Volume 4

Methane, CH4, the Simplest Organic Molecule • How Can We Explain That CH4 Is Tetrahedral? Hybridization • Alkanes with Two or More Carbon Atoms. Homology. Nomenclature of Linear Alkanes • Isomerism of Alkanes. Nomenclature of Branched Alkanes Cycloalkanes • Natural Occurrence, Preparation, and Applications of Alkanes • Chemical Properties of Alkanes • Exercises

2.2.1. Methane, CH4, the Simplest Organic Molecule. The main component of natural gas (up to 90%) is methane, the simplest organic compound. The formula of methane is CH4. Pure methane is a nontoxic, odorless, and colorless gas. The boiling and melting points of methane are −162 and −183 oC, respectively. As follows from the molecular mass of CH4, 16 a.m.u., methane is almost two times less dense than air.
Digression. Surprisingly many people think that methane has a rotten egg smell and is poisonous. Neither is true. Just like nitrogen, inert gases, and carbon dioxide, methane is an asphyxiant, a nontoxic gas that, unlike oxygen, does not support breathing. The unpleasant smell of natural gas used in the stoves does not come from methane itself, but is rather due to the odorant, a chemical compound that is added to the gas on purpose, as a safety measure to alert the user to a leak. The greatest risk of a stove gas leak is the explosive nature of mixtures of methane and air in a greater than 5:95 ratio by volume. The smell of the odorant used for stove gas is strong enough to spot the problem with the human nose as a highly sensitive detector long before too much gas has leaked into the air to create an explosion hazard.

The most common odorant used for natural gas is isopropyl mercaptan, (CH3)2CHSH. The odor of this organosulfur compound is so strong that just a tiny harmless amount of it added to the gas does the job well. Since the human nose can detect mercaptans at concentrations as low as a few ppb (parts per billion!), about 1 ppm (parts per million) of the odorant is usually added to natural gas. The odor is totally destroyed as the additive is burnt along with the gas in the stove.

To get an idea of the odor strength of mercaptans, here is a story from my personal experience. I once had the "privilege" to take a ride in a van that the day before was used to transport chemicals. One of those chemicals was isobutyl mercaptan, CH3CH2CH(SH)CH3, a liquid that boils at 98 oC, just 2 degrees below the boiling point of water. The amount of the chemical was 1 g and it was hermetically confined in a flame-sealed glass ampoule. As the chemicals were unloaded, the ampoule accidentally cracked inside the van. There was no spill, but vapors of the mercaptan naturally entered the atmosphere inside the vehicle. The cracked ampoule was immediately removed, after which the van was kept with all of its doors wide open for at least 16 hours. Even after that, the stench inside the van was hardly bearable.
4.2.2. How Can We Explain That CH4 Is Tetrahedral? Hybridization. It has been experimentally established that the molecule of methane is tetrahedral (Figure 4-3).
Figure 4-3. The molecule of methane (CH4) is tetrahedral (left: modified from source; right: source).

How could we rationalize the ideally symmetrical tetrahedral structure of methane with all four C-H bonds of exactly the same length (1.54 Å) and all four identical H-C-H angles of 109.5o (109o 28' to be precise)? For that, we use the concept of orbital hybridization. First, let us recall that the electron configuration of a carbon atom C in the ground state is 1s2 2s2 2p2. Exciting the ground state with energy produces the more reactive, excited state C*, in which one of the two 2s electrons has moved to the vacant 2p orbital, which lies higher in energy (Figure 4-4). The electron configuration of the excited state C* is: 1s2 2s1 2p3.
Figure 4-4. Excitation of a carbon atom in the ground state (C) to the excited state (C*).

In the excited state C*, the three 2p orbitals and one 2s orbital bearing one electron each mix together to give four identical hybrid orbitals (Figure 4-5). This process is called orbital hybridization or just hybridization. This particular type of hybridization is denoted sp3 because it involves one s and three p orbitals. An sp3 orbital is similar in shape to a p orbital, except the dumbbell is asymmetric with one lobe being larger in size than the other (Figure 4-5).
Figure 4-5. Formation of four identical sp3 orbitals from one s orbital and three p orbitals (sp3 hybridization).

Now that we have four identical sp3 orbitals with one electron in each, we have to arrange them in 3-dimensional space in such a manner that (a) they all remain equal and (b) the repulsion between the negatively charged electrons in each of these four orbitals is minimized. There is only one way to fulfill these two requirements, namely the arrangement must be tetrahedral, as displayed in Figure 4-6.
Figure 4-6. Tetrahedral arrangement of four sp3 orbitals of a carbon atom in the excited state (C*).

The sp3-hybridized carbon atom interacts with four H atoms, whose 1s orbitals bearing one electron each overlap with the sp3 orbitals on the C* to give a molecule of methane, CH4 (Figure 4-7).
Figure 4-7. Building a molecule of methane (CH4) from one sp3-hybridized carbon atom and four hydrogen atoms.

Orbital hybridization is a conceptual mathematical model. The mixing of atomic orbitals into hybrids is a hypothetical rather than physical process, originally proposed by the outstanding American scientist Linus Carl Pauling (1901-1994), who won the Nobel Prize in Chemistry in 1954 and the Nobel Peace Prize in 1962. Since Pauling published his theory in 1931, the concept of hybridization has been widely used in chemistry to rationalize many experimental observations. We will appreciate the helpfulness of orbital hybridization models many more times in this course.

4.2.3. Alkanes with Two or More Carbon Atoms. Homology. Nomenclature of Linear Alkanes. Suppose we have two sp3-hybridized carbon atoms, and these two atoms bond to each other. One of the four sp3 orbitals of one C atom and one of the four sp3 orbitals of the other C atom overlap to form a covalent bond by sharing their electrons. This process is illustrated in Figure 4-8.
Figure 4-8. The formation of a C-C bond between two sp3-hybridized carbon atoms. Only the overlapping orbitals are shown for clarity. All other sp3 orbitals are presented as dotted lines.

After the C-C bond has been formed by two sp3-hybridized carbon atoms, the remaining six sp3 hybrid orbitals bearing one electron each are available for overlap with 1s orbitals of hydrogen atoms. Each of these six sp3 orbitals (depicted as dotted lines in Figure 4-8) overlaps with the 1s orbital of a hydrogen atom, the way it happens in the formation of a molecule of methane (Figure 4-7). As a result, a molecule of ethane, C2H6, is produced. All six C-H bonds of ethane are equivalent and the geometry around the carbon atoms in C2H6 is tetrahedral, like in methane.

As shown in Figure 4-9, one can similarly build molecules of hydrocarbons using more sp3-hybridized C atoms, such as CH3-CH2-CH3 (propane) and CH3-CH2-CH2-CH3 (butane). Methane, ethane, propane, and butane all belong to the class of organic compounds that are called alkanes. The general chemical formula for alkanes is CnH2n+2. From this general formula, we see that the alkane is methane (CH4) for n = 1, ethane (C2H6) for n = 2, propane (C3H8) for n=3, and so on and so forth. Like in methane and ethane, carbon atoms in all other alkanes display a tetrahedral geometry.
Figure 4-9. Formulas of alkanes containing 2 (ethane), 3 (propane), and 4 (butane) carbon atoms.

In chemistry, and especially in organic chemistry, it is often important to portray a 3-dimensional molecule on a sheet of paper or computer screen. This is conventionally done by drawing chemical bonds using a wedge and dash projection, as illustrated at the bottom of Figure 4-9. To exhibit the tetrahedral geometry around the carbon atoms of ethane, propane, and butane (Figure 4-9), we draw the regular solid lines for the bonds that lie in the plane of the page, the wedge-shaped lines for the bonds protruding from the plane toward us, and the dashed lines for those that extend away from us.

By conceptually inserting a CH2 group in a C-H bond of methane, CH4, we produce a molecule of ethane, C2H6. If we insert a CH2 group in the C-C bond or a C-H bond of ethane, a molecule of propane, C3H8, will be the result. Repeating such CH2 insertions over and over again leads to alkanes with longer chains. Organic molecules that differ from one another by one or more -CH2- bridges are called homologues. Long-chain alkanes composed of dozens of carbon atoms are known. In this course, we will deal with alkanes CnH2n+2 with n up to 10. The chemical names of the first ten alkanes along with their formulas are listed in Table 1. These names must be memorized.

Table 1. Nomenclature of alkanes CnH2n+2 (n = 1 – 10).
4.2.4. Isomerism of Alkanes. Nomenclature of Branched Alkanes. For the first three alkanes, CnH2n+2, methane (n = 1), ethane (n = 2), and propane (n = 3), there is only one way to connect the carbon and hydrogen atoms in the molecules. This is no longer the case for alkanes with n = 4 (butane) and higher. Let us consider possible carbon skeletons that can be built using four carbon atoms. As shown in Figure 4-10, there are two ways four C atoms can be connected to yield two distinct skeletons, linear and branched.
Figure 4-10. Linear and branched skeletons that can be built using four carbon atoms.

Let us now put the "flesh" of hydrogen atoms on each of these two carbon skeletons to make each carbon atom tetravalent. Having done that (Figure 4-11), we can see that two molecules are produced, which have the same composition (C4H10) but different structures (connectivity of atoms). Chemical compounds whose molecules have the same number of atoms of each element (same composition) yet display a different arrangement of these atoms (different structure) are called isomers. This is the first time that we have encountered the phenomenon of isomerism in our course.
Figure 4-11. Linear and branched isomers of butane.

Do isomers have different properties despite the fact that their composition is the same? The answer is "yes", because it is not just the composition but also connectivity of the same set of atoms that determines intermolecular interactions, stability, reactivity, and other physical and chemical properties of substances. For example, linear butane boils at ~0 oC, whereas the boiling point of the branched isomer, isobutane, is approximately -12 oC.

There are many types of isomerism. The one we are currently considering is called skeletal isomerism because such isomers differ in the structure of the carbon skeleton, the way carbon atoms are connected to one another in isomeric molecules. Later in our course, we will encounter other types of isomerism.

Isomerism in organic chemistry is somewhat similar to allotropy in inorganic chemistry. Allotropes are simple substances that are composed of atoms of the same element yet have different structures. As we already know, this difference can have a dramatic effect on many properties of allotropes. Just recall diamond and graphite, the two allotropes of carbon, or highly toxic and pyrophoric white phosphorus and nontoxic and air-stable red phosphorus. Likewise, isomers of the same compound can display different properties, although the difference is seldom as dramatic as between allotropes. As a notable example, while many isomers of C20H12 hydrocarbons are low-hazard chemicals, one, benzo[a]pyrene, is extremely carcinogenic.
Digression. The anti-inflammatory drug ibuprofen, so widely used that it has dozens (!) of brand names, is a 1:1 mixture of two isomers, of which only one is active. The other one is inactive but, fortunately, totally harmless. Naturally, the question arises as to why do we have to take both isomers of ibuprofen? Is it not possible to use only the active one? In that case we would not be taking the chemical substance that is useless, even though it does not do any harm to our bodies. Furthermore, the bottling, transportation, storage, and many other costs to bring ibuprofen to the pharmacies and hospitals would be reduced dramatically. The problem is that the production cost of just the right isomer of ibuprofen is prohibitively high. It is still much more economical to manufacture a 1:1 mixture of the two isomers while absorbing all other costs associated with only 50% of the product being active. In many other cases, isomers should be thoroughly separated prior to use though, especially for some pharmaceuticals, agrochemicals, and materials for electronics.
As the number of carbon atoms constituting an alkane molecule grows, the number of isomers grows, too. And how! While for butane (n = 4 in CnH2n+2) only two isomers are possible (Figure 4-11), pentane (n = 5) has three isomers, hexane (n = 6) five, and decane (n = 10) already 136. The number of all possible isomers of higher alkanes is mind boggling.

Do we have to come up with a common name for each of these isomers and memorize them all? That is impossible. We can and should remember the names of the two isomers of butane (Figure 4-11) and all three isomers of pentane (Figure 4-12), but to name the rest of branched alkanes we should know and use the rules of organic chemistry nomenclature.
Figure 4-12. Three possible isomers of pentane.

To name a branched alkane, we have to follow the rules presented below, step by step.

Step 1. Find the longest continuous chain of carbon atoms in the structure. Pay attention, as in some cases branched alkanes are drawn in a confusing manner, such that the longest chain is not lined up in a row but rather snakes through branches. For example, a first quick glance at the structure below might tempt one to think that the longest continuous chain is positioned horizontally and has 8 carbon atoms, drawn in red below.
That would be a mistake. On taking a closer look at the structure, we can find a longer, 9-carbon atom straight chain in the molecule, drawn in blue below.
For clarity and convenience, the structure can be redrawn as follows.
Note that the originally drawn and redrawn structures are the same compound. We always keep in mind that all atoms bonded to each carbon in alkanes do not lie in the same plane but form a tetrahedron. Moreover, there is free rotation about the carbon-carbon bond in alkanes. This rotation is very fast, vastly faster than the cartoon in Video 4-1 displays. At room temperature, the C-C bond of ethane rotates roughly a hundred billion (1011) times per second!
Video 4-1. Free rotation around the C-C bond of ethane (source).

Step 2. Now that the longest straight chain has been identified, we number all carbon atoms in it, starting at the end nearest a branching point. In our case, the numbering will look as follows.
Or like the one below, if you wish to use the redrawn formula.
Since the number of carbon atoms in the longest straight chain is 9, a part of the name of our molecule will be the name of the parent linear alkane containing 9 carbon atoms, which is nonane (Table 1).

Step 3. Find the branches hanging off the main chain. In our molecule, there are two of those. One is at carbon atom number 2 and the other at carbon atom number 6.

Step 4. Count the number of carbon atoms in each branch. In our case, the branching group attached to carbon #2 is CH3 and the one attached to carbon #6 is CH2CH3 (C2H5). These groups, called alkyls, are derived by conceptually removing a hydrogen atom from the corresponding alkane. Alkyl groups are named by adding the suffix -yl to the alkane prefix. The CH3 group is called methyl, CH3CH2 (or C2H5) ethyl, CH3CH2CH2 (or C3H7) propyl, etc. Note that two types of propyl are possible. One derives from removal of a hydrogen atom from a terminal carbon of propane. This type of the propyl group, CH3CH2CH2, is called just propyl or sometimes normal propyl (n-propyl) to emphasize that it is linear. The other type of propyl is isopropyl, (CH3)2CH, which is formed upon removal of a hydrogen atom from the central atom of propane. The names of the groups (substituents) at C #2 of our molecule is methyl and at C #6 ethyl. Now we can name the compound.

Step 5. Start naming the compound by placing the first branching number, followed by a hyphen, followed by the name of the group and another hyphen: "2-methyl-". Next we place the second (and last in our case) branching number, followed by a hyphen and the name of the group: "2-methyl-6-ethyl". Finally, we add the name of the parent linear alkane, the one with the same number of carbon atoms as in the longest carbon chain of our molecule, which is nonane (see Steps 1 and 2 above): 2-methyl-6-ethylnonane. This is the name of our branched alkane.

Now let us consider another molecule that is similar to 2-methyl-6-ethylnonane, except that it has one more methyl group at carbon #4, as shown below.
Would we name this molecule 2-methyl-4-methyl-6-ethylnonane? No. If the same alkyl group is present on the main chain more than once, the name of that alkyl group is not repeated. Instead, (a) the numbers of the C atoms bearing the identical groups, separated by commas, are included and (b) the number of these identical groups is specified using the prefixes di-, tri-, tetra-, penta-, etc. In accordance with this rule, the name of our alkane is 2,4-dimethyl-6-ethylnonane.

Now let us practice the reverse thing, the drawing of the molecule of an alkane from its name. For example, we need to draw the structure of 2-methyl-4,5-diethyloctane. First, we draw the longest straight carbon atom chain of eight carbon atoms, as indicated by "octane" at the end of the name. Then we number the atoms in the chain and attach one methyl to carbon #2, one ethyl to carbon #4 and one ethyl to carbon #5.
4.2.5. Cycloalkanes. Three or more CH2 units can form cyclic molecules that are called cycloalkanes (Figure 4-13). The general formula for cycloalkanes is obviously CnH2n. To name a cycloalkane, just count the number of CH2 groups in the cycle, name the corresponding linear alkane (Table 1), and add the prefix "cyclo".
Figure 4-13. Cycloalkanes composed of 3 to 8 CH2 units.

As we know from the above, carbon and/or hydrogen atoms bonded to a carbon atom in an alkane form a tetrahedron around that carbon. Except for cyclopropane, all cycloalkanes are nonplanar, striving to maintain a tetrahedral geometry to the greatest extent possible. However, a tetrahedral geometry around the carbon atoms of cyclobutane and — obviously — cyclopropane is merely impossible. These smallest 3- and 4-membered cycloalkane rings are strained, which makes their chemistry rich and distinct from the chemistry of other cycloalkanes and alkanes. In cyclopentane and especially in cyclohexane and larger cycles, the geometry around the carbon atom is tetrahedral. Consequently, their chemical behavior is similar to that of alkanes. Cyclohexane is produced on a very large scale for making nylon.

4.2.6. Natural Occurrence, Preparation, and Applications of Alkanes. Natural gas and oil are the main sources of a broad variety of alkanes. Natural gas consists of methane (up to >90%) and smaller quantities of ethane, propane, and isomers of butane and pentane. Oil is a mixture of various hydrocarbons, including both linear and branched alkanes. Some pure liquid alkanes are isolated from oil by distillation to be employed as solvents. Methane is used in huge quantities to produce hydrogen in the nickel-catalyzed high-temperature (700–1,100 oC) steam reforming process according to the equation below.

CH4 + H2O = CO + 3 H2

The largest amounts of alkanes, however, are burnt to produce energy. Although there is no need to isolate pure individual alkanes for use as fuels, different engines require different alkane fractions from crude oil distillation (Figure 4-14). For instance, gasoline (petrol) is needed for internal combustion engines in cars, kerosene (jet fuel) for aviation engines, diesel fuel for diesel engines, and fuel oil for heating furnaces.
Figure 4-14. Fractional distillation of crude oil (source).

It is, however, not just the boiling point range that is important for a particular type of fuel, but also the composition of the fraction. You have seen the octane rating numbers at many gas stations, but do you know what these numbers mean? In the U.S.A. and Canada, there are usually three grades of gasoline: 87 (regular), 88–90 (mid-grade), and 91–94 (premium). These numbers are a measure of how much compression a particular grade of gasoline can withstand before it detonates spontaneously (engine knocking). The higher the octane number, the higher compression is possible in the combustion chamber without the knocking and, consequently, the higher the performance of the engine.

For octane number determination, mixtures of isooctane (2,2,4-trimethylpentane) and heptane in various ratios are used. Isooctane and heptane are assigned octane numbers of 100 and 0, respectively (Figure 4-15). First, the anti-knocking capacity is measured for a gasoline sample, using a special test engine. It is then determined in what ratio isooctane and heptane should be blended in order for the mixture to have the same anti-knocking capacity in the same engine. The volume percentage of isooctane in the mixture is the octane number of the fuel tested. For instance, the regular gasoline with the octane number 87 can withstand the same compression as an 87:13 (volume/volume) blend of isooctane and heptane.
Figure 4-15. The two alkanes used in octane number determination.

The branch of chemistry that deals with studies of crude oil is called petrochemistry. Industrial plants for converting crude oil to useful products, called oil refineries, employ numerous chemical transformations as well as separation techniques and purification processes. It is worth to note that, in principle, alkane fuels can be produced without oil. During World War II, Germany, a country without natural oil reserves, was blockaded from external oil supplies and had to use synthetic oil, which was produced from coal by various chemical methods.

Although methods exist for small-scale preparation of alkanes, these methods are rarely used in chemical laboratories for two main reasons. First, many pure alkanes are commercially available from petrochemical and fine chemical industries. Second, laboratory methods to prepare alkanes are rather limited. If you are interested, you may familiarize yourself with two older methods, the Wurtz reaction and the Kolbe electrolysis, developed by Charles Adolphe Wurtz (1817-1884) and Hermann Kolbe (1818-1884), respectively. In organic chemistry, particularly important chemical transformations are often named after their discoverers. Such transformations are called Organic Name Reactions. We will encounter more of such reactions in due course.

4.2.7. Chemical Properties of Alkanes. An alternative name for alkanes is paraffins. The word paraffin is derived from two Latin words parum and affinis, together meaning "insufficient affinity" (insufficient reactivity). Indeed, alkanes are very poorly reactive compounds. Three factors contribute to the high chemical inertness of alkanes.

First, Lewis structures of alkanes show that all C atoms in their molecules have 8 electrons in the outer shell (Figure 4-16). In accord with the octet rule, the carbon atoms of alkanes are "happy". So are the hydrogen atoms, which all have two electrons in their outermost shells. There are neither lone electron pairs on alkane molecules, which would make them reactive toward electron acceptors, nor low-lying vacant orbitals for reactions with electron donors.
Figure 4-16. Lewis structures of methane, ethane, and propane.

Second, all C-C and C-H bonds of alkanes are nonpolar covalent, as follows from the very close Pauling electronegativity values for C (2.5) and H (2.1). Consequently, alkane molecules lack significant partial positive or negative charges, which would attract electron-deficient and electron-enriched reagents.

Third, the C-H and C-C bonds of alkanes are very strong and hard to break. The strength of a chemical bond is conventionally measured as the energy required to homolytically cleave this bond (Bond Dissociation Energy, BDE). The BDE value for a C-H bond of methane is 105 kcal/mol and for the C-C bond of ethane 90 kcal/mol (Figure 4-17). These are huge numbers, which you do not have to memorize. Just remember that alkanes have exceptionally strong bonds and that it takes a lot of energy to break them.
Figure 4-17. Highly energy-demanding homolytic cleavage of a C-H bond of methane and the C-C bond of ethane.

Given the high inertness of alkanes, there are two ways to make them react. One strategy is to employ extreme reaction conditions, such as burning. Methane and other alkanes readily burn in air.

CH4 + 2 O2 = CO2 + 2 H2O

The other strategy employs milder conditions and highly aggressive reagents. Methane and other alkanes react with chlorine gas, Cl2. To occur, however, this reaction must be initiated by light. Chlorine is yellow-green and methane is colorless. A greenish-yellow gaseous mixture of Cl2 and CH4 shows no sign of reaction in the dark. But, once such a mixture is exposed to diffused light, the color gradually fades due to the consumption of the chlorine in a chemical reaction. On exposure to direct bright sunlight, a mixture of CH4 and Cl2 may even explode.

The light-promoted reaction between methane and chlorine occurs according to the following equation.

CH4 + Cl2 = CH3Cl + HCl

How does this reaction occur? Why is light needed for this reaction to happen? In other words, what is the mechanism of this reaction, the sequence of chemical events that lead up to the observed products? The mechanism of the light-induced chlorination of methane (Figure 4-18) and other alkanes is well known from many detailed experimental studies.

The first step of the reaction of Cl2 with CH4 is the light-induced dissociation of a small fraction of the Cl2 molecules to chlorine atoms (Figure 4-18, Step 1). The energy of the light is needed to break the Cl-Cl homolytically. The chlorine atoms generated in this initiation step are highly reactive, vastly more reactive than molecules of chlorine, Cl2. If you do not remember why, go back to Volume 2. In Cl2, each chlorine atom has 8 valence electrons and is therefore "happy". In contrast, a Cl atom has only 7 valence electrons and is therefore "unhappy", seeking any opportunities to expand its incomplete outer shell to an octet. In their "pursuit of happiness", these "hot" chlorine atoms pluck a hydrogen atom from a molecule of methane (Figure 4-18, Step 2).
Figure 4-18. Mechanism of light-induced chlorination of methane.

The hydrogen atom abstraction (Step 2) gives rise to a stable molecule of HCl along with a highly reactive species, a methyl radical, •CH3. Radicals (often called free radicals) are molecules that contain at least one unpaired electron. A middle or superscript dot placed right before or immediately after a chemical formula is conventionally used to denote a radical. Numerous atoms and ions also bear unpaired electrons and are therefore radicals. While some of them are rather stable, the vast majority of organic radicals such as the methyl radical (•CH3) are highly unstable and extremely reactive.

Once produced in Step 2, the •CH3 radical needs to find a way to stabilize by forming a forth bond to the carbon atom. This happens when •CH3 collides with a molecule of chlorine and abstracts a chlorine atom from it (Step 3). As a result of this abstraction, a stable molecule of chloromethane, CH3Cl, is produced along with a chlorine atom. The chlorine atom thus formed then abstracts an H atom from another molecule of methane to give HCl and •CH3, exactly in the same way as in Step 2.

The continuous formation of •CH3 in Step 2 and chlorine atoms (Cl•) in Step 3 is like a growing chain of two alternating events. In chemistry, such processes are called chain reactions. Because our reaction involves free radicals as key intermediate reactive species, it is categorized as a free radical chain reaction.

As the reaction between CH4 and Cl2 occurs, the amount of CH3Cl produced increases. This reaction product starts to react with chlorine in the same manner, and by the same mechanism, as CH4. The highly reactive chlorine atoms can abstract hydrogen atoms not only from CH4, but also from CH3Cl (Figure 4-19).
Figure 4-19. Abstraction of a hydrogen atom from chloromethane by a chlorine atom.

Like •CH3, the •CH2Cl radical produced can and does abstract a chlorine atom from Cl2 to give a molecule of dichloromethane, CH2Cl2 (Figure 4-20).
Figure 4-20. Abstraction of a chlorine atom from Cl2 by a chloromethyl radical.

Similarly, the CH2Cl2 formed is then chlorinated to chloroform, CHCl3, which, in turn, also reacts with chlorine to give carbon tetrachloride, CCl4. Unless a huge excess of either CH4 or Cl2 is used for the reaction, all four chlorinated products are formed. With Cl2 in excess, only CCl4 is produced, whereas with methane in large excess, CH3Cl is the main product.

Chlorination of methane is performed on a huge industrial scale to make all four possible chlorinated products, CH3Cl, CH2Cl2, CHCl3, and CCl4. These four products have sufficiently different boiling points to be easily separated by distillation (Table 2).

Table 2. Names and boiling points of chlorinated methanes.
Some important uses of chlorinated methanes are listed below.

- CH3Cl: as a reagent in industrial synthesis of organic compounds and silicon polymers.

- CH2Cl2: as an excellent solvent capable of dissolving a broad variety of compounds. Dichloromethane is also used to decaffeinate tea and coffee.

- CHCl3: as a solvent and reagent in organic synthesis. Chloroform is the starting material for the manufacturing of Teflon®, a unique fluoropolymer for making nonstick cookware and a variety of parts for cars, aircraft, and spacecraft. Chloroform used to be a widely employed anesthetic before safer and more effective inhalation anesthesia agents were developed.

- CCl4: was broadly used as a solvent in organic synthesis, as a powerful fire suppressant, and for dry cleaning. The high toxicity of carbon tetrachloride to the liver, however, has seriously limited its applications in the modern world.

4.2.8. Exercises.

1. The experimentally established tetrahedral structure of methane proves the existence of sp3 orbital hybridization. True or false? Answer

2. C-H bonds of alkanes are (a) ionic and highly reactive; (b) polar covalent and highly reactive; (c) nonpolar covalent and highly reactive; (d) ionic and highly inert; (e) polar covalent and highly inert; (f) nonpolar covalent and highly inert. Answer

3. Using the hybridization concept, explain why the geometry of an sp3-hybridized C atom is tetrahedral. [Answer: See 4.2.2]

4. Draw structural formulas for: 2-methylhexane; 3,3-dimethylpentane; 3-methyl-5-ethyldecane; 2,2,4-trimethylpentane, methylcyclohexane, ethylcyclopentane.

5. Name the following alkanes:

6. What is wrong with the following chemical names? (a) 2-ethyl-3,3-dimethylhexane; (b) 3,3-dimethylbutane; (c) 3-ethyl-5,5,5-trimethyloctane? Draw the structures and give correct names. Answer
7. Is cyclopentane a homologue or isomer of pentane? Answer

8. Write balanced chemical equations for combustion of (a) ethane and (b) pentane. Answer

9. Methane reacts with Br2 just like with Cl2. Write the mechanism of light-induced free-radical bromination of methane. How many products can be formed in this reaction? Why is radical bromination (and chlorination) of methane (and other alkanes) called a chain reaction? Answer