6.4     Properties of Substituted Alkanes

6.4a Polarity Substituted alkanes are polar molecules. Halogens are relatively electronegative atoms and strongly attract the electrons in the carbon–halogen bond of alkyl halides, which have dipole moments of approximately 2 D (Fig. 6.23; Table 6.2).

 

These polar molecules can associate in solution, and therefore the boiling points tend to be higher than those of their hydrocarbon counterparts (compare Tables 2.4 and 6.1).

TABLE 6.2 Dipole Moments of Some Simple Alkyl Halides

 

In alcohols (ROH), the presence of the electronegative oxygen atom ensures that bonds are strongly polarized and that substantial dipole moments exist. In the liquid phase, alcohols are strongly associated, both because of dipole–dipole attractions and because of hydrogen bonding, in which the basic oxygen atoms form partial bonds to the acidic hydroxyl hydrogens. As you probably remember from general chemistry, hydrogen bonding is the noncovalent interaction between a hydrogen that is covalently attached to an oxygen (or nitrogen or fluorine) and another oxygen (or nitrogen or fluorine) to which it is not covalently attached. Such interactions can be quite strong, approximately 5 kcal/mol (21 kJ/mol), a value not high enough to make permanent dimeric or polymeric structures but quite sufficient to raise the boiling points of alcohols far above those of the much less strongly associated alkanes. Figure 6.24 shows an ordered, hydrogen-bonded structure for an alcohol. In the absence of hydrogen bonding, water (with a molecular weight of only 18) and the low molecular weight alcohols would surely be gases. The polarity of alcohols makes them quite water soluble, and most of the smaller molecules are miscible (soluble in all proportions) with, or at least highly soluble in, water. Table 6.3 gives some physical properties of alcohols. For comparisons with the parent alkanes, see Table 2.4 (p. 83).

TABLE 6.3 Some Physical Properties of Alcohols

 

Even though these intermolecular hydrogen bonds are relatively weak (~5 kcal/mol) they are critically important. For example, life as we know it is clearly not possible without water, and there would be no liquid water without hydrogen bonds. Moreover, hydrogen bonds are used to maintain the proper structures in proteins and nucleic acids, polymeric structures essential to our existence that we will meet later in this book.

Amines, like alcohols, have much higher boiling points than those of hydrocarbons of similar molecular weight, although the effect is not as great as with alcohols (Table 6.4; compare to Table 6.3). The reason for the increased boiling points is the same as for alcohols. Amines are polar compounds and aggregate in solution. Like alcohols, small amines are miscible with water. Boiling requires overcoming the intermolecular attractive forces between the dipoles. Moreover, amines form hydrogen-bonded oligomers in solution, effectively increasing the molecular weight and further increasing the boiling point (Fig. 6.25).

When hydrogen bonding is impossible, as with tertiary amines, boiling points decrease (Table 6.4).

TABLE 6.4 Some Physical Properties of Amines and a Few Related Alkanes

It is impossible to discuss the properties of amines without mentioning smell. The smell of amines ranges from the merely fishy to the truly vile. Indeed, the odor of decomposing fish owes its characteristic unpleasant nature to amines. Diamines are even worse. Some indication of how bad-smelling they often are can be gained from their names. 1,4-Butanediamine is called putrescine, and 1,5-pentanediamine is called cadaverine.

 

Ethers are polar, but only for the smallest members of the class does this polarity strongly affect physical properties. Other ethers are sufficiently hydrocarbon-like so as to behave as their all-carbon relatives. For example, diethyl ether has nearly the same boiling point as pentane and is only modestly soluble in water. Table 6.5 gives some physical properties for a few common ethers.

6.4b Acidity and Basicity Alkyl halides are neither substantial acids nor bases. They are not proton donors and have no appreciably basic electron pairs. Alcohols and amines, however, can be proton donors and, because of their lone-pair electrons, can be bases as well.

We have already seen water and alcohols acting as Brønsted acids and bases. Hydrogen bonding is a perfect example of such behavior. In Chapter 7, we are going to look closely at several reactions of alcohols. These processes depend on the ability of alcohols to act as both acids and bases. Accordingly, let’s first make sure that we can see these roles clearly.

TABLE 6.5 Some Physical Properties of Ethers

A Brønsted acid is any compound that can donate a proton. A Brønsted base is any compound that can accept a proton. Let’s look first at a very simple process, the reaction of solid potassium hydroxide (KOH) with gaseous hydrochloric acid (HCl). This reaction is nothing but a competition of two Brønsted bases (Cl⁻ and HO⁻) for a proton, H⁺. The stronger Brønsted base (HO⁻) wins the competition (Fig. 6.26).

 

Hydrochloric acid (HCl) is called the conjugate acid of Cl⁻, and Cl⁻ is the conjugate base of hydrochloric acid. Conjugate acids and bases are related to each other through the gain and loss of a proton (Fig. 6.27).

 

As we saw in Figure 6.24, a molecule may be either a Brønsted acid or a Brønsted base. Consider water, for example. Water can both donate a proton (act as a Brønsted acid) and accept a proton (act as a Brønsted base). Figure 6.28 shows two such reactions: (a) the protonation of water by hydrogen chloride, in which water acts as a base, and (b) the protonation of hydroxide ion by water, in which water acts as the acid.

 

Many reactions begin with a protonation step. Because strong acids are better able to donate a proton than are weaker acids, it will be important for us to know which molecules are strong acids and which are weak acids. The dissociation of an acid HA in water can be described by the equation

HA + H₂O ⇄ H₃O⁺ + A⁻

This equation leads to an expression for the equilibrium constant, K.

Because it is present as solvent, in vast excess, the concentration of water remains constant in the ionization reaction. This equation is usually rewritten to give the acidity constant, K_a,

By analogy with pH, we can define a quantity pK_a.

pK_a = −log K_a

The stronger the acid, the lower is its pK_a. Table 6.6 gives pK_a values for some representative compounds. This short list spans about 80 powers of 10 (pK_a is a log function). In practice, any compound with a pK_a lower than about +5 is regarded as a reasonably strong acid; those with pK_a values below 0 are very strong. As we discuss new kinds of molecules, more pK_a values will appear. Should you memorize this list? We think not, but you will need to have a reasonable idea of the approximate acidity of different kinds of molecules.

TABLE 6.6 Some pK_a Values for Assorted Molecules^a

 

In alcohols and water, Brønsted basicity is centered on the nonbonding electrons of the oxygen atom, which form bonds with a variety of acids. Protonation converts an alcohol, ROH, into an oxonium ion (RO⁺H₂). This conversion has profound chemical consequences, which we can preview here. Breaking the R―OH bond to form ions is very difficult, because both positive (R⁺) and negative (HO⁻) ions must be formed. Hydroxide, HO⁻, is very difficult to form—we say that hydroxide is a poor leaving group—and alcohols do not ionize easily (Fig. 6.29). After protonation, however, the “leaving group” is no longer hydroxide, but water. Formation of the neutral molecule water is a much easier process (Fig. 6.29).

 

Protonated alcohols are very strong Brønsted acids. Therefore, in order to protonate an alcohol, it takes a very strong acid indeed. The conjugate base (B:⁻) of the protonating acid (HB) must be a weaker base than the alcohol. The base must be an ineffective competitor in the equilibrium of Figure 6.29.

Oxonium ions are much better proton donors (acids) than are alcohols themselves. Table 6.7 gives the pK_a values for a few protonated alcohols and water. Oxonium ions formed from alcohols have pK_a values in the −2 range, and those from ethers (R―O―R′) are even more acidic.

TABLE 6.7 Some pK_a Values for Simple Oxonium Ions

 

Brønsted acidity is centered on the hydroxyl hydrogen. Loss of this hydrogen to an acceptor, a Brønsted base, leads to the conjugate base of the alcohol, the alkoxide ion (RO⁻). The proton is not just lost as H⁺ but removed by a base (Fig. 6.30).

 

Alcohols are approximately as acidic as water. Table 6.8 gives the pK_a values for some simple alcohols in aqueous solution and for water itself.

TABLE 6.8 Some pK_a Values for Simple Alcohols in Water

 

Note that the acidity order in water is CH₃OH > CH₃CH₂OH > (CH₃)₂CHOH > (CH₃)₃COH. Apparently, the more alkyl groups on the alcohol, the weaker an acid it is. For many years, this pK_a order was explained by assuming that alkyl groups are intrinsically electron donating. If this were true, the product alkoxides would be destabilized by alkyl groups, and the acidity of the corresponding alcohol would be reduced (Fig. 6.31).

Inductive effects, which result from polarized sigma bonds, are important. For example, the pK_a of 2,2,2-trifluoroethanol (CF₃CH₂OH) is 12.5, whereas the pK_a of ethyl alcohol is 15.9. The fluorinated alcohol is more than a thousand times as acidic as ethyl alcohol (remember the pK_a scale is logarithmic). Fluorinated alcohols are always stronger acids than their hydrogen-substituted counterparts.

 

 

Yet this traditional, and simple, explanation that treats alkyl groups as intrinsically electron-donating is not correct. In 1970, Prof. John Brauman (b. 1937) and his co-workers at Stanford University showed that in the gas phase, the opposite acidity order is obtained. The intrinsic acidity of the four alcohols of Table 6.8 is exactly opposite to that found in solution. The acidity order measured in solution reflects a powerful effect of the solvent, not the natural acidities of the alcohols themselves. Organic ions are almost all unstable species, and the formation of the alkoxide anions depends critically on how easy it is to stabilize them through interaction with solvent molecules, a process called solvation. tert-Butyl alcohol is a weaker acid in solution than methyl alcohol because the large tert-butyl alkoxide ion is difficult to solvate. The more alkyl groups, the more difficult it is for the stabilizing solvent molecules to approach (Fig. 6.32). In the gas phase, where solvation is impossible, the natural acidity order is observed.

In the gas phase, the alkyl groups actually operate so as to stabilize the charged alkoxide ions by allowing electron delocalization into alkyl group empty antibonding orbitals, exactly the opposite from what had been thought (Fig. 6.33).

 

What lessons are to be drawn from this discussion? First, solvation is important and not completely understood. We are certain to find other phenomena best explained in terms of solvation in the future. Second, the gas phase is the ideal medium for revealing the intrinsic properties of molecules. (Some would go further and say that calculation is the best way.) However, the practical world of solvated reactions is certainly real, and for us to understand reactivity we cannot afford to ignore the solvent. We will return to the subject of solvents and solvation in Section 6.5.

Like water and alcohols, amines are Brønsted bases (proton acceptors). Ammonia and primary, secondary, and tertiary amines are all weak bases (Fig. 6.34). Protonation of the nitrogen results in the formation of an ammonium ion. This process is significant in biological chemistry. When biological reactions need a base, it is often an amine that plays that role (Fig. 6.35).

 

The better competitor an amine is in the proton-transfer reaction, the stronger is its basicity. A way of turning this statement around is to focus not on the basicity of the amine but on the acidity of its conjugate acid, the ammonium ion. If the ammonium ion is a strong acid, the related amine must be a weak base. If it is easy to remove a proton from the ammonium ion to give the amine, the amine itself must be a poor competitor in the proton-transfer reaction, which is exactly what we mean by a weak Brønsted base. Strongly basic amines give ammonium ions from which it is difficult to remove a proton: ammonium ions with high pK_a values. So the pK_a of the ammonium ion has come to be a common measure of the basicity of the related amine. Ammonium ions with high pK_a values (weak acids) are related to strongly basic amines, and ammonium ions with low pK_a values (strong acids) are related to weakly basic amines (Fig. 6.36).

 

Table 6.9 relates the basicity of ammonia and some very simple alkylamines by tabulating the pK_a values of the related ammonium ions.

TABLE 6.9 Some pK_a Values for Simple Ammonium Ions in Solution

 

The general trend of the first three entries in Table 6.9 is understandable if we make the analogy between ammonium ions and carbocations. Remember (Chapter 3, p. 142): The more substituted a carbocation, the more stable it is. Similarly, the more substituted an ammonium ion, the more stable it is. The more stable the ammonium ion, the less readily it loses a proton and the higher its pK_a. At least for the first three entries of Table 6.9, increasing substitution of the ammonium ion carries with it a decreasing ease of proton loss.

Actually, even the first three entries in Table 6.9 should make you suspicious. Why is there a much bigger change between ammonia and methylamine (1.4 pK_a units) than between methylamine and dimethylamine (only 0.15 pK_a units)? The structural change is the same in each case—the replacement of a hydrogen with a methyl group. When we get to the fourth entry, we find that things have truly gone awry: trimethylammonium ion is a stronger acid than the dimethylammonium ion and the methylammonium ion and is almost as strong as the ammonium ion itself, which implies that trimethylamine is a weaker base than dimethylamine or methylamine. There seems to be no continuous trend in these data.

Indeed there isn’t, and it takes a look at matters in the gas phase to straighten things out. In the gas phase, the trend is regular. The basicity of amines increases with substitution, and the acidity of ammonium ions decreases with substitution (Fig. 6.37).

There must be some effect present in solution that disappears in the gas phase, and this effect must be responsible for the irregularities in Table 6.9. The culprit is the solvent itself. Ions in solution are strongly stabilized by solvation, by interaction of the solvent molecules with the ion. These interactions can take the form of electrostatic stabilization or of partial covalent bond formation, as in hydrogen bonding (Fig. 6.38).

In solution, an alkyl group has two effects on the stability of an ammonium ion. It stabilizes by helping the ammonium ion to disperse the charge, but it destabilizes the ion by interfering with solvation by making intermolecular solvation (charge dispersal) more difficult. These two effects operate in opposite directions. Little steric interference with solvation occurs when we replace one hydrogen of the ammonium ion with a methyl group, and the pK_a change is more than a full pK_a unit. This change reflects the stabilizing effect of the methyl group on the ammonium ion. However, when a second hydrogen is replaced by a methyl, the stabilization and interference with solvation nearly balance. The result is practically no change in the overall stability of the ammonium ion. When a third hydrogen is replaced with methyl, the destabilizing effects outweigh the stabilizing forces, and the result is a less stable, more acidic ammonium ion. In the gas phase, where there can be no solvation, only the stabilizing effects remain, and each replacement of a hydrogen with a methyl is stabilizing.

In addition, as we progress along the series ⁺NH₄, ⁺NRH₃, ⁺NR₂H₂, ⁺NR₃H, ⁺NR₄, there are fewer hydrogens available for hydrogen bonding and therefore less stabilization of the ammonium ion in solution.

It must be admitted that this traditional discussion of base strength of amines in terms of the pK_a of the related ammonium ion is indirect. Another, less common but undeniably more direct method is to frame the discussion in terms of the pK_b (−log K_b), where K_b is the basicity constant of the amine (Fig. 6.39).

The higher the pK_b, the weaker is the base, just as the higher the pK_a, the weaker is the acid. Typical amines have pK_b values of about 4, which makes them quite strong bases compared to many other organic molecules.

Amines are much stronger bases than alcohols are. For example, it is much harder to deprotonate ⁺NH₄ than ⁺OH₃. The pK_a value for the ammonium ion (9.2) is much higher than that for the oxonium ion (–1.7) (Fig. 6.40).

 

Why is the oxonium ion much more acidic than the ammonium ion? Oxygen is a more electronegative atom than nitrogen and bears the positive charge less well. We can see the results of the increased stability of the ammonium ion in practical terms. Stable ammonium salts are common, but salts of oxonium ions, though known, are rare and usually unstable (Fig. 6.40).

Primary and secondary amines are Brønsted acids, but only very weak ones. Amines are much weaker acids than alcohols. Removal of a proton from an alcohol gives an alkoxide in which the negative charge is borne by oxygen. An amine forms an amide² in which the less electronegative nitrogen carries the negative charge (Fig. 6.41).

 

It takes a very strong base indeed to remove a proton from an amine to give the amide ion, which is itself a very strong base. Alkyllithium reagents (p. 257) are typically used (Fig. 6.42).

So, amines are good bases but relatively poor acids. Amide ions, the conjugate bases of amines, are even stronger bases than the amines themselves (Fig. 6.42).

 

²Be careful—there is another kind of amide that has the structure R―CO―NH₂. There is no written or verbal distinction made, so you need to know the context before you know which kind of amide is meant. We will use “amide ion” when referring to ⁻NR₂ to help avoid confusion.