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Chlorofluorocarbons ($$CFCs$$​) are molecules whereby all the hydrogen atoms have been replaced by fluorine and chlorine.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

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