OXIDATION OF ALKANES

Alkanes: Oxidation

Alkanes can be oxidized to carbon dioxide and water via a free‐radical mechanism. The energy released when an alkane is completely oxidized is called the heat of combustion. For example, when propane is oxidized, the heat of combustion is 688 kilocalories per mole.

In a homologous series like the straight‐chain alkanes, the energy liberated during oxidation increases by approximately 157 kilocalories for each additional methylene (CH ₂) unit.

Heat of combustion data is often used to assess the relative stability of isomeric hydrocarbons. Because the heat of combustion of a compound is the same as the enthalpy of that compound in its standard state, and because potential energy is comparable to enthalpy, the differences in heats of combustion between two alkanes translate directly to differences in their potential energies. The lower the potential energy of a compound, the more stable it is. In the alkanes, the more highly branched isomers are usually more stable than those that are less branched.

KINETICS OF ALKANES

Alkanes: Kinetics and Rate

Most reactions require the addition of energy. Energy is needed for molecules to pass over the energy barriers that separate them from becoming reaction products. These energy barriers are called the activation energy, or enthalpy of activation, of the reactions.

At room temperature, most molecules have insufficient kinetic energy to overcome the activation energy barrier so a reaction can occur. The average kinetic energy of molecules can be increased by increasing their temperature. The higher the temperature, the greater the fraction of reactant molecules that have sufficient energy to pass over the activation energy barrier. Thus, the rate of a reaction increases with increasing temperature.

The rate of a reaction also depends on the number of interactions between reactant molecules. Interactions increase in solutions of greater concentrations of reactants, so a reaction rate is directly proportional to the concentration of the reactants. The proportionality constant is called the rate constant for the reaction. Not every collision is effective in producing bond breakage and formation. For a collision to be effective, the molecules must have sufficient energy content as well as proper alignment. If all collisions were effective, every reaction would proceed with explosive force.

Activation energy. The change in structure of each of the reactants as a reaction proceeds is very important in organic chemistry. For example, in the reaction of methane and chlorine, the molecules of each substance must “collide” with sufficient energy, and the bonds within the molecules must be rearranged for chloromethane and hydrogen chloride to be produced. As reactant molecules approach each other, old bonds are cleaved, and new bonds are formed. The cleavage of bonds requires a lot of energy, so as the reaction occurs, the reactant molecules must remain in high‐energy states. When new bonds form, energy is released, and the resulting products possess less energy than the intermediates from which they were formed. When reactant molecules are at their maximum energy content (at the crest of the activation energy curve), they are said to be in a transition state. The energy necessary to drive the reactants to the transition state is the activation energy (Figure 1).

Many organic reactions involve more than one step. In such cases, the reactants may proceed through one or more intermediate stages (either stable or unstable arrangements), with corresponding transition states, before they finally form products (Figure 2).

The overall rate of the reaction is determined, for the most part, by the transition state of highest energy in the pathway. This transition state, which is usually the slowest step, controls the rate of reaction and is thus called the rate‐determining step of the mechanism.

Energy of reaction. The energy of reaction is the difference between the total energy content of the reactants and the total energy content of the products (Figure 3). In ordinary organic reactions, the products contain less energy than the reactants, and the reactions are therefore exothermic. The energy of reaction has no effect on the rate of the reaction. The greater the energy of reaction, the more stable the products.

Effects of temperature on rate of reaction. The rates of organic reactions approximately double with each 10°C rise in temperature. A more quantitative relationship between reaction rate and temperature is given by the Arrhenius equation

HALOGENATION OF ALKANES

Alkanes: Halogenation

The reaction of a halogen with an alkane in the presence of ultraviolet (UV) light or heat leads to the formation of a haloalkane (alkyl halide). An example is the chlorination of methane.

Experiments have shown that when the alkane and halogen reactants are not exposed to UV light or heat, the reaction does not occur. However, once a reaction is started, the light or heat source can be removed and the reaction will continue. The mechanism of the reaction explains this phenomenon.

Halogenation mechanism. In the methane molecule, the carbon‐hydrogen bonds are low‐polarity covalent bonds. The halogen molecule has a nonpolar covalent bond. UV light contains sufficient energy to break the weaker nonpolar chlorine‐chlorine bond (∼58 kcal/mole), but it has insufficient energy to break the stronger carbon‐hydrogen bond (104 kcal/mole). The fracture of the chlorine molecule leads to the formation of two highly reactive chlorine free radicals (chlorine atoms). A free radical is an atom or group that has a single unshared electron.

The bond that is ruptured is said to have broken in a homolytic fashion; that is, each of the originally bonded atoms receives one electron. This initial reaction is called the initiation step of the mechanism. The chlorine free radicals that form are in a high‐energy state and react quickly to complete their octets and liberate energy. Once the high‐energy chlorine free radicals are formed, the energy source (UV light or heat) can be removed. The energy liberated in the reaction of the free radicals with other atoms is sufficient to keep the reaction running.

When a chlorine free radical approaches a methane molecule, a homolytic fission of a carbon‐hydrogen bond occurs. The chlorine free radical combines with the liberated hydrogen free radical to form hydrogen chloride and a methyl free radical. This is called a propagation step, a step in which both a product and a reactive species, which keeps the reaction going, are formed.

A second propagation step is possible. In this step, a methyl free radical reacts with a chlorine molecule to form chloromethane and a chlorine free radical.

When a reaction occurs between free radicals, a product forms, but no new free radicals are formed. This type of reaction is called a termination step because it tends to end the reaction. There are several termination steps in the chlorination of methane.

1. A methyl free radical reacts with a chlorine free radical to form chloromethane.

2. Two methyl free radicals react to form ethane.

3. Two chlorine free radicals react to form a chlorine molecule.

To summarize, this free‐radical chain reaction initially contains few free radicals and many molecules of reactants. As the reaction proceeds, the number of free radicals increases, while the number of reactant molecules decreases. Near the end of the reaction, many more free radicals exist than reactant molecules. At this stage of the overall reaction, termination steps become the predominant reactions. All of the halogenation mechanism reactions occur very rapidly, and the formation of the products takes only microseconds.