Alkanes are non-polar hydrocarbons because the electronegativity difference between carbon and hydrogen is not significant enough to create polar bonds. They are also known as paraffins, a term derived from the Latin words parum (little) and affinis (affinity/reactivity), which aptly describes their low reactivity under normal conditions.
The low reactivity of alkanes is attributed to two main factors:
The electronegativity values of carbon and hydrogen are 2.5 and 2.1, respectively. The small difference in electronegativity means the bonding electrons in bonds are shared almost equally, making the bonds non-polar.
Because alkanes are composed of non-polar and bonds, they do not react with polar reagents such as acids, bases, oxidizing agents, or reducing agents under normal conditions.
Alkanes contain only single covalent bonds, which are sigma (σ) bonds. They do not have any pi (π) bonds.
A sigma (σ) bond is very strong because the shared electrons are located directly between the two bonded nuclei, resulting in a strong electrostatic attraction. These strongly held electrons are not easily available for chemical reactions, which contributes to the chemical inertness of alkanes. For more on orbital overlaps, see Shapes of Orbitals→.
Despite their general lack of reactivity under normal conditions, alkanes undergo reactions at high temperatures or in the presence of ultraviolet (UV) light. Two important reactions of alkanes are:
Free radical substitution reactions are initiated by UV light, which generates highly reactive free radicals. Cracking involves heating large alkane molecules to high temperatures to break them down into smaller, more useful molecules.
Not all halogens react with alkanes at the same rate. The order of reactivity is:
| Halogen | Reactivity | Notes |
|---|---|---|
| Explosive | Reaction is uncontrollable; highly exothermic | |
| Vigorous (UV light) | Most commonly used in laboratory reactions | |
| Slow | More selective; requires UV light or heat | |
| Very slow / reversible | Practically does not occur under normal conditions |
This type of reaction involves highly reactive halogen free radicals — species with one or more unpaired electrons — that attack ethane molecules, replacing hydrogen atoms with halogen atoms. The mechanism occurs in three distinct steps.
The reaction is initiated by ultraviolet (UV) light, which provides the energy to break the chlorine-chlorine bond through homolytic fission. Each atom retains one electron from the shared pair, creating two highly reactive chlorine free radicals.
(Here, represents a photon of UV light)
This step is a self-sustaining chain reaction. Once a chlorine radical is formed, it attacks an ethane molecule, abstracting a hydrogen atom to form hydrogen chloride () and an ethyl free radical.
The newly formed ethyl free radical is also highly reactive. It attacks a neutral chlorine molecule, abstracting a chlorine atom to form chloroethane and a new chlorine free radical, which continues the chain.
Further Substitution: If excess chlorine is used, the propagation steps can continue, substituting more hydrogen atoms. This leads to a mixture of chloro-substituted products like 1,1-dichloroethane, 1,1,1-trichloroethane, and eventually, 1,1,1,2,2,2-hexachloroethane ().
The chain reaction eventually stops when two free radicals collide and combine to form a stable molecule. This removes the reactive radicals from the system.