In a covalent bond, a pair of electrons is shared between two atoms. In a reaction, bond breaking is known as bond fission. Electrons within the bond can be redistributed in one of two ways:
Both electrons go to one of the atoms, leaving a positive and negative atom. This is common when a bond is already polar. \[H\div Cl \Rightarrow H^+ + {^{\bullet}_{\bullet}Cl}^-\]
In this kind, one electron goes to each of the atoms. \[Br \div Br \Rightarrow Br^\bullet + Br^\bullet\] The dot depicts a radical - an unpaired electron. This has a strong tendency to pair up with another electron. Radicals are most commonly formed from a non-polar bond being broken, but not exclusively, particularly in thepresnece of light and in the gaseous phase.
\[O+O_2\Rightarrow O_3\]
Oxygen atoms can be formed from the dissociation of Dioxygen molecules in the stratosphere: \[O_2 +hv \Rightarrow O + O\] This requires alot of energy, which can be provided by UVR or electric discharge.
These oxygen atoms may then collide with other molecules: \[\begin{aligned}
O+O_2\Rightarrow O_3 && \text{Termination}\\
O+O\Rightarrow O_2 && \text{Termination}\\
O+O_3\Rightarrow O_2+O_2 && \text{Termination}\end{aligned}\] When Ozone absorbs UV radiation, it can undergo photodissociation: \[O_3+hv\Rightarrow O_2+O\] This reaction is responsible for the screening effect of ozone.
Filled outer shells are most stable. Radicals are reactive because they try to fill their outer shells. This often results in radical chain reactions. These have three stages:
Initiation
Propagation
Termination
For hydrogen and chlorine: \[\begin{aligned} Cl_2+hv\Rightarrow Cl^\bullet + Cl^\bullet && \text{Initiation due to photodissociation}\\ Cl^\bullet + H_2 \Rightarrow HCl + H^\bullet && \text{Propagation}\\ H^\bullet + Cl_2 \Rightarrow HCl + Cl^\bullet && \text{Propagation}\\ H^\bullet + H^\bullet \Rightarrow H_2 && \text{Termination}\\ Cl^\bullet + Cl^\bullet \Rightarrow Cl_2 && \text{Termination}\\ Cl^\bullet + H^\bullet \Rightarrow HCl && \text{Termination}\\ \\ H_2+Cl_2 \Rightarrow 2HCl && \text{Overall effect of the reaction}\end{aligned}\]
Alkanes react with halogens in the presence of light. Radical symbols are omitted in the following reaction. \[\begin{aligned} Cl_2 + hv \Rightarrow Cl + Cl && \text{Initiation}\\ Cl + CH_4 \Rightarrow HCl + CH_3 && \text{Propagation}\\ CH_3+Cl_2\Rightarrow CH_3Cl+Cl && \text{Propagation}\\ Cl+Cl\Rightarrow Cl_2 && \text{Termination}\\ CH_3+Cl \Rightarrow CH_3Cl && \text{Termination}\\ CH_3 + CH_3 \Rightarrow C_2H_6 && \text{Termination}\end{aligned}\] Overall, Hydrogen Chloride, Chloromethane, and small amounts of ethane are produced.