives are being looked into and used, to ensure damage to the Earth's atmosphere is minimised.

Free-radical substitution

Free-radical substitution is when a free-radical replaces another atom or group of atoms in a molecule. Alkanes can react with halogens to give a haloalkane, in photochemical reactions. Free-radical substitution has a three step mechanism.

Example

$CH_4 \, + \, Cl_2 \xrightarrow{UV} CH_3Cl \, +\, HCl$

Mechanism for the reaction between methane and chlorine.

Step 1 - Initiation: production of free radicals. This is done via homolytic fission where the energy for the breaking of the bond is provided by energy from sunlight (photodissociation).

$Cl_2 \xrightarrow{UV} 2\, Cl \bullet$ 

Step 2 - Propagation: chain reaction of free radical formation and consumption.

The free radical produced is very reactive and goes onto react with a molecule of methane. The methyl radical can then go on to attack the chlorine molecule forming another chlorine radical. And this process continues.

$Cl\bullet + \,CH_4\rightarrow \bullet CH_3 \,+\, HCl \\\bullet CH_3 \, + \, Cl_2 \rightarrow CH_3 Cl \ +\, Cl \bullet$

Step 3 - Termination: free radicals react together to form a stable molecule. There are many possible reactions for termination. For example:

$Cl\bullet + \,\bullet CH_3 \rightarrow \,CH_3Cl\\\bullet CH_3 \, + \, \bullet CH_3 \rightarrow \,C_2H_6$

The problem is that a mixture of products is obtained. A chlorine radical may go on to react with chloromethane to form dichloromethane, which in turn may go on to form trichloro- and tetrachloromethane. This is not ideal as it requires time and energy to separate the mixture and obtain the desired product. For longer carbon chains, the free-radical attack may happen anywhere along the chain leading to isomers.

Example

Reaction between butane and bromine can lead to the formation of $1$ -bromobutane and $2$ -bromobutane.

Chlorofluorocarbons

Chlorofluorocarbons ( $CFCs$ ) are molecules whereby all the hydrogen atoms have been replaced by fluorine and chlorine.

Example

The molecule chlorotrifluoromethane $(CClF_3)$ contains fluorine and chlorine atoms only.

The inert, non-flammable nature of $CFCs$ pose a problem as they breakdown the ozone layer, by creating holes. The ozone layer $(O_3)$ is a natural sunscreen for the Earth which absorbs UV radiation. This protects the inhabitants of Earth from severe sunburns and skin cancer.

Ozone is made naturally in the atmosphere via free-radical substitution.

Firstly, oxygen in the atmosphere is split into radicals which then go on to react with free oxygen molecules to form ozone.

$O_2 \xrightarrow{UV}O\bullet \ + \ O\bullet \\O_2 \ + \ O\bullet \rightarrow O_3$

Below is a mechanism on how $CFCs$ breakdown ozone in the upper atmosphere.

Step 1 - Initiation

$CCl_3F \xrightarrow{UV} CCl_2F\bullet \ + \ Cl\bullet$

Step 2 - Propagation

$Cl\bullet \ + \ O_3 \rightarrow O_2 \ + \ ClO\bullet \\ClO\bullet \ + \ O_3 \rightarrow 2 O_2 \ + \ Cl\bullet$

You can see that a chlorine radical is used and regenerated during this reaction which means it is acting as a catalyst. Therefore the overall reaction is:

$2O_3 \xrightarrow{UV} 3O_2$

$CFCs$ used to be widely used but the negative environmental impact led to a ban of their use. Alternatives have been developed, such as $HFCs$ and hydrocarbons, which do not contain chlorine. Hydrocarbons contribute to environmental issues as some are greenhouse gases however, they cause less damage than $CFCs$ . Safer alternatives are being researched. The good news is that the ozone is reforming as the holes are slowly becoming smaller.

Nitrogen oxides

Nitrogen oxides are formed by vehicles and thunderstorms. They are not haloalkanes but they do also form free radicals such as $NO\bullet$ which also breaks down the ozone layer. $NO\bullet$ acts as a catalyst.

$NO\bullet \ + \ O_3 \rightarrow NOO\bullet \ + \ O_2 \\NOO\bullet \ + \ O_3 \rightarrow 2O_2 \ + \ NO\bullet$