An essay at explaining the wide field of organic chemistry and its reactions
Organic Chemistry is the study of the compounds of carbon, their reactions and their preparation. It is a fundamental course for the life sciences and chemical engineering. Traditionally, it is presented as a description of a large number of compounds classified by structure, and a few hundred typical reactions that will show how they can be prepared, and how they behave with respect to one another. This constitutes a very large amount of empirical information that is usually committed to memory as a sort of mental handbook, supplemented by actual handbooks, and by laboratory practice with a relatively small number of important or illustrative general reactions. It is unusually difficult to assimilate this mass of information because of the lack of a general pattern (only empirical rules) and heuristic aids. Nevertheless, organic chemists become quite skilled in the use of this gigantic amount of information, and can work wonders in the laboratory with processes of which they have no fundamental understanding, but a wonderful empirical mastery. In fact, a fundamental understanding, were it generally possible, would be of little help to them in their practical endeavors, even if it could be brought to a more functional level.
However, a fundamental understanding is intellectually much more satisfying than a collection of empirical or semi-empirical rules. Just what goes on in a chemical reaction is fully known in principle, but predicting what will happen in any particular case is not possible, for more reasons than the difficulties of numerical calculation. Digital computers have brought much computing power to bear, but even this is much too feeble for the complexities of organic chemistry. Even if the results for a particular case were known, they would be only numbers giving no understanding: each problem would be unique. This might be technologically useful, but very boring from an intellectual point of view.
The great variety of organic compounds, and the enormous amount of investigation that has been done on them, can be appreciated in no better way than by a look at Beilstein's handbook of organic chemistry (see References). The fourth edition, which appeared in the 1920's, covers work up to 1910, and later work is quoted in supplements. Several suggestions are made in this article about compounds to look up, and the exercise is highly recommended. Proper use of Beilstein requires a knowledge of German, which was once a sufficient reason for all chemistry postgraduates to learn to read the language. Similar references exist in English, but they are not of the standard of Beilstein. Incidentally, a Handbuch in German was not one you could hold in your hand (all of Beilstein now seems to weigh about a ton), but one that is to be kept "at hand" for reference. Among the sciences, chemistry uniquely requires a reference like Beilstein.
The qualifier "organic" comes from the times when all organic chemicals were thought to be produced by organisms. In 1828, Wöhler made urea, H2N-CO-NH2, by heating its isomer ammonium cyanate, NH4OCN, showing that "life force" was not necessary for making "organic" compounds. Etymologically, the name comes from Greek organon, a "tool," via "organisms," which are collections of tools (organs). The Greek adjective organikos means "serving as an instrument," which is not what the "organic" in organic chemistry means.
Carbon alone gives us the giant molecules of diamond and graphite, the hollow balls of buckminsterfullerenes, and nanotubes. For information on these subjects, and all the lore of carbon, including oxides, gems and fuels, please see Carbon.
Hydrogen and oxygen play important roles in organic chemistry, in addition to carbon. Hydrocarbons are composed of carbon and hydrogen only, while carbohydrates join carbon with hydrogen and oxygen in the atomic ratio 2:1 as in water. These compounds include fuels and important biological substances such as sugars, cellulose and starch. Many additional compounds involve only C, H and O. With the addition of nitrogen, phosphorus and sulphur, most of the important biological substances are included, such as the amino acids, proteins, and many others. I will not consider the compounds of biological origin here for lack of space. The halogens, F, Cl, Br and I, are important in many reactions, including those of synthesis. These 10 atomic species: C, H, O, N, P, S, F, Cl, Br and I, are enough to explain nearly all of organic chemistry. Petroleum is a major source of aliphatic hydrocarbons (those not involving the benzene ring) and coal is a major source of aromatic hydrocarbons (those which do). From these sources most of the organic compounds of non-biological importance can be synthesized.
The formula expresses the atomic constitution of a substance. The empirical formula merely states the ratio of the atoms, as CH2O for acetic acid. This says that the ratio of carbon, hydrogen and oxygen atoms is as 1:2:1 in the molecule. From the atomic weights, the proportions by weight are found to be 12/30, 2/30 and 16/30 or 6:1:8. More useful is the molecular formula, which gives the numbers of each atom in one molecule. For acetic acid, this is C2H4O2. The molecular weight is 2x12 + 4x1 + 2x16 = 60. 60 grams of acetic acid contains 6.023 x 1023 molecules. The structural formula suggests the way in which the atoms are connected. This is CH3COOH in the present case. The carbon of the methyl radical -CH3 is connected to the carbon of the carboxyl radical -COOH by a single bond. We are supposed to recognize these structural elements from these abbreviations. Finally, the graphic formula explicitly shows the bonds between atoms, but does not usually show the spatial arrangement of the atoms. A stereographic formula shows the arrangement of the atoms in space as a perspective drawing, a stereopair, or an actual model. Examples of all kinds of formulas except the empirical will be found in this article. Familiarity with formulas and chemical equations is a basic skill.
Two substances with the same molecular formula but different structures are called isomers. Isomers are distinct substances with different properties. Two substances with the same chemical properties but different structures are called metamers, and are rare enough to be a curiosity. Two substances with structures that are mirror images of each other are called stereoisomers. This can happen, for example, when four different groups are bonded to a carbon. An interchange of two groups results in a molecule that cannot be superimposed on the original, but is a mirror image of it. Such molecules are called chiral, or "handed," and exist in different forms. A molecule with more than one chiral carbon can exhibit multiple stereoisomers. Stereoisomers usually have identical chemical properties, however. Molecules differing only in conformation are the same molecule, except that there may be rotation about certain bonds. Proteins, for example, are very flexible, and their conformation is important to their function.
Molecular models are a powerful aid to visualization, especially since they can be handled and moved to show the possibilities of molecular configurations. The best ones are correctly scaled with accurate angles and bond distances, the atoms represented by (truncated) spheres that show the spatial extent of electron density, and colored to show the different species. I assume such model sets still exist (a search of the Fisher catalog did not produce any photos) but are probably expensive. Photographs of such models are included on this page. There are many less expensive alternatives that may give a degree of satisfaction, generally using sticks and balls. The stereopair for free fusion is an excellent, but little-used, technique of illustration that would be very useful for organic chemistry. Pauling presents several examples in his general chemistry text, if you would like to see what this technique is like. I have developed computer graphic programs that produce stereo views, either stereopairs or red-green anaglyphs, but at this time it is not possible to give examples on this page. It is important to have a good appreciation of the three-dimensional structure of molecules to understand their properties and reactions.
Examples of structural formulas are shown at the left. The full structural formula shows each bond. Double bonds are represented by two lines, and triple bonds by three. The formula is usually abbreviated, as shown in the middle. The arrangement of the hydrogens is evident in H3C or CH3, and it is clear that the bond is with the carbon, whether the H3 is written before or after the C. The carboxyl radical is clearly meant by COOH, with a double C=O bond followed by -OH. Sometimes double bonds are shown for clarity even in an abbreviated formula. At the right is the Lewis structural formula, where the dots represent valence electrons. Each C contributes 4 dots, each H one dot, and each O six dots. The dots are arranged in a stable molecule so that each atom kernel, represented by its chemical symbol, is surrounded by eight dots. The kernel is the nucleus together with all the electrons out to the last completed noble gas shell of 2 or 8 electrons. The remaining electrons are the valence electrons, equal in number to the charge of the kernel in electronic units.
The carbon kernel consists of the nucleus of charge +6 and the two 1s electrons, so its charge is +4. The oxygen kernel consists of the nucleus of charge +8 and the two 1s electrons, so its charge is +6. The hydrogen kernel consists of the proton of charge +1. The formal charge on each atom is found by starting with the charge of the kernel and subtracting one for each electron that "belongs" to the atom. This includes one of the two electrons in each bond, and any valence electrons that are not bonded, called lone pairs. If this is done for each atom in the Lewis formula, we find that the formal charge of each atom is zero. The usual signed valence of inorganic chemistry (H = +1, O = -2) that works for ionic compounds means nothing here.
An ion is an atom with fewer or more electrons than necessary for electrical neutrality. We can extend the concept to refer to an assembly of atoms that is not electrically neutral, which we'll call a radical. For example, OH- is the hydroxyl radical. It has a kernel of charge +7 and 8 electrons, 6 from O, 1 from H, and 1 to give it a negative charge. Write the Lewis formula for the hydroxyl radical. The oxygen will be surrounded by 8 electrons, the proton by 2, so both will be happy and the radical will be stable. If a proton comes up, it can stick easily to one of the three electron pairs that are free, and make H2O, a well-known stable molecule. There are still two free pairs, and another proton could stick to one of them quite easily, forming H3O+, the hydrated proton. This is certainly what happens to any odd protons floating about. A fourth proton for the remaining free pair would give an excessive charge for such a small molecule, which is not stable considering the alternatives, so H4O++ is hypothetical. In principle, though, it would work. Any free pair is a target for a proton, though unattractive to a neutral hydrogen atom.
A free radical is a neutral atom or molecule with an "unsatisfied valence," or an odd unpaired electron that could engage in some arrangment of lower energy when encountering another molecule. A simple example is the H atom, a proton and an electron. Should it encounter another free radical, combination is certain, because in that way the total energy can be lowered. If it finds an atom to which it can bind more strongly than the atom's current engagement, it can steal the atom away, leaving another free radical in its wake.
Inorganic chemistry is characterized by ionic reactions in water solution. If we dissolve silver nitrate in water, we get solvated Ag+ and NO3- ions in solution. Now if we add salt, Na+ and Cl- are added to the solution. It happens that if silver and chlorine ions meet, they form AgCl molecules that ionize very little and are insoluble in water because they clump together. The silver chloride precipitates until either silver or chlorine runs out. Similarly, if a gas is evolved a reaction will go to completion. Usually, however, we have an equilibrium between reactants and products with both present at the same time. Inorganic reactions also occur in the solid state, but the nature of these reactions is much less well-known.
Because organic compounds usually do not ionize, or even dissolve in water, they seem unreactive. This, however, is not actually so. Gasoline burns quickly with the release of a great deal of heat in internal combustion engines, for example. Instead of an ionic reaction, we are dealing with a free radical reaction in this case. Suppose methane and chlorine are mixed in a vessel. In the absence of light or a catalyst, there is no reaction although one would be expected on energy grounds. These reagents would love to rearrange themselves as methyl chloride and hydrogen chloride. When the vessel is exposed to light, the reaction is quick, forming mainly CH3Cl until either the chlorine or the methane is used up (other chloromethanes are formed in smaller amounts). The mechanism of this reaction is: first the formation of free chlorine radicals by photochemical homolysis: Cl2 → 2Cl., where the . represents an unpaired electron. Now, CH4 + Cl. → HCl + H3C., so a free methyl radical is produced when the chlorine grabs a hydrogen. Then, H3C. + Cl2 → CH3Cl + Cl. and the chlorine free radical is recovered. The C-H bond is only a little less stable than the H-Cl bond, but still the chlorine has a good chance of stealing the hydrogen. Every photolysis, the splitting of Cl2 by light, is the start of a chain reaction that ends only when two free radicals encounter each other and combine. If the heat produced in the reaction is enough to create free radicals by itself, the chain reaction may branch by thermolysis and proceed even more rapidly. Free radical chain reactions are almost always involved in rapid reactions, and are much more common that might be supposed.
Returning to equilibria, we might review a few points that are treated in general chemistry. The ionization of acetic acid is an example of a chemical equilibrium. In water, CH3COOH ↔ CH3COO- + H+. The equilibrium between the three species is expressed by [Ac-][H+]/[HAc] = 1.753 x 10-5 = K, where the [HAc] is the concentration of acetic acid in moles per litre, and so forth. In 1M acetic acid (60g per litre, about 6%, like household vinegar), the hydrogen ion concentration x will be the solution of x2 = K, since the hydrogen ion and acetate ion concentrations will be equal in this case. Then, x = 4.3 x 10-3, or pH = 2.4, since pH = -log [H+]. Also, [H+][OH-] = 1 x 10-14 in aqueous solution, where the concentration of water is about constant (55.6 mol/litre) and is figured into the equilibrium constant. The pH of water is 7. In water, either the hydrogen ion concentration can be large, or the hydroxyl ion concentration can be large, but not both at the same time. These ions often have an important effect on reactions, and even on the structure of dissolved molecules. If you are rusty, it would be a good idea to review equilibrium and solutions.
Organic chemistry (and all other chemistry at that) is completely described by the quantum mechanics of nuclei and electrons interacting by electrostatic forces. The small contributions of the magnetic interactions of orbital currents and spin magnetic moments is of interest in spectroscopy, but of no consequence in chemistry. Charges of like sign repel one another, so electrons are repelled by electrons and nuclei are repelled by nuclei, but electrons and nuclei attract one another strongly and at large distances. The classical forces are F = z1z2(e2/r2), where F is the force in dynes, e the magnitude of the electronic charge (4.803 x 10-10) in esu, r the distance between the charges z1e and z2e. For electrons, z = -1.
The description of matter by means of classical mechanics, with point particles interacting by the electrostatic force just given, in accordance with Newton's laws of motion, is impossible. Classically, our world cannot exist. Quantum mechanics must be used instead, and it gives a full and accurate description of matter, on the level of chemistry. It is a good enough approximation to consider the heavy nuclei as occupying definite positions, and to provide an electric field in which the electrons move, subject to their mutual repulsion. The motion of the nuclei can be described classically, but the electrons are a different matter. The electrons occupy most of the volume in a molecule. The volume that an electron occupies depends on the attractive electric field. A nucleus draws the electrons neutralizing its charge closely around it, in a sphere of about 0.1 nm radius. All atoms are about the same size, since z electrons are attracted by a nuclear charge of ze, so that more electrons means stronger attraction, so the electrons huddle more closely.
The motion of the electrons is described by a wave function that is a solution of the Schrödinger equation with the electrostatic potentials of nuclear attraction and electronic repulsion in the Hamiltonian. In this way, an electron "feels" all of the region in which it moves. Generally, we take as a first approximation a single electron moving in the average field of all the other charged particles. These approximations are necessary for finding solutions in a practical case, but are not fundamental. In a particular stationary state the electron is described by a wave function called an orbital, to which corresponds an energy. That is, an electron "occupying" the state has the energy corresponding to the orbital. In any problem, we can find (in principle) many states, with energies from some minimum upwards. Then we "assign" electrons to the states to determine the overall state of the atom or molecule. The states themselves are dependent on which other states are occupied, and which are not, so we must adopt some kind of iterative approach (called the self-consistent field).
An atom or molecule is usually electrically neutral: it contains equal amounts of negative and positive charge. If this is not true, very strong forces are exerted on other charged atoms or molecules and something vigorous occurs. When salt is dissolved in water, Na+ and Cl- ions are formed in large numbers. If they were not very well mixed, extraordinary things would happen. Without the Cl-, the Na+ ions would instantly rush to the limits of the solution, like a charge on a metal sphere, and a powerful spark would attract electrons from the air. Not only are the ions well mixed, but each is surrounded by an atmosphere of polar water molecules, the positive hydrogens clustered around a chlorine ion, the negative unbonded electrons of the oxygens clustered around a sodium ion. These micelles spread the ionic charge over a larger volume, and the ionic charges are distributed to screen one another well. An applied electrostatic field will cause the ions to drift in the appropriate directions, the sodium to the cathode (-), the chlorine to the anode (+). We also note that the sodium ions do not behave anything like sodium metal, and the chlorine ions do not behave anything like chlorine gas.
The energy of an isolated system is constant. However, if a collection of nuclei and electrons can rearrange themselves to give a lower energy, and the excess energy can be carried away somehow, we have a chemical reaction. In most practical cases, systems are in thermal contact, so that the criterion for stability is a minimum of the free energy a = u - Ts + Pv, and not just a minimum of the internal energy u. Let's simply talk of u instead here, letting it refer to whatever thermodynamic potential that is appropriate. Therefore, if we bring two systems into contact, they may rearrange themselves to form a new system of lower energy, dissipating the excess energy by collision or by breaking into two or more parts with kinetic energy. This is the only reason for a chemical reaction. All binding energy in molecules is electrostatic. A molecule is stable if its energy is less than that of any other configuration accessible to it. When different molecules come together, there may be rearrangements of smaller energy, and a reaction may occur.
For a reaction to occur, the reactants must be brought together physically, and in such an orientation that the reaction may proceed, rather than repulsion. This means, generally, that there are only two reactants (unless by some means we can bring more than two reactants to bear at the same time) and they must be properly oriented so that they attract. Some of the kinetic energy of the molecules may be used to attain a favorable configuration. Two free hydrogen atoms, for example, are each electrically neutral and quite happy. They exert no forces on each other until they approach closely, so that their electrons overlap. With the nuclei at a certain distance apart, the Schrödinger equation gives two orbitals, one of low energy (the bonding orbital) where the two electrons are between the nuclei, and one of high energy (the antibonding orbital) where the two electrons are on the opposite sides of the nuclei. One electron of the two available put in the bonding orbital gives a stable system, H2+. We can put the second electron in the same orbital, provided its spin is opposite to the spin of the electron already there. Since the spin angular momenta are opposite, the net spin angular momentum is 0 (a singlet state). This is the stable H2 molecule. If we happened to have a third electron, it could go nowhere but in the antibonding orbital, so the H2- ion would have about the binding energy of H2+. H2 in the triplet state, where the electron spins are in the same direction, would not be bound at all. To a (good) approximation, the spins have only a bookkeeping significance.
These things are illustrated at the right in an energy-level diagram. The energy ε is about 51 kcal/mol. The antibonding state is assumed to be as high above the zero level as the bonding state is below it. The Pauli Principle is that no two electrons may occupy the same state, since the wave function must change sign if the coordinates of any two electrons are interchanged. The "state" is assumed to include the spin. Electrons are exactly the same, and can be identified only by the quantum numbers specifying a state. An organic chemistry textbook can be happy saying in one place that "electrons tend to pair," and in another that "electrons repel." The statement that "electrons tend to pair" is not only wrong, it is misleading. The two electrons in H2 are in the same state because all the other possibilities give higher energy. They actually would rather not be so close together, and their mutual repulsion decreases the binding energy somewhat. They form a pair out of necessity, not out of desire for one another. When electrons are being added to the 2p levels in B, C and N the spins are parallel in the ground state. In O, F and Ne they are forced to double up.
An equally misleading statement is that "electrons tend to form octets." The only reason for this is that when the eight 2s and 2p levels are filled, it is quite a jump in energy to the next available state, 3s. When the 3p levels are filled, the next are the 4s, not the 3d, and after the 4p levels, the next are the 5s. It just happens that at Ne, A, Kr, Xe and Rn, which have just completed s2p6 "shells" of 8 electrons, there is an energy gap to the next s level. No energy advantage can be gained by promoting an electron from the filled shell in most cases, so a kernel with a noble-gas structure generally remains untouched. The charge on the kernel is equal and opposite to the charge on the additional valence electrons. Electrons have no "tendency" to form groups of 8; it just works out that way. Of course, this is just a manner of speaking, but it should not be allowed to hide the real reason. There is a search for lowest energy, not for the figure 8. These "rules" are typical of organic chemistry, ascribing nonexistent wills and desires to electrons and atoms as explanations of physical phenomena. The tetrahedral angle of carbon bonding does not result from any love of carbon for 109° 28', but simply from the fact that this angle puts equivalent bonds as far from each other as possible. When conditions are not the same at the ends of the bonds, the angles are not exactly tetrahedral.
In H2, the pair of bonding electrons occupy the bonding orbital, staying as far away from each other as they can in the limited volume. They are mainly between the two protons, screening them from each other, and symmetrically disposed relative to the horizontal plane of symmetry. The molecule is quadrupolar, with positive charges on the ends and a negative charge in the middle. Its electrical field decreases rapidly with distance from the molecule. Both the molecule and its bond are described as nonpolar, and the bond is called covalent. Suppose that one of the protons is replaced by a chlorine atom. This would affect the bonding orbital by pulling the electrons closer to the increased charge, and away from the other proton. The chlorine atom would "like to add one electron to its octet" in the usual parlance, meaning only that this would be energetically favorable. Now one end of the molecule would have an increased positive charge, the other end a negative charge of equal and opposite amount. This is a dipole, whose external field is stronger and of longer range than a quadrupole field. The molecule would be polar, and the bond has no good name in organic chemistry, except to call it a polar covalent bond.
The bonds in many organic compounds can be considered independent of one another, so that the total bond energy, the heat of formation, is approximately the sum of the individual bond energies. A table of bond energies is shown at the left. In many cases, the energy of a bond is affected by neighboring bonds. We will see that the energy of a C-H bond is affected by how many other carbons are bonded to the carbon (more carbons makes the bond weaker), and that the C-OH bond in the carboxyl radical COOH is affected by the presence of the C=O nearby. Also, energies depending on the state of the substance are not included. The heat of evaporation of water, for example, is about 9.7 kcal/mol, and the heat of fusion is about 1.4 kcal/mol. These rather unusually large figures show the limits of these effects. The energies of bonds not in the table can be estimated by comparison, or can be found from tabulated heats of formation in handbooks.
If we imagine an electron taken from one hydrogen atom and transferred to the chlorine, what results would be the ions H+ and Cl- that would attract one another electrostatically. In HCl gas, this does not quite happen, but HCl dissolves in water to form the ions, each surrounded by water molecules, or solvated. In salt, NaCl, the electron in the bond is similarly attracted to the Cl, or would be if this molecule existed. In solution, we have sodium and chlorine ions, as in the case of HCl, and if the water is evaporated, we have a crystal of salt, which contains Na+ and Cl- ions in a cubic lattice. The salt crystal is the state of least energy, and is one huge molecule. Chemists speak of an ionic bond in this case, though there is really no bond at all, unless it is a polar covalent bond or an expression for the crystal binding. As best I can tell, an ionic bond is one that breaks up into ions in solution or a crystal. It is purely a matter of degree.
The covalent radii of atoms making single bonds are shown at the right. 10Å = 1 nm. The C-H bond is 0.77 + 0.29 = 1.06Å in length. The C-C bond is 2 x 0.77 = 1.54Å long. The results are approximate, of course, but the marvel is that these values are relatively consistent. The C=C bond is 1.33Å, and the C≡C bond is 1.20Å. The C=O bond is 1.24Å, and the N=O bond 1.18Å. Multiple bonds are shorter than single bonds. By means of the tables just given, the sizes and heats of formation of organic compounds can be estimated. A good set of molecular models also makes clear the spatial configurations.
Three possible structures can be drawn for the carbonate ion, CO3--. The kernels have charge 22, and there are 24 valence electrons. In each structure, there is a double bond with one oxygen, and single bonds to the others. The state of the ion is a linear combination of equal amounts of the three structures shown. The result is that the three oxygens are equivalent, and located at the corners of an equilateral triangle. The C-O bond length is 1.31Å, intermediate between the C-O and the C=O lengths. The ion is more stable than any one of the contributing structures, the result of having the freedom of making a linear combination of them to construct the actual state. In chemistry, this is called resonance, although there is no actual resonance involved. The name is derived from an early semiclassical model. Protons can adhere to any of the free pairs, two protons making the molecule neutral, which is carbonic acid, H2CO3. If the two protons carry away an oxygen, the result is CO2 and a molecule of water.
Exactly the same thing occurs with the nitrate ion, NO3-, but here the kernel charge is +23 and there are 24 electrons, as before. Three possible structures of reasonable stability can be drawn as before, and the actual state is the linear combination of them. The nitrate ion is triangular, stable, and with an N-O bond length of 1.22Å.
The carboxyl group -COOH ionizes to -COO- and H+ in water, as we have already seen with acetic acid. Two structures are possible for the ionized group, shown at the left. The actual state is a linear combination with equal weights, making the two oxygens equivalent. A bond picture is shown in the diagram; the reader may want to construct the Lewis structures. Resonance in these structures makes the acetate ion more stable, and increases the probability that the -COOH will ionize. In chloroacetic acid, CH2ClCOOH, the chlorine pulls electrons towards it, and away from the oxygen in the -OH group. This makes it easier for water to pull off the hydrogen, so chloroacetic acid is a stronger acid than acetic, with K = 1.396 x 10-3. Dichloroacetic acid is still stronger, at K = 5 x 10-2, and trichloroacetic acid, K = 0.13, is as strong as sulphuric acid.
Another form of bonding is shown at the left for SO2. The kernel charge is +18, and there are 18 electrons, so the molecule is neutral. The 18 electrons can be distributed in two ways so that each atom has a completed octet. In each way, one of the oxygens is bonded by a pair of electrons donated by the sulphur. The two structures resonate so that the oxygens are equivalent and the molecule is symmetric. Note that each Lewis structure has a formal charge of +1/2 on the S, and -1/2 on the O. In the actual molecule, the S will have a formal charge of +1/2, and each oxygen a formal charge of -1/4. The oxygens pull the electrons towards themselves, leaving the sulphur slightly positive. This type of bond is called a "semi-polar double bond," or a "coordinate covalent bond," neither name being particularly descriptive or fortunate. Although the electrons "donated by the sulphur" are distinguished in the figure, electrons arrange themselves for the lowest energy, and cannot remember to what atoms they belong.
Carbon atoms can join to each other to form chains and rings of any desired complexity. A few examples are shown at the right. In them, carbon makes four bonds at tetrahedral angles (109° 28'). We assume that all the bonds not shown in the figure are with hydrogen atoms, so that all these molecules are hydrocarbons. Two major classes are the molecules without rings, which form tree-like branching molecules, and those with rings. Of course, the two classes may be combined in any way. Only the simpler structures are commonly found, however. The compounds are named according to the number of carbon atoms in them. The branched molecules can also be named by the IUC system, in which the longest chain of carbons gives the basic name, and the side chains are considered as substituents. The carbons of the longest chain are given serial numbers so that the smallest numbers appear in the name. For example, 2-methylpropane has a methyl group (CH3) attached to carbon 2 of the propane chain.
All compounds like those shown in the figure are called paraffins or alkanes. The acyclic paraffins have the molecular formula CnH2n+2 (no matter how they are branched), and the cycloparaffins, CnH2n. A series of compounds differing only by CH2 groups is called a homologous series. The first three cyclic paraffins are planar, and the bond angles are considerably changed from the tetrahedral value, which makes them less stable than if they had tetrahedral angles. The angle 108° is shown for cyclopentane, but the molecule is actually not planar and the actual angles are close to the tetrahedral value. Cyclohexane has the tetrahedral bond angles, and is not planar. It exists in two conformations, called boat and chair, depending on whether two ends are both tipped in the same direction, or opposite directions. The cycloparaffins are not very important compounds, though interesting ones.
Carbons with three hydrogens attached are called primary, with two hydrogens secondary, with one hydrogen tertiary, and when attached only to carbon, quaternary. Primary hydrogens are more strongly bound than secondary, and secondary more than tertiary. Free radicals attack tertiary hydrogen most effectively, secondary hydrogens less effectively, and primary hydrogens least effectively. Straight-chain hydrocarbons without branches are called normal. Normal butane, for example, may be written n-butane. Usually, if the "n" is omitted, the paraffin is considered to be normal. There is only one structure for methane, ethane and propane, but all higher paraffins have isomers with the same molecular formula but different structures. Pentane has 3 isomers, but dodecane (12 carbons) already has 355 isomers.
The paraffin chains are not actually straight lines, but zigzag ones as each carbon maintains tetrahedral bond angles. They can rotate freely about any bond, so they represent more a plate of spaghetti than a pile of rods. All the paraffins have about the same chemical properties. They do not react with any of the common reagents, which gave them the name "paraffin" from Latin parum, "slight" and affinis, "affinity." They do react readily with free radicals, which displace the hydrogens, notably in the reaction of burning. They are colorless and odorless. They dissolve in organic (nonpolar) solvents, but not in water. The melting and boiling temperatures increase with the number of carbon atoms, and are lower for the more compact molecules that do not get entangled so easily. At room temperature, methane to butane are gases, pentane to nonadecane (19 carbons) liquids, and eicosane and above solids. Eicosane, with 20 carbons, is a soft gel at room temperature, typical of petroleum jelly. Cetane, or n-hexadecane, with 16 carbons, melts at just below room temperature and boils at 287.5°C. Examples of paraffins are easy to find in daily life: natural gas (methane), gasoline, kerosene, lubricating oils, petroleum jelly and wax.
The primary source of paraffins is petroleum. Some high-quality crude oils consist almost entirely of acyclic paraffins. They can be separated by fractional distillation into fractions with small ranges of numbers of carbons. The bubble tower fractionating column was introduced in the 1920's, replacing the simple still and condenser. Straight fractionation yields naphtha (gasoline, boiling 40°C-180°C, 5-10 carbons), distillate (kerosene, boiling 180°C-230°C, 10-13 carbons), gas oil (fuel oil, boiling 230°C-400°C, 13-25 carbons), heavy oils (lubricating oils, boiling 400°C-500°C, 20-30 carbons) and bottoms (wax, tar and asphalt, more than 30 carbons). The predominant demand for motor fuels prompted a search for ways of increasing the amount of gasoline produced from a barrel of crude, which was originally only about 25% (it was then a waste product). In 1913, cracking of the heavier fractions by catalysis and heat was introduced, which doubled the gasoline yield. This scrambling of the structures introduced more branched paraffins, which burn more smoothly in an internal-combustion engine than straight chains. This alone boosted the octane rating from about 50 to over 70. Catalytic polymerization or reforming (catalytic cyclodehydrogenation), together with hydrogenation, adds additional compounds with high anti-knock value, such as "iso-octane." Isooctane is 2,2,4-trimethylpentane. The octane number of a fuel is the percentage of isooctane in a mixture with n-heptane that has the same anti-knock properties. Isooctane, then, represents a 100-octane fuel, n-heptane a 0 octane fuel. Isooctane melts at -107.4°C and boils at 99.2°C, while n-heptane melts at -90.6°C and boils at 98.4°C. These figures are typical of "gasoline," which must be liquid and volatile at ordinary temperatures.
Petroleum is the only cheap natural source of paraffin hydrocarbons, supplying ethylene and propylene, the raw material for plastics, alcohol, acetylene, phenol, detergents and other basic industrial chemicals. It is a waste to squander petroleum as a fuel, but because of its present cheapness this is its principal use. If anything, the parallel misuse of natural gas is even worse. When petroleum becomes scarce and expensive (and it will, not far in the future and quite suddenly), the effect on the chemical industry will be far worse than the effect on fuel supplies. There are alternative fuels, but no alternative cheap source of paraffins.
Two carbon atoms can be joined by double or triple bonds, and this gives further prospects for molecules. Some important examples are shown at the left. A compound that contains double or triple bonds between carbons is called unsaturated, while those with only single bonds are saturated. Just as saturated compounds are called alkanes, those with double bonds are alkenes and those with triple bonds alkynes. Aliphatic compounds with multiple bonds between carbons are, in general, called olefins. The most important are the alkene ethylene and the alkyne acetylene, which are important starting points for making other organic compounds because of the activity of the double bonds, which also make polymerization possible. The thermoplastics we see every day are an example. Ethylene polymerizes into a saturated compound called polyethylene, or PE. This is made in two forms, high-density HDPE and low-density LDPE. Propylene polymerizes to polypropylene, PP, which is stronger than PE. If one hydrogen is replaced by a chlorine to form ethylene chloride, more commonly called vinyl chloride, the result is the saturated polymer polyvinyl chloride, or PVC. If a benzene ring is used instead of chlorine, the result is polystyrene, PS. Plastic articles usually carry a recycling symbol indicating the polymer used by a number in a triangle, plus the symbols given above. HDPE is 2, PVC is 3, LDPE is 4, PP is 5 and PS is 6.
The addition of oxygen to our toolbox gives further possibilities. We can base a series of compounds on a fundamental compound, such as water, HOH. If an H is replaced by any paraffin, denoted by R, we get ROH, which is an alcohol. The simplest alcohol is CH3OH, methanol. We can make alcohols by sticking one or more OH, hydroxyl, groups on any of the paraffins. (CH2OH)2 (1,2 dihydroxyethane) is glycol, famous from car antifreezes. (CH2OH)(CHOH)(CH2OH) (1,2,3-trihydroxypropane) is glycerol or glycerine. Mannitol has 6 hydroxys, with 3 more CHOH groups added to glycerol. The OH makes these molecules soluble in water if the R part is small enough. The H can even come off as H+, and ionic reactions can occur. The essential thing here is the -OH group, which is called the functional group of the alcohols, and is responsible for their particular properties. An alcohol never lets go of the O, however.
If both of the hydrogens in HOH are replaced by paraffin groups, which may be the same or different, the result is an ether. The simplest ether is CH3OCH3, methyl ether, a gas. Ethyl ether, (C2H5)2O, a volatile liquid, is the familiar ether that was used as a general anesthetic, and is a useful solvent. It boils at 34.6°C and is very inflammable. On long standing, or if the bottle is exposed to sunlight, explosive peroxides can form. The vapor explodes in concentrations of 1.85% to 36.5%, and the ignition temperature is only 180°C. Ether is more hazardous than gasoline.
The convenient abbreviations Me = CH3 for the methyl group, and Et = C2H5 for the ethyl group, will often be used in what follows.
Alcohols and ethers are both derived from water, HOH, but they act very differently. In alcohols, the OH group may donate a hydrogen: ROH → RO- + H+. Then it is a Brönsted-Lowry acid, or proton donor, and its anion RO- a base, or proton acceptor. MeOH is methanol, and MeO- is the methoxide anion, a strong base. The lone pairs of electrons on the O may attract the H from another molecule, forming a hydrogen bond. The hydrogen bond is what gives water its peculiar properties, such as high melting and boiling temperatures. The lighter alcohols share this peculiarity: even MeOH is a liquid. Higher alcohols become less and less like water as their alkyl parts dominate.
None of this can happen with an ether, ROR', so there is no acid-base behavior, and no hydrogen bonding, even with the lightest memeber, MeOMe. Indeed, methyl ether is a gas, and ethyl ether a very volatile liquid. With no hydrogen, there is no easy point of chemical attack in an ether, and they are as inert to ionic reactions as paraffins.
Carbon can form a double bond with oxygen, to make the carbonyl group =C=O, in which the carbon has two free valences. The basic compound is HCHO, formaldehyde, where the two free valences are satisfied by hydrogens. Then we can add a hydrocarbon R to make the series of compounds RCHO, which are called aldehydes. A second hydrocarbon R', that may be different from R or the same, then gives the compounds R(CO)R', called ketones. The reason the two kinds of compounds are distinguished is that the CHO group of an aldehyde behaves somewhat differently than the CO group of a ketone, just as the OH of an alcohol behaves differently from the O of an ether. Further examples are CH3CHO, acetaldehyde (a gas), and CH3COCH3, acetone or dimethyl ketone (a liquid). Methyl ethyl ketone, called MEK, is a popular solvent, as is acetone.
The carbonyl and hydroxyl groups can be combined into the carboxyl group, (C=O)OH or COOH. The two groups modify each other's behavior considerably. The H comes off more easily than in an alcohol, leaving a COO- ion in which the oxygens are equivalent (a case of resonance, mentioned above, which is the reason for the behavior), and a H+ ion. They are stronger acids than the alcohols. Conversely, their bases, such as MeCOO-, the acetate ion, are weaker. They will form salts with cations, such as MeCOONa, sodium acetate. Compounds containing the COOH group are called organic acids, of which CH3COOH, acetic acid, is a familiar example. The simplest organic acid is HCOOH, formic acid (found in ants, formicae). HOOC-COOH is oxalic acid. HOOC-CH2-COOH is malonic acid.
Cyclohexane, C6H12, is a typical paraffin, containing only single bonds. We have mentioned above that it is not a planar molecule, and has two conformations. The very different compound C6H6, called benzene, was first found in the liquids obtained by distilling coal. In German, it is known as benzol. Benzine is an impure benzene that can be used as a motor fuel. See Beilstein, Bd. V, p 179. It is a colorless, volatile liquid with a characteristic odor and density 0.88. It is carcinogenic. Faraday studied benzene in 1825, but it was known to Schelenz in 1785. Benzene is found in some petroleums, in addition to coal tar. It has only half as much hydrogen, and is a planar molecule of high stability. The fewer hydrogens suggest that it is unsaturated, but it does not act like any of the unsaturated hydrocarbons. It is the basis of a large number of compounds, called aromatic compounds from the pleasant aroma of benzene and of many compounds containing the benzene ring, including many flavorings. The best traditional structural formula, the Kekulé formula, is shown at the upper left. It features alternating double and single bonds, called conjugated double bonds. (As usual, hydrogen is not shown explicitly.) A space-filling model is shown at the left. However, benzene does not behave like linear molecules with conjugated double bonds in any way. There are six electrons in addition to those required for the single bonds, and these six electrons are free to move in circles above and below the plane of the molecule. This mobility, similar to the "resonance" we have encountered above, allows a large binding energy, which makes benzene very stable. This is indicated by the hexagon at the right, where a carbon and hydrogen are understood at each vertex, and the circle represents the six ring electrons. A hexagonal skeleton is often used to represent a benzene ring.
Substituting a methyl group for a hydrogen gives toluene, also an aromatic liquid. Toluene is not carcinogenic and you can safely smell it. Its odor resembles that of benzene, but is not as pleasant. A second methyl group could go onto any one of three vertices to make distinct compounds. These vertices are denoted by ortho, meta and para, which are abbreviated o-, m- and p-. Xylene has two methyl groups, and three isomers, o-xylene, m-xylene and p-xylene. In the figure, m-xylene is shown. The functional groups OH and COOH can be added to the benzene ring, producing phenols and aromatic acids. Adding an OH to toluene gives a cresol. Phenol was discovered in 1834 by Runge, who named it carbolic acid. In 1841, Laurent obtained it in pure form, calling it acide phénique. In the same year, Gerhardt named it phenol. The OH in a phenol acts differently from the OH in an alcohol, since the attraction of the benzene ring for electrons makes it easier for phenol to ionize as an acid. In fact, it is known popularly as benzoic acid. The group BzO- resulting from this ionization is called the phenoxy group (Bz represents a benzene nucleus). There can be more than one OH group, as in resorcinol and hydroquinone. These can be named systematically as m-dihydroxybenzene and p-dihydroxybenzene. If the benzene ring is not adjacent to the OH group, the compound acts like an alcohol, and is called an aromatic alcohol. For example, Bz(CH2)OH, benzyl alcohol. The COOH groups behave similarly, making stronger acids than when attached to an alkane. Again, more than one group can appear on a benzene nucleus. Terephthalic acid has two COOH groups in the para position. In the meta position, the compound is called isophthalic acid, and in the ortho position, phthalic acid. Mellitic acid has six carboxyl groups, hemimellitic acid three.
Phenol is a colorless solid, melting at 40.9°C and boiling at 181.8°C, with a characteristic odor suggesting disinfection. It is an excellent antiseptic, but is also corrosive to tissues and poisonous, so it has largely been replaced by other compounds, some of them related to phenol. It is found in wood and coal smoke, and with formaldehyde gives the preserving power of smoking. The name comes from the Greek fainos, "shining," because it was found in the distillate from making illuminating gas. Benzol is not phenol, but the German name for benzene. Benzene itself has a similarly tortuous etymology, being originally from Arabic luban jawi, the "frankincense of Java," through "benjoin" after dropping the article lu-, then to the aromatic "gum benzoin" and finally to benzene. Gum benzoin contains cinnamic acid, which is ethylene with Bz on one carbon and COOH on the other.
The benzene ring can also act as the R groups in ethers, aldehydes and ketones. These are called aryl radicals, Ar, as the R were alkyl radicals. Benzaldehyde, BzCHO, a colorless liquid smelling like almonds, is an example. Adding an OH in the para position, and OMe (methoxy) in the meta position gives vanillin, a very popular flavoring. Benzene rings can also fuse, as in naphthalene (mothballs), anthracene and phenanthrene. The typical inertness of benzene disappears as more rings are fused. The limit is reached in graphite, which consists of sheets of fused benzene rings. Benzene, and some of the fused ring compounds, are powerful carcinogens, and contact with them should be avoided. Other aromatic compounds have no carcinogenic properties at all.
We can take a closer look at naphthalene, which forms many useful compounds. The carbons are not all equivalent, as in benzene. They are numbered as shown in the diagram. Carbons 1, 4, 5 and 8 are called α-carbons or 1-carbons, carbons 2, 3, 6 and 7 are called β-carbons or 2-carbons. The bonds of the other two carbons are all involved in the naphthalene core, so substitutions cannot take place at these locations. An -OH substitution can make two phenol-like molecules, α naphthol (1-naphthol) or β naphthol (2-naphthol), with closely similar but not identical properties. α-naphthol smells faintly of phenol, but β-naphthol is odorless. β-naphthol is triboluminescent. Both naphthols irritate the nose, have a burning taste, and crystallize in the monoclinic system. The melting points are 96°C and 122°C, respectively. Their densities are about 1.22. The melting point of naphthalene itself is 80.2°C, and it boils at 218°C; its density is 1.1517. From alcoholic solution, it crystallizes in colorless plates. Naphthalene has a strong characteristic odor, and is not carcinogenic. Naphthalene is not only found in coal tar, but also in some petroleums, such as that from Puente Hills, California.
We have now looked at many carbon compounds that can be made with C, H and O only, as well as at the functional groups OH, CO, and COOH. We have met the alkanes, alkenes, alkynes, and aromatics, and have become acquainted with alcohols, ethers, aldehydes, ketones and acids. These building blocks can be put together in limitless ways. We have not even considered the compounds of biological importance, such as sugars, amino acids, fats and oils, starch and cellulose. For the biological compounds, the reader is advised to refer to Stryer. We have not even considered the halogens: fluorine, chlorine, bromine and iodine, that can replace hydrogens in any of the compounds so far considered. Rings containing nitrogen (heterocyclics) have not been mentioned, nor the functional groups -NO2 (nitro), -NH2 (amino), -C≡N (cyanide or nitrile), =C=N- (imine), and many others. Compounds in which S replaces O are yet to be mentioned.
The reactions of inorganic chemistry usually take place in aqueous solution, where the water pulls apart the reagents into ions and allows them to move and interact freely. Occasionally reactions occur in fused reagents, between gases, or in the solid state where exchanges with the gaseous state can take place. The reactions are generally fast, and often go to completion. In organic chemistry, there are often other solvents than water, since most of the compounds are not soluble in water and do not form ions. Even where molecules are mobile, as in liquids and gases, reactions are often slow and incomplete. Even when a reaction is very exothermic, the reactants may not react. For a reaction to occur, the reagents must be able to try out various alternatives to see which ones will lead to a reduction in free energy, so that the reaction may proceed. A reaction must have a definite mechanism that permits it to proceed with reasonable speed. The reaction path may be reversible, so that the products can react to produce the reagents as well as the reagents can react to produce the products. This situation is well-known in chemistry, and is expressed by the law of mass action and equilibrium constants. An organic reaction may follow several paths, with different probabilities and different products. Often the particular products are a clue to the reaction mechanism.
In the rest of this section, some reaction mechanisms will be discussed from a theoretical point of view in an effort to find out what is going on. Organic chemistry has a very different practical side, however. The best conditions for a particular reaction have been empirically determined, and a definite procedure prescribed for it. An organic chemistry review book lists 258 named reactions. This "cookbook" approach is very practical and effective, but tends to obscure basic principles. The reader is referred to texts on organic chemistry for the practical details of reactions. One is struck by the fact that discussions of organic reactions involve practically no mathematics or fundamental theory, and instead are based on qualitative theories and empirical rules. In principle, every question could be answered by numerical calculation on the basis of quantum mechanics, but this procedure is impossibly complicated besides being unilluminating. The wonder is that the empirical approach of organic chemistry is so useful and illuminating. The reader is warned, however, that this is a very complex and confusing subject, with an appallingly large literature. I will do my best to reveal some order in it, but with no guarantee that the result will be either clear or correct. I rely wholly on my sources, which I shall interpret to the best of my ability.
If we drop sodium metal in water, the vigorous reaction evolves hydrogen gas and produces sodium and hydroxide ions: 2Na + 2HOH → 2Na+ + 2OH- + H2. If we drop sodium metal in EtOH, a milder reaction has this result: 2Na + 2EtOH → 2Na+ + 2EtO- + H2 (Note: Et = ethyl). The ethoxide ion, EtO- is produced. If we evaporate the remaining alcohol, we get crystals of EtONa, sodium ethoxide. When dropped in water, EtONa is hydrolyzed to make EtOH, Na+ and OH-. The solvent in any reaction using EtONa should be EtOH, not water. This reaction can be imagined for any alcohol, so we have available in principle any alkoxyl ion we desire.
We have already mentioned that a mixture of methane and chlorine gases does not react in the dark, though the reaction is strongly favored energetically. If a chlorine molecule is split by absorbing light, however, we now have free-radical Cl. (the dot represents an electron available for a bond) and a reaction mechanism for producing methyl chlorides. Probability makes MeCl the predominant product. Using this rapid chain reaction makes alkyl chlorides available to us, and alkyl chlorides are more subject to chemical attack than the alkanes. The C-H bond is tight and nonpolar; there is no place to gain an electrical hold on it. Only the free-radical attack is successful, since Cl. can tear the H away from the C to form HCl. Now we can make other compounds by substitution, where the Cl is replaced by some group of our choosing.
When one of the hydrogens of methane is replaced by chlorine to make chloromethane, the carbon-chlorine bond becomes slightly polar, the carbon aquiring a net charge +δ and the chlorine a net charge -δ This creates a dipole, with a rather long-range electrostatic field. A group with negative charge is attracted to the carbon, and this group could be OH-. It is not necessary for the attracted group to be negatively charged, since a neutral group could be polarized by the field and still attracted. It is necessary that this group be a strong base, eager to donate an unbonded pair of electrons to make a bond. Such a group that seeks out a positive charge is called nucleophilic. The number of unbonded pairs, and the number available for nucleophilic attack, is shown at the right for the atoms of the first period. The importance of nitrogen and oxygen is evident. When the nucleophile approaches the carbon, electrons are repelled by its charge onto the chlorine atom, which becomes even more negative and independent. In the final stage, the OH can bond to the carbon by means of a lone pair of electrons on the oxygen, while the chlorine drifts away taking a pair of electrons with it. The net reaction is OH- + CH3Cl → CH3OH + Cl-. This reaction is called nucleophilic substitution and denoted SN2. The S means substitution, the N nucleophilic, and the 2 means that the rate depends on the collision of two species, and so is proportional to the product of their concentrations. This is a type mechanism that describes a family of similar reactions. From stronger to weaker, some important nucleophiles, or bases, are NH2-, EtO-, OH- and CH3COO-.
Suppose we have some ethoxide anion, EtO-. This anion is strongly nucleophilic, so when ethyl bromide, EtBr is available, it may attack the positive carbon bound to the bromine. It donates its free electron pair, and the bromine goes off with the displaced electron pair, in accordance with the general mechanism of the SN2 reaction. This is shown schematically in the diagram at the left. The products are ethyl ether, EtOEt and a bromine ion. If we used MeO- instead, we would get MeOEt, methyl ethyl ether. It is clear that, in principle, we now have a means of synthesizing any alkyl ether we may want. If we use the cyanide (nitrile) ion CN- instead, we obtain EtCN. Now we have succeeded in lengthening the carbon chain by one carbon. If we can replace the nitrogen, we will have a propyl compound. Since there is a large variety of nucleophilic groups, we can make a large variety of compounds. Water will not work in this reaction, because HOH is too weak a base (it is a fairly strong base, actually, but not strong enough). F- will not work either, since it is too strongly solvated and will not donate its unbonded pairs, although it has three. Groups containing nitrogen work very well, since nitrogen is willing to donate its unbonded pair to make a new bond. NH3, in particular is a very strong base, forming the ammonium ion NH4+ if protons are available. In the substitution reaction, ammonia yields amines, such as EtNH2, primary ethylamine. The nitrogen can be further attacked in the same way, until (Et)4NBr, tetraethylammonium bromide, is formed. Now the nitrogen has no unbonded pairs left. These amines are bases, forming, for example, ions like EtNH3+, ethylammonium, that form salts.
However, the ethyl bromide also contains hydrogens, and the ethoxide anion could pick one off the neutral carbon that is not bound to the bromine and go off as EtOH, releasing an electron pair that then makes another carbon-carbon bond, again displacing the bromine on the other end. The net result in this case is a molecule of ethyl alcohol, a molecule of ethylene, and a bromine ion. This second reaction is called an elimination, denoted E2. The elements of HBr have been eliminated from the ethyl bromide. By means of this reaction, we can synthesize ethylene, an important source of further compounds, as well as further alkenes. The double bond will now be the point of attack.
Another kind of elimination is based on quaternary amines, such as CH3CH2N(CH23)3, where a nucleophile, say OH-, picks off one of the hydrogen atoms on the CH3 group. This creates a double bond between the carbons, and releases an electron pair to the electrophile nitrogen, which now happily goes its way as N(CH3)3, trimethyl amine (which smells like rotting fish). The result is ethylene, water and the amine. This is a general reaction for creating a double bond in whatever you attach to the amine, called Hofmann elimination. This reaction is an example of one of the many named reactions in organic chemistry, which are an organic chemist's stock-in-trade. The mechanism of the reaction may not be understood, but it still works.
There is another way to perform a substitution, which more closely resembles an inorganic ionic reaction. If an alkane should ionize like MeCl → Me+ + Cl- to form a carbonium ion Me+, this ion could then react quickly with any anion around, such as OH- in water, to form MeOH and the ions of HCl. Methane will not form a carbonium ion even with all the blandishments of polar water, but tertiary carbons (those bound to only one hydrogen) hold their hydrogens more weakly, and may be induced to form a carbonium ion, such as (Me)3C+ (tertiary butyl carbonium ion), which then rapidly forms tert-butyl alcohol if it finds any OH- around. The rate of this reaction depends only on the concentration of the carbonium ion, so it is effectively monomolecular. This kind of substitution reaction is denoted SN1. Carbonium ions are quite rare, so this reaction is also rare.
An unsaturated molecule such as ethylene can undergo reactions in which a pair of electrons is removed from the double bond and is used to attach new groups to the two carbons. Such reactions are called additions. When X-Y is added to H2C=CH2, the product is XH2C-CH2Y. Of course, the addition can occur across any C-C double bond, not just the one in ethylene. If X-Y is H2, the reaction is called hydrogenation. This is an exothermic reaction, as we can see by comparing the heats of formation of ethane and ethylene, and deducting the energy required to dissociate the H2. Using the table of bond energies given above, the heat of formation of ethane is 84 + 6 x 98 = 672 kcal/mol. The heat of formation of ethylene is 150 + 4 x 98 = 542, to which we must add the 102 kcal/mol required to break the H-H bond, for a total of 644 kcal/mol. The hydrogenation of ethylene is exothermic to the amount of 28 kcal/mol, so it is energetically favored.
However, if you mix ethylene and hydrogen, nothing will happen even if you wait forever. There is no mechanism that will lead to the desired final state of ethane from this initial state. Even light will not work, since it requires 4.4 eV to dissociate hydrogen and produce free radicals, and this is more than ordinary daylight can provide. If we had free radicals, however, they would certainly be able to pick apart the double bond, and this is one mechanism for the addition reaction.
An ionized proton, H+, however, actively seeks out electrons where it can find them, and when it does, it will form a bond of considerable strength, since the proton is small. The proton, and similar electron-seekers such as NH4+, are called electrophiles. They will attack the large electron density in a double bond, and will muscle in to share the bounty. A proton will go right ahead and make a new C-H bond, leaving the other carbon with a positive charge (since it has effectively lost an electron). This carbon will then woo any nucleophiles that happen to be around in its search for neutrality. HBr will add to ethylene to form H3C-H2Br, ethyl bromide. The mechanism of this reaction is shown at the right. Even Br2 will add in this way, since it ionizes slightly to form Br+ + Br-. To see this, you must be in the dark, since light easily dissociates bromine to make free radicals with much greater probability. The dissociation into ions is called heterolysis, while dissociation into free radicals is called homolysis. Either way we have addition to the double bond, in the first case by electrophilic attack, and in the latter by a free-radical reaction.
Sulphuric acid, HOSO3H (written this way to exhibit the hydrogen sulphate nucleophilic anion) also adds across the double bond in ethylene to form ethyl hydrogen sulphate, an ester that is hydrolyzed in water to form ethanol and sulphuric acid (recall that esterification is a reversible reaction). This reaction is the basis for the industrial process that makes ethanol from "cracked" petroleum, which contains ethylene. The demand for alcohol is so large that it can no longer be satisfied by alcohol from fermentation.
An alcohol can be dehydrated by a strong acid by electrophilic attack by the H+ on the oxygen of the -OH group of the alcohol, to add the proton and make an oxonium ion. This oxonium ion is then attacked by the nucleophile (the anion of the acid) which picks off a hydrogen from the end farthest from the oxygen, and electron pairs then cascade onto the HOH at the other end, releasing it as the amine was released above. The result, if ethanol was dehydrated, is water, acid and ethylene. We shall see that ethanol can be made from ethylene, so we can move easily from one to the other, depending on what is required. This is an example of addition followed by elimination.
Electrophilic attack by something larger than a proton forms a positively-charged bridge structure between the two carbons, which is easily disrupted by the approach of a nucleophile. Since the double bond draws electrons away from the carbons, they may fall victim to nucleophilic attack at the beginning. This makes a negatively charged carbon that will attract a positive ion, and again the addition is completed. In alkenes, this reaction usually does not occur, since the carbons are not positive enough to be attractive to a nucleophile. It does occur, however, for the C=O double bond, where the oxygen sucks electrons away from the carbon.
Boron, with its lack of valence electrons, is a strong electrophile, although not positively charged. Diborane, B2H6 can add to the double bonds in six molecules of ethylene to form two molecules of triethylene boron (or ethylene borane). This molecule hydrolyzes in acidic water to form ethane and boron hydroxide, or in alkaline hydrogen peroxide to form ethyl alcohol and boron hydroxide. The boron attacks the double bond, and bonds to one carbon, while the proton released bonds with the other carbon. This is yet another way of making alcohol out of petroleum.
Hydrogenation of olefins is an important reaction, in petroleum refining and other industrial processes. It will proceed catalytically in the presence of an active metal surface. Ni, Pt or Pd are often used, in "black" colloidal form. Free bonds on the metal do the job of dissolving the double bond when the olefin is adsorbed on its surface, acting like free radicals. Hydrogen is dissociated in the same way. The molecules can diffuse on the metal surface, and when they meet the hydrogens (protons) prefer to bond with the adsorbed olefin rather than with the metal. The hydrogenated molecule is then freed to go on its way.
The C=O double bond appears in aldehydes, ketones and acids, so it is very important. Unlike the C=C bond, it is polarized, with the O end negative and the C end positive, just like the C-Cl bond we met above. Nucleophiles will add to the C, and electrophiles to the O, and both kinds of reactions will be common with the C=O bond. For example, when chloral, trichloroacetaldehyde, Cl3C-CHO, is treated with water, OH- will attack the C and stick to it, while H+ will attack the O and stick to it, so that we have two OH groups on the same carbon. In this case, we get chloral hydrate ("knockout drops"), which is quite stable, but in most cases the product is not stable and exists only in solution.
If we mix formaldehye, HCHO, and ammonia, NH3, NH3+CH2O- will be formed. The charge distribution on this molecule is as indicated, which easily persuades the molecules to polymerize. Four of these polymerize to a curious molecule called hexamethylenetetramine, also called urotropine, shown at the right, where the three-dimensional structure is suggested. This was the first organic molecule whose structure was determined by X-ray diffraction, in 1923. Its symmetry is tetrahedral, with the carbons at the vertices of a regular octahedron. The bond angles are tetrahedral, and the bond lengths agree with the covalent radii in the table above. This molecule forms colorless cubic crystals, and is soluble in water because of the unbonded electrons on the nitrogens. It irritates the skin, burns readily, is a urinary antiseptic and diuretic, and dissolves uric acid in cases of kidney stone and gout. It was used in an alkaline solution with phenol and glycerine as an antidode to phosgene, (COCl2), which is closely related to formaldehyde. 7 pages are devoted to it in the 4th edition of Beilstein (V. I, p. 583). A picture of a space-filling model is shown at the left. Two nitrogens (blue) marking 3-fold axes, can be seen. When it is nitrated, electrophilic attack by the NO2+ on the nitrogens cleaves the molecule at the nitrogens and adds a nitro. The fragments reform to make cyclotrimethylenetrinitramine, with a ring of alternating N and C, a nitro at each N and H2 at each C. This is a powerful explosive, also called cyclonite or RDX, which is easily detonated with mercuric fulminate while being more powerful than TNT. It is notable as not being an nitrate ester, but on the other hand it has N-N bonds, which are also explosive.
The proton from HCl can mount an electrophilic attack on the O atom of dimethyl ketone, forming OH there and making the carbon atom positive. This resembles a carbonium ion, and is reactive enough to attract an electrophilic attack by the unbonded electrons of an alcohol, say ethanol, which become bonded to the carbon of what was the C=O group. Now the positive charge on the O of the alcohol OH group migrates down to the other O as the OH group ionizes. Now this oxygen can be attacked by a proton, and after it adds and H, HOH separates readily. Now we are left with (CH3)2COEt with a + charge on the central carbon. This carbonium-like group can now attract a second alcohol OH, and now Cl- ion takes away its hydrogen and departs as HCl. The result is (CH3)2C(OEt)2, acetal, with two ether-like oxygens. Since the HCl is restored, the net result is the addition of two OEt's across the carbon-oxygen double bond.
The C=O double bond can also be catalytically hydrogenated like the C=C double bond, but the reaction is more sluggish, since it is not as favored energetically. The addition of hydrogen is called reduction in organic chemistry, like the addition of electrons in inorganic chemistry. An aldehyde or ketone is reduced to an alcohol in this way. An oxidizing agent such as the permanganate ion, MnO4- (Mn +7), can add an oxygen to the carbon atom of the carbonyl group of RCHO and leave it there by becoming manganate, MnO3- (Mn +5). Now we have RCOO- + H+ + MnO3-, and the aldehyde has been reduced to a carboxylic acid.
The carboxyl, COOH, group contains the C=O of a ketone and the OH of an alcohol, but the two groups interact so strongly that the reactions of the carboxyl group are distinct from those of ketones and alcohols. We have already mentioned that the carboxyl group ionizes more readily than the OH of an alcohol, and that resonance causes the two oxygens of COO- (the carboxylate anion) to be equivalent, with a positive charge on the carbon and a negative charge on the oxygens.
An important reaction of the carboxyl group is esterification, whose mechanism is shown in the figure. When acetic acid and ethanol are mixed in an acidic solution, which can be aqueous since both are water-soluble, the ester ethyl acetate (with the characteristic aroma of bananas) and water are created. Both reagents ionize, the acid more than the alcohol, and the ethoxy group makes a nucleophilic attack on the carbon of the acetate ion. This releases an oxygen to form OH-, which is then quickly protonated to form water. The reaction is reversible, so that ethyl acetate is hydrolyzed by water. Ethyl acetate is not an ionic compound, since the carbonium ion Et+ does not exist. Esterification is not analogous to NaOH + HCl → NaCl + HOH. NaCl is not hydrolyzed by water to sodium hydroxide and hydrochloric acid. An ester, therefore, bears no similiarity to an inorganic salt.
An ester is the compound formed by taking the H from an alcohol, and the OH from the carboxyl group of an acid, and joining the remainders. For example, EtOH + CH3COOH → HOH + EtO-OCCH3. Ethyl alcohol and acetic acid yield the ester ethyl acetate (which smells like bananas). This might look like the H from the COOH combined with the OH from the alcohol, as in MeCOOEt, but that is not the case. The COOH can be induced to give up the OH by nucleophilic attack by the ethoxide anion on the carbon. This reaction is reversible, and does not go to completion unless water is removed. It goes very slowly indeed unless a catalyst is present, such as a strong acid or base.
A similar thing happens in the nitration of glycerol, (CH2OH)2CHOH, with nitric acid, HO-NO2. The result is (CH2ONO2)2CHONO2, an ester called glyceryl trinitrate or nitroglycerine. The ONO2 or NO3 groups make the substance explosive. Methyl nitrite, or nitromethane, CH3NO2, is a nonexplosive gas, while methyl nitrate, CH3ONO2, an ester of methanol and nitric acid, is an explosive liquid. Nitration is carried out in a mixture of concentrated nitric and sulphuric acids. The sulphuric acid takes up any water produced, as well as acting as a catalyst. In the mixture, nitric acid is very little dissociated, and the nitronium ion, NO2+ is formed that reacts readily with the ionized glycerol. Schönbein discovered this method of nitration when making guncotton in 1845, and it was used by Sobrero in 1846 to produce the ester glyceryl trinitrate, or nitroglycerine.
Other nucleophiles add to the carboxyl group equally readily as EtO-. For example, ammonia adds to the C=O carbon in ethyl acetate in a typical nucleophilic addition. Then the leaving of EtO- reestablishes the double bond, producing acetamide, Me(CO)NH2. Further reaction with ammonia adds another NH2, producing urea, CO(NH2)2. The amides are like carboxylic acids, but with an NH2 group instead of OH. The group -(C=O)-(NH)- is called a peptide link, which joins the amino acids in proteins. Carboxylic acids will not succumb to nucleophilic attack on the C of the COOH group. They simply form the stable carboxylate ion and throw a proton at the hungry nucleophile. Ammonia and acetic acid give ammonium acetate, which is an ionic salt and not an ester.
Metal atoms will bond covalently to organic groups to form organometallic compounds. We don't include ionic compounds like lead acetate, Pb(CH3COO)2, in this category. Tetraethyl lead, Pb(Et)4, is a typical organometallic compound. In 1849, Frankland found that by combining an alkyl halide RX with powdered zinc in ether, a stable compound RZnX was produced in which the carbon attached to the Zn acquired a negative charge and so could be attacked by electrophilic reagents. You will recall that in RX the carbon attached to the X acquires a negative charge and is exposed to nucleophilic attack. The new organometallic compound made many additional reactions possible. The negative carbon is also a nucleophile, and can mount nucleophilic attacks itself.
In 1901, Victor Grignard found that Mg was even better than Zn for making these compounds, which are now called Grignard reagents. Bromine is the best halogen to use, since Cl is less reactive, and I gives too many side reactions. A typical Grignard reagent, then, is EtMgBr, made by mixing ethyl bromide and zinc powder in pure ether. The reagent can be crystallized, but it is generally used in the same solution in which it was prepared. A Grignard reagent such as EtMgBr is readily hydrolyzed by water, since the nucleophilic carbon attracts the electrophilic H+ of HOH to give C2H6 + Mg(OH)Br.
To see how this could be useful, consider the reaction of EtMgBr with formaldehyde. The Et adds to the carbon of the formaldehyde by electrophilic addition to the C=O double bond, and the MgBr to the oxygen at the other end, making C3H7OMgBr. This product can be hydrolyzed, as just explained, to C3H7OH, propyl alcohol, and Mg(OH)Br. We have added CH2 to the carbon chain, and this is a general method of ascending a homologous series.
By using a ketone instead, we get a tertiary alcohol by the same process. If we start with dimethyl ketone and MeMgBr, we end up with 2-isobutanol, Me(CCH3OH)Me. Methaldehyde instead of dimethyl ketone gives 2-propanol. If we start with a carboxylic ester, such as ethyl acetate, we first get dimethyl ketone (by rearrangement of the addition product that eliminates EtO and recreates the double bond) and EtOMgBr. A further attack by EtMgBr produces a Grignard reagent that will hydrolyze to tertiary butanol.
Grignard reagents even attack the positive carbon in CO2. CO2 + EtMgBr → EtCOO- + MgBr+ → EtCOOMgBr, which hydrolyzes to EtCOOH, propionic acid. To avoid further reaction of the desired acid product, the reaction is carried out by pouring the fresh Grignard reagent on powdered solid CO2. Then the addition of water creates the acid before the slow second reaction can take place.
EtMgBr will react with PbCl2 to make tetraethyl lead: 4EtMgBr + 2PbCl2 → Pb(Et)4 + 4MgClBr + Pb. This is not the industrial process, which uses EtCl and PbNa alloy instead. Tetraethyl lead is a liquid, which was added to motor fuels so that ethyl free radical would be released at the proper time to smooth out combustion in Otto cyle engines. Incidentally, Diesel engines prefer a fuel that ignites immediately upon injection, so straight-chain hydrocarbons are the best Diesel fuel, while branched hydrocarbons are preferred by Otto-cycle engines.
The benzene molecule is a strong fortress, symmetrical and tightly bound. The ring carbons and their hydrogens are unpromising targets. The one weakness is the six ring electrons, three on one side and three on the other, which may be captured by strong electrophiles. An electron pair may be donated to a Br+ ion, to bond it with the molecule with the six hydrogens. If one of the hydrogens is then sacrificed to an electrophile, the bromine will have succeeded in muscling into the ring in place of a hydrogen. Note that the loss of the hydrogen re-establishes the stable ring electrons. The bromobenzene that results is now open to chemical attack of many kinds. This reaction is illustrated at the right, together with the usual means of obtaining positive bromine. The ferric bromide acts as a catalyst, since it is not consumed in the reaction. The reaction is best described as addition followed by elimination, not as a substitution, since the addition and elimination are consecutive, not simultaneous. An activation energy is required to attain the middle state, then a much smaller activation energy when the hydrogen is emitted.
A very important reaction makes use of aluminium chloride in the same way that ferric bromide was used in bromination. Methyl chloride will form a complex with aluminium chloride through the reaction MeCl + AlCl3 → Me(+)-AlCl4(-). The Me(+) is a powerful electrophile, like the corresponding carbonium ion, and can successfully attack the benzene ring, and add Me to it, producing methyl benzene, or toluene, BzMe (Bz = C6H5). This is called the Friedel-Crafts reaction. Compare with the Grignard reaction described above.
Benzene can also be nitrated by nitric acid in anhydrous sulphuric acid, in which the very electrophilic nitronium ion, NO2+, exists, making nitrobenzene. Toluene is easily nitrated to nitrotoluene, where the nitro group is ortho to the methyl. Further nitration results in trinitrotoluene or TNT, Bz(NO2)3, with the three nitros ortho and para to the methyl. This is not a nitrate ester, but is still quite explosive, like the other two explosive arrangments we have found, O-NO2 and N-NO2. Further additions to the benzene ring can be made after the first one, and are generally easier than the first one. The additions may be directed to a particular position. With toluene, the nitro goes ortho, but a second methyl will go meta, to form m-xylene.
Potassium permanganate will not oxidize benzene, but it will oxidize the methyl group in toluene, making a (C=O)OH while eliminating hydrogen. Even the methyl group is more readily attacked than the benzene nucleus. The result in this case is benzoic acid, BzCOOH. Nitrobenzene can be reduced (hydrogenated) to aminobenzene, or aniline, BzNH2, while the benzene ring is unaffected. Chlorine free radicals will attack the benzene ring when they have fully conquered all easier targets. The result is 1,2,3,4,5,6-chlorocyclohexane, an aliphatic compound. There are many stereoisomers, since the chlorines and hydrogens can be above or below the ring ad libitum. Usually all the chlorines are on one side (cis), or chlorines and hydrogens alternate (trans). One of the stereoisomers is a powerful insecticide. This compound should not be confused with hexachlorobenzene, C6Cl6, which is aromatic. It is sometimes confusingly called benzene hexachloride.
Halobenzenes form Grignard reagents with magnesium, which render the benzene ring susceptible to electrophilic attack, and allow the addition of benzene rings by nucleophilic addition. For example, BzMgBr is formed the usual way by treating powdered magnesium with benzene in ether.
OH- will not attack BzCl as it attacked RCl, because no positively charged carbon exists. However, the reaction will go at the high temperatures and pressures of an industrial process, and phenol, BzOH, is formed. The OH group behaves differently when attached to the benzene nucleus than when attached to an alkane. Because of resonance in the BzO- ion, which is not as strong as with the carboxyl group but still acting, the H is more easily ionized and phenol is a weak acid. When the H+ leaves, the negative charge on the oxygen can be spread around the ring. The ionization constant of ethanol is only about 10-18, but that of phenol is 1.3 x 10-10.
If we treat ethyl alcohol with nitrous acid, HO-N=O, we get the ester ethyl nitrite, EtO-N=O by the normal esterification reaction. With ethylamine, we get instead: EtNH2 + HO-N=O → EtNH-N=O + H2O. The product rearranges by electrons moving from left to right to Et-N=N-OH and then ionizes as Et-N=N+ + OH-, forming a diazonium cation. This has two resonating forms, the one shown, and Et-N+≡N. The diazonium cation is unstable, decomposing to N≡N and a carbonium ion. In this case, the carbonium ion forms EtOH. The word "azo" refers to nitrogen (the French name for nitrogen is azote), since all the other good words for nitrogen are already taken.
There is an aromatic amine, aminobenzene, ususally called aniline, BzNH2. When we treat it with nitrous acid and hydrochloric acid, we get benzenediazonium, BzN+≡N, as the chloride. This ion is stabilized by having several resonance structures that spread the charge over the benzene nucleus, so it is a reasonably stable molecule. It is a strong electrophile, so it can attract OH-, Cl-, CN- and other anions, then eliminate N2 and tack the anions onto the benzene nucleus to get phenol, chlorobenzene and benzonitrile, respectively. Aniline is prepared from benzene by nitration with HNO3 in H2SO4 to make nitrobenzene, followed by reduction with H2 and Ni catalyst, or by chemical reduction by Sn in HCl.
Benzenediazonium will add phenoxide, BzO-, but to the outer nitrogen and form the stable molecule Bz-N=N-Bz-OH, p-hydroxyazobenzene. This is called a coupling reaction. The azo group -N=N- is stabilized by having a benzene nucleus on each end, and less triple bond character. The result of this reaction has a bright orange color, and is the simplest of the azo dyes. The -N=N- group is the chromophore, responsible for the absorption of light, while the -OH is the polar auxochrome group, that binds the dye to the fiber, or to the mordant for nonpolar fibres. Methyl Red has an N(Me)2 on one benzene ring, and COOH on the other. Methyl Yellow has only the methamine group. Methyl Orange has the methamine group on one ring, and -SO2OH on the other. All these dyes have the same chromophore, but different auxochromes. Furthermore, in aqueous solution they change color depending on the pH (they have two forms depending on the pH), and so are used as indicators. Aniline is the starting point for the synthesis of many dyes because of the general coupling reaction that introduces a chromophore.
While we are talking about dyes, it may be interesting to consider a very old dye that is still used, the blue dye indigo, familiar from blue jeans. It was prepared by oxidation of the substance indican, a glucoside from the Indigofera tinctoria and related plants, which are legumes. At the present time indigo is prepared synthetically. The chromophore is the O=C-C=C-C=O chain in the keto form of the molecule. This insoluble compound is rendered soluble by reduction with NaHSO3, which alters it to the colorless but soluble enol form. After the cloth is dyed, exposure to the air restores the blue form by oxidation. This procedure is typical of a vat dye, of which indigo is an example. Another ancient dye was red madder, from the roots of Rubia tinctorum, of which a close modern relative is alizarin, an anthraquinone dye. These dyes are based on a nucleus of two benzene rings joined by two C=O ketone bridges. Indigo and madder are good dyes, but monotony of color was the rule until artificial dyes were introduced. The first synthetic dye was yellow picric acid, discovered in 1771. Coal-tar dyes (those based on benzene rings) were discovered in 1856 by Perkin.
A dye must be insoluble, or else bond strongly with the fabric, or it will be washed out. On the other hand, it must be soluble to be distributed over the fabric. These conflicting requirements may be met as described above for vat dyes, which have two suitable forms. Some dyes are so insoluble that they must be made "in place" on the fabric. These are called developed dyes. A good example is the azo dye Para Red, which is brilliant red, whose structure is shown at the right. It is made by the coupling reaction of para-nitrodiazobenzene (hence the name Para) and β-naphthol, a phenol-like derivative of naphthalene (see Organic Compounds for its structure). The fabric is first soaked in a solution of the sodium salt of β-naphthol and coated with the BzBzO- ion. Then it is passed through a solution of nitrobenzenediazonium chloride, which forms Para Red in intimate contact with the fibres. Para Red can be prepared in powder form for use in coloring smokes or for dispersion in plastics.
A bright yellow dye, auramine, is made as shown at the left. Without the chlorine, the compound is colorless, but the addition of chlorine induces the structure shown, which has numerous alternating double bonds. Auramine is an example of a ketoimine dye. It is not very fast, but is an excellent yellow that is used to color smoke. With indigo, it makes a green smoke by additive color mixing. The chromophore in this case seeems to be the entire upper ring system. The -NH2 and -N(CH3)2 groups are auxochromes.
Methylene blue, whose structure is shown at the right, is an important biological stain. Tissues are almost transparent under the microscope, but stains tint tissues selectively, allowing them to be distinguished. This was first reported in 1838. Methylene blue stains cell nuclei and bacteria. It is an example of a thyazine dye. The important dyes indigo, para red, auramine and methylene blue have now been introduced. There are still many other kinds of dyes, but these examples should show how organic compounds can exhibit infinite variety.
Although α- and β-naphthol occur in coal tar, the amounts available from this source are too small to support the dye industry. They can, however, be synthesized from the more abundant naphthalene with the aid of sulphuric acid. Before discussing this, let's look at some important organic compounds containing sulphur. Sulphur with valence -2 can replace oxygen in many compounds, which are similar to the oxygen analogues. Sulphur also has a great affinity for electrons, like oxygen, but its unbonded pairs are not as tightly held, and do not form hydrogen bonds of the same strength. Sulphur exhibits positive valence states of +2, +4 and +6, notably in compounds with oxygen, and can be a reducing agent as well as an oxidizing agent.
Two series of compounds can be derived from hydrogen sulphide, H2S, the analogue of water. Unlike water, hydrogen sulphide melts at -83.9°C and boils at -59.6°C, because of the lack of hydrogen bonding. It is a weak acid in aqueous solution, giving H+ and SH- ions, and a little S-- by further ionization. It is a dangerous poison, though the effects of a little wear off quickly in fresh air. It smells, as is well known, like rotten eggs.
If one hydrogen is replaced by an organic radical R, a thioalcohol or mercaptan is formed. MeSH is methyl mercaptan or methanethiol, boiling at 6°C. EtSH is ethyl mercaptan, boiling at 35.1°C. Propyl mercaptan can also be described as 1-propanethiol, and isopropyl mercaptan as 2-propanethiol. The alkyl mercaptans have foul smells, and are added in small amounts to natural gas to give it a detectable odor. Phenyl mercaptan, BzSH, is better known as thiophenol. It is a liquid over the wide range -15° to 170°C. Its odor must be peculiar. I expect the -SH will overpower the -Bz.
If both hydrogens are replaced, a thioether is formed. Dimethyl sulphide, (Me)2S, boils at 37.3°C, while dimethyl ether boils at -23.7°C. This shows the attraction between the sulphur and the methyl group is greater than the attraction between the sulphur and the hydrogens in hydrogen sulphide, and also stronger than the attraction between oxygen and a methyl group.
Sulphur burns in air or oxygen to form sulphur dioxide, which dissolves slightly in water to form sulphurous acid. The first ionization of sulphurous acid gives the bisulphite ion, which appears in several useful inorganic salts. Sulphur dioxide oxidizes slowly in oxygen to form the trioxide. Processes that speed up the oxidation by catalysis are those that make sulphuric acid, the most useful industrial chemical. Sulphur dioxide dissolves readily in water to form sulphuric acid, which is only partly ionized. 100% sulphuric acid can even dissolve more SO3, becoming fuming sulphuric acid, or oleum. The bisulfate ion is analogous to the bisulfite ion, forming acid sulphate salts. Of course, both sulphite and sulfate ions are also found, resulting from complete ionization of the acids.
The SO3 in fuming sulphuric acid is a very strong electrophile, like the NO2+ ion, and can likewise attack the benzene ring electrons. The result is replacement of a hydrogen by the sulphone radical, as shown in the figure, producing benzenesulphonic acid. Note that the ring carbon is bound directly to a sulphur, so this is not a sulphate ester. There is a similar sulphine group, but we are not concerned with it here. Sulphonic acids are reasonably strong, as may be expected. There is an homologous series of alkyl sulphonic acids beginning with methanesulphonic acid, HMeSO3. Note that these are not sulphites, but have a different structure. They can be produced by reaction with sodium sulfite, however: RCl + Na2SO3 → RSO2OH + NaCl. This way of writing the sulphone distinguishes it from a sulphite. The other hydrogen can also be replaced, giving sulphones like Ar-SO2-Ar, where Ar is an aromatic radical.
The great advantage of the sulphonic acids is that the defenses of the benzene ring have been broken down, and further reactions are possible. Among the many possibilities, we are interested in alkali fusion, which has the effect of eliminating the bracketed atoms in Bz[SO2ONa + H]ONa, producing BzONa which hydrolyzes to BzOH, or to the BzO- ion. This is the way to make phenol from benzene. Phenol can be triply sulphonated, from which the trinitrate, picric acid, can be made.
Naphthalene, as we have seen, consists of two fused benzene rings. It can be sulphonated in exactly the same way as benzene. From 80° to 100°C, &alpha-naphthol is produced; at 160°C the product is &beta-naphthol. We actually make the negative ions, but that is just what is desired to carry out the coupling reactions of dye synthesis.
Organic chemistry provides fuels and industrial chemicals from petroleum, coal and biological materials. It supplies plastics, dyes, fibres, pharmaceuticals, detergents, rubber and paints. Chemical engineering is largely organic chemical engineering, carrying out organic reactions on an industrial scale.
G. Herzberg, Atomic Spectra and Atomic Structure (New York: Dover, 1944). Essential familiarization with the properties of electrons.
L. Pauling, The Chemical Bond (Ithaca, NY: Cornell U. P., 1967). Elementary quantum mechanics applied to chemical bonds.
L. Pauling and R. Hayward, The Architecture of Molecules (San Francisco: W. H. Freeman, 1964). Hexamethylenetetramine is No. 31.
J. M. Tedder and A. Nechvatal, Basic Organic Chemistry (London: John Wiley & Sons, 1966). An introductory text based on reaction mechanisms.
Ed. F. Degering (editor), Organic Chemistry, 6th ed. (New York: Barnes & Noble, 1951). A College Outline Series edition and excellent reference.
P. W. Atkins, Molecules (New York: W. H. Freeman, 1987). Interesting facts molecule by molecule, excellently illustrated with pictures of space-filling molecular models. Reactions, of course, are not explained.
Beilsteins Handbuch der organischen Chemie, 4te Aufl. (Berlin: Springer, 1922). The library of any university with a chemistry department will have this reference, but only the better public libraries.
L. Stryer, Biochemistry, 3rd ed. (New York: W. H. Freeman, 1988). A wonderful account of a prime application of organic chemistry. So much is now known about genetics, protein synthesis, respiration, nerve and muscle function, and sensory perception that it is clear how much we have still to learn. It is worth learning some organic chemistry just to be able to enjoy this book.
Composed by J. B. Calvert
Created 4 January 2003
Last revised 8 January 2003