Overview

Methane (CH₄ ), the simplest of all organic compounds, is the first in the homologous series of alkanes. We shall study this single compound at some length, since most of what we learn about it can be generalized in both theory and application (with minor modification) to any other member of the alkane family.

Hybridization & Structure

Hybridization describes the bonding atoms from an atom's point of view. That is, for a tetrahedrally coordinated carbon, the C atom should have 4 orbitals with the correct symmetry to bond to the 4 atoms. But carbon has an unpaired electron in each of the 2 p orbitals. On this basis, we might expect the C atom to form a compound such as CH₂ which we find to be highly unstable in nature when it is actually observed. What we find is the natural tendency to form as many bonds as possible, which results in the formation of four identical C-H bonds.

The problem with the existence of methane is that the ground-state configuration is of carbon is given by:

In order to provide four unpaired electrons, we promote one of the 2s electrons into the empty 2p orbital. Thus, the first step in our hybridization scheme is the excitation of one of the 2s electrons, resulting in the presence of four unpaired electrons.

The proton that forms the nucleus of a hydrogen atom attracts one of the valence electrons on carbon. This causes an excitation, moving a 2s electron into a 2p orbital. This, however, increases the influence of the carbon nucleus on the valence electrons by increasing the effective core potential..

The combination of these forces creates new mathematical functions known as hybridized orbitals. In the case of 1 C atom attempting to bond with 4 H atoms, four orbitals are required. Thus, the 2s orbital mixes with the three 2p orbitals to form four sp³ hybrids.

In CH₄, four sp³ hybridized orbitals are overlapped by the 1s orbital of the H atom, yielding four covalent  sigma (σ) bonds. How are these four equivalent orbitals to be arranged in 3-dimensional space ? The answer lies in accordance with the Valence Shell Electron Pair Repulsion (VSEPR) theory. In order to minimize the forces of electrostatic repulsion, the electron pairs will arrange themselves in the configuration which allows them to be as far away from each other as possible. They are directed to the corners of a regular tetrahedron. 

Overlap of each of the four sp³ hybridized orbitals of carbon with a 2s orbital of hydrogen results in the methane molecule, with a C atom at the center of a regular tetrahedron, and the 4 H atoms at he corners.  The four bonds are of the same length (1.10 Angstroms) and bond strength (104 kcal / mol). The angle between any two of the orbitals is the tetrahedral angle 109.5 degrees.

This theory fits our requirements. And indeed, this tetrahedral structure of methane has been confirmed by electron diffraction. Note that, unlike the ionic bond (which has equal strength in all directions), the (hybridized) covalent bond is highly directional.

We often write methane in 2-dimensional space as a square with a centered C atom, 4 H atoms at the center of each of the four edges, and connecting C - H dashes to represent each pair of shared electrons. To focus our attention on individual electrons, we may sometimes indicate a pair of electrons by a pair of dots. When we wish to indicate the actual shape of the molecule, we choose simple illustration such as that shown below.

Here the darkened wedge is to indicate a C-H bond emerging from the plane of the screen, while the dashed line is indicative of the C-H bond receding from the plane of the screen. An ordinary line would represent a bond lying in the plane of the screen. 

Physical Properties

Because the methane molecule is highly symmetrical, the polarities of the individual carbon-hydrogen bonds cancel each other out. As a result, the molecule itself is non-polar.

Attraction between such non-polar molecules is limited to van der Waals forces. These forces are extremely weak compared to ionic bonds. Thus, melting and boiling occur at very low temperatures.

                              MP:  - 183 °C            BP:   - 162 °C

Thus, methane is a gas @ standard temperature and pressure (STP).

Methane is odorless and, when liquefied, is less dense than water (relative density = 0.4). In agreement with the rule of thumb that "like dissolves like", it is only slightly soluble in water. But, like most other alkanes, it is quite soluble in organic liquids such as gasoline, ether and alcohol.

Source

Methane is an end product of the anaerobic ("without air") decay of plants and other complex, naturally occurring organic matter. As such, it is the major constituent (up to 97 %) of natural gas. It is the hazardous firedamp of the coal mine, and can be seen as marsh gas bubbling to the surface of swamps. It can be separated from other natural gas components by fractional distillation. But most of it is consumed as fuel without purification.

According to popular theory, the origins of life trace back to a primitive earth surrounded by an atmosphere of methane, water, ammonia and hydrogen. Energy in the form of radiation from the sun and/or lightning discharges broke these simple molecules into reactive fragments or free radicals. Through myriads of combinations of trial and error (with numerous short-lived unsuccessful attempts), these simple organic precursors combined to form larger molecules which eventually yielded the enormously complicated biochemical compounds (amino acids, proteins, nucleic acids: DNA) that currently constitute living organisms.

In summary, the methane generated in the final decay of a once-living organism may well be the very substance from which - in the final analysis - the organism was derived. 

                        "....earth to earth, ashes to ashes, dust to dust........"

Reactions

In its chemical and physical properties, methane sets the pattern for the alkane family. Chemically speaking, it is a very stable compound, and reacts only with highly reactive substances. We will focus primarily on its capacity to undergo oxidation. This includes oxidation reactions with 1) oxygen, 2) any of the halogens (Group VII), and 3) water. 

In the case of reaction with oxygen, we will see that it constitutes the classical exothermic combustion (burning) reaction, in which a hydrocarbon fuel combines with oxygen, resulting in the consumption of oxygen and the production (or evolution) of heat. By-products of combustion include carbon dioxide (the major cause for global warming) and water.

I. Oxidation

The burning of hydrocarbons takes place only at high temperatures, as provide, for example, by a flame or a spark. Once initiated, however, the reaction gives of  heat which serves to mainitin the high temperature and propagate the reaction. The heat of combustion for methane is 213 kcal / mole. 

The combustion process can also be used  in the production of acetylene, C2H6, as follows.

This is a high temperature reaction which requires a heat source to maintain the conditions of the reaction @ 1500 °C.

II. Halogenation

Under the influence of UV light or at e temperature of 250 - 400 °C, a mixture of the two gases methane and chlorine react vigorously. Halogenation is a typical example of a broad class of reactions known as substitution. If we use chlorine as a prototype example of a halogen, then in this reaction scheme, a chlorine Cl2 atom is substituted for a hydrogen atom in methane. We call the product, CH3Cl, chloromethane or methyl chloride (CH3 = methyl group). The hydrogen atom combines with another chlorine atom to form HCl (hydrogen chloride or hydrochloric acid).  

The methyl chloride can then itself undergo further substitution to form more hydrogen chloride and the compound CH2Cl2, dichloromethane or methylene chloride (CH2 = methylene)

Similarly, chlorination may continue to yield CHCl3, trichloromethane or chloroform, and CCl4, tetrachloromethane or carbon tetrachloride.

Chlorination of methane may yield any one of four hydrocarbons , depending on the stage to which the reaction is carried.  Can we control this reaction so that methyl chloride is the principal reaction product ? The reaction may be limited to almost entirely monochlorination if we use a large excess of methane. In this case, at all stages of the reaction, methane greatly exceeds methyl chloride in concentration. (Chlorine is more likely to successfully replace an H atom on methane than an H atom on methyl chloride). The use of a large excess of one reactant is a common device of the organic chemist who wishes to limit a reaction to only one of a number of reactive sites in the molecular conformation of that reactant.  

Bromination takes place somewhat less readily than chlorination, and iodination simply does not occur. Alternatively, fluorine gas reacts so vigorously that, even in the dark at room temperature, the reaction must be carefully controlled. We can therefore arrange the halogens in order of relative reactivity. 

                                                         F2   >   Cl2   >   Br2   >   I2

This same order of reactivity holds for the reaction of the halogens with other alkanes, and indeed, with most other organic compounds. The spread of reactivities is so great that only chlorination and bromination proceed at such rates as to be generally useful.

Mechanism of Reaction

It has been shown experimentally that the wavelength of light that induces chlorination is that known independently to cause dissociation of chlorine molecules. It my therefore be reasonable to assume that the chlorination process is initiated by the breakdown of chlorine molecules into chlorine atoms.

The breaking of the chlorine molecular bond requires energy which we call the bond dissociation energy. In this case, it is equal to 58 kcal/mole. The energy is supplied as either heat or light. In this case, each atom retains one electron of the pair that formed the covalent bond. This odd electron is unpaired, as opposed to the other 6 valence electrons surrounding each chlorine atom. An atom or group of atoms possessing an odd (unpaired) electron is called a free radical. In writing the symbol for a free radical, we generally include a dot to represent the odd electron.   

( 1 )                                                                           

Once formed, what is a highly reactive chlorine atom most likely to do ? Of the likely collisions, that with a reactive chlorine atom and a methane molecule is both probable and productive: that is, it causes a net change (vs. a collision with another chlorine molecule).

( 2 )                                  

The chlorine atom abstracts a hydrogen atom, with one electron, to form a molecule of hydrogen chloride. The methyl group is left with an odd, unpaired electron (i.e. the carbon atom has only 7 electrons in its valence shell). Thus, the chlorine atom free radical has been consumed, and the methyl radical has been formed in its place. This is step (2) in the reaction mechanism.

Now, what is the methyl radical most likely to do ? The collision of a methyl radical with a chlorine atom is the most productive possibility. Thus, the methyl radical abstracts a chlorine atom, with one of the bonding electrons, to form a molecule of methyl chloride. The other product is a chlorine atom. This is step (3) in the reaction mechanism.

(3)                                     

Then (2), (3), (2), (3), etc.

The chlorination of methane is an example of a chain reaction -- where a product of a reaction is itself a reactive particle which can cause more similar reactions. In a chain reaction, each one of a series of steps generates a reactive substance that brings about the next step in the reaction sequence. Chain reactions all include a chain-initiating step (1), chain propagating steps (2) and (3), and finally a chain terminating step (4) , in which reactive particles are consumed but not generated. In this case, the termination proceeds as follows:

    (4)                                                         

or  (5)                                                   

or  (6)                                                 

Under one set of conditions, about 10,000 molecules of methyl chloride are formed for every single photon of light absorbed. Each photon cleaves one chlorine molecule to form two chlorine atoms, each of which initiates a chain reaction. Each chain consists of ~ 5000 repetitions of the chain propagating cycle before it is finally terminated.

Inhibitors

It has been shown experimentally that even a small amount of oxygen introduced into the reaction vessel will slow down the reaction for a period of time, after which the reaction proceeds normally. The length of the period depends on the amount of oxygen present.

Oxygen is believed to react with a methyl radical in order to form a new free radical:

The CH3OO radical is much less reactive than the methyl radical, and can do little more to propagate the chain. By combining with a methyl radical, one oxygen molecule breaks a chain, and thus prevents the formation of thousands of molecules of methyl chloride. This slows down the rate of reaction enormously. After the oxygen has depleted itself, the reaction returns to its normal rate based as determined by the particular conditions of chemical equilibrium. 

A substance that slows down or stops a reaction even when present in small amounts is called an inhibitor. The period of time during which inhibition lasts, and after which the reaction proceeds normally, is called the inhibition period. We shall frequently encounter the use of oxygen to inhibit free radical reactions in organic chemistry.

The Methyl Radical

Let us consider the structure and properties of the highly reactive methyl free radical. Our observations and conclusions will apply not only to this species, but also to numerous other free radicals we will encounter in the numerous reaction mechanisms covered in the field of organic chemistry. 

Let us begin by rolling back the discussion of methane to the electron configuration of carbon. 

In order to provide four unpaired electrons, we promote one of the 2s electrons into the empty 2p orbital. Thus, the first step in our hybridization scheme is the excitation of one of the 2s electrons, resulting in the presence of four unpaired electrons.

Ina similar manner to boron on boron trifluoride, carbon here is bonded to three other atoms. Hybridization of the 2s electron orbital and 2 of the 2p orbitals provides the necessary hybridized orbitals: three strongly directed sp2 hybridized orbitals which all lie in a single plane that includes the carbon nucleus, and are directed to the corners of an equilateral triangle.  

If we arrange the carbon and three hydrogen atoms of a methyl radical to permit maximum overlap of orbitals, we obtain the following structure. It is planar, with the carbon atom at the center of a triangle and the three hydrogen atoms at the corners. Every bond angle is 120 degrees.        

So where is the odd electron ? In forming the sp2 hybridized orbitals, the carbon atom has used only two of its three p orbitals. The remaining p orbital consists of two equal lobes, one lying above and the other lying below the plane of the three sp2 orbitals. This 'split' p orbital is occupied by the odd electron.

*Note: While it has been conjectured that the methyl radical might adopt a pyramidal configuration (e.g. ammonia and its lone pair of electrons), spectroscopic studies indicate that the methyl radical is actually planar (or nearly so). In addition, there is substantial stereoscopic evidence that most free radicals are planar. If they are  pyramidal, then they undergo rapid inversion like that of the ammonia molecule. We will see how, in a vast number of organic reactions, these conformations lend themselves particularly well to the ideal structure of a user friendly reactive intermediate or catalyst.