Jones M., Fleming S.A. Organic Chemistry

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6.5 Solvents in Organic Chemistry 239

aprotic; hydrocarbons are obvious examples you know. Table 6.7 collects a few

common polar solvents and their descriptions.

TABLE 6.7 Some Common Polar Solvents

Dielectric

Compound Name Constant (ε) Description

O Water 78 Protic, polar

HCOOH Formic acid 59 Protic, polar

OH Methyl alcohol 33 Protic, polar

OH Ethyl alcohol 25 Protic, polar

(CH

)

SO Dimethyl sulfoxide (DMSO) 47 Aprotic, polar

CN Acetonitrile 38 Aprotic, polar

(CH

)

NCHO Dimethylformamide (DMF) 37 Aprotic, polar

Nitromethane 36 Aprotic, polar

6.5b Solubility: “Like Dissolves Like” Oil spreads on water, but very little

dissolves.Two solids, salt (the prototypal ionic solid sodium chloride) and the sugar

sucrose (C

, an organic material bristling with OH groups), both dissolve

completely in water (Fig. 6.24). What is going on?

Sodium chloride, Na





, dissolves in the polar solvent water because both the

positive sodium ion and the negative chloride ion are well solvated by the protic,

polar solvent water (Fig. 6.25). Nonpolar solvents cannot solvate ions efficiently and

do not dissolve sodium chloride.

Sodium chloride

+–

Sucrose

HOCH

FIGURE 6.24 Two very different

solids that dissolve easily in water.

(Only two

hydrogen

bonds shown)

–

FIGURE 6.25 Solvation of the positive

sodium cation, and hydrogen

bonding to the negative chloride ion

(only two hydrogen bonds shown).

Protic solvents are particularly good at dissolving molecules capable of hydro-

gen bonding. Formation of hydrogen bonds in solution is highly stabilizing. Thus

sucrose, a molecule with many OH groups, dissolves easily in water. The polar sol-

vent water stabilizes sucrose both through formation of many hydrogen bonds and

through stabilization of the polarized groups of sucrose. Figure 6.26 shows

a schematic picture of hydrogen bonding (only one of many shown) and dipolar

stabilization of the polyhydroxy compound sucrose by water.

Sucrose

Dipole–dipole

stabilization

Hydrogen

bonding

–

FIGURE 6.26 Solvation of an

bond by water.

240 CHAPTER 6 Alkyl Halides, Alcohols, Amines, Ethers, and Their Sulfur-Containing Relatives

Polar aprotic solvents can help disperse the partial charges in other polar apro-

tic molecules and thus dissolve them well. Although they cannot act as proton

donors, such solvents can align their negative ends toward the positive charges in a

polar molecule and their positive ends toward the negative charges (Fig. 6.27). For

example, Figure 6.27 shows the polar solvent THF solvating a water molecule by

aligning dipoles in a stabilizing way. Water and THF are miscible.

By contrast, hydrocarbons cannot form hydrogen bonds (they are neither proton

donors nor acceptors) and are not very polar. Typical ε values for hydrocarbons are

around 2.The more nonpolar and “greasy”the solvent,the more easily the similarly non-

polar hydrocarbon molecules dissolve in it.Thus,hydrocarbons (oil is a mixture of many

hydrocarbons) are almost completely insoluble in water,but very soluble in other hydro-

carbons.The polar molecule acetone [dimethyl ketone, (CH

)

CO],with an ε value of

21, is completely soluble in water, but the alkene isobutene, with the same number of

nonhydrogen atoms, is almost completely insoluble in water. Like does dissolve like.

THF

–

FIGURE 6.27 Stabilization (solvation)

of water by tetrahydrofuran, THF.

Summary

In the practical world of making chemical reactions work, we need to allow

molecules to “find each other”and react. Dissolving molecules in a fluid medium

provides an opportunity for molecules to move around and do just that—find

other reactive molecules. The choice of solvent is critical because we need to

dissolve the molecules we want to react together. In making that choice, the

rather simple notion that “like dissolves like” really works! As we work through

chemical reactions in the next few chapters, we will have to take account of

solvents and solvation.

6.6 Diols (Glycols)

Although molecules containing two OH groups are logically enough called diols,

there is an often-used common name as well, glycol.

Some 1,2-diols or 1,2-glycols are familiar and are often important molecules.

1,2-Ethanediol, or ethylene glycol, is the commonly used antifreeze, for example.

Note that in naming these compounds, the final “e”of the alkane parent compound

is not dropped, as it is in naming simple alcohols, and that the almost universally

used common name can be very misleading. Ethylene glycol implies the presence of

an alkene, and propylene glycol (1,2-propanediol) suggests a three-carbon unsaturat-

ed chain. In neither case is the implied unsaturation present (Fig. 6.28). Be aware

that the real structures are based on saturated hydrocarbon chains.

HO OH

1,2-Ethanediol

or ethylene glycol

1,2-Propanediol

or propylene glycol

WEB 3D

FIGURE 6.28 1,2-Ethanediol and

1,2-propanediol are called ethylene

glycol and propylene glycol,

respectively.

6.7 Amines

6.7a Nomenclature We first saw amines in Chapter 1, albeit very briefly.

Amines are derivatives of ammonia,NH

(Fig. 6.29). Successive replacement

of the hydrogens of ammonia leads to primary, secondary, and tertiary amines,

NRH

,NR

H, and NR

. Quaternary nitrogen compounds are positively

charged and are called ammonium ions. Many systems for naming amines exist,

and the chemical world seems to be resisting attempts to bring order out of this

minichaos by keeping to the old common names.

:::

R NHR

mmonia Primary amine Secondary amine Tertiary amine Ammonium ion

FIGURE 6.29 Substituted amines.

6.7 Amines 241

Primary amines are commonly named by using the name of the substituent, R,

and appending the suffix “amine.” Secondary and tertiary amines in which the R

groups are all the same are simply named as di- and trialkylamines. Amines

bearing different R groups are named by ordering the groups alphabetically.

Ethylmethylamine is correct, methylethylamine is not, although both are surely

intelligible (Fig. 6.30).

Methylamine

(CH

)

CNH

tert-Butylamine

Propylamine

Cyclohexylamine

Dicyclopentylamine

Ethylamine

PhCH

Benzylamine

(CH

)

Dimethylamine

Benzylethyl-

methylamine

(CH

)

Triethylamine

Ethylmethylamine,

(not methylethylamine)

PhCH

WEB 3D

FIGURE 6.30 A naming scheme for primary, secondary, and tertiary amines.

The IUPAC system names amines analogously to alcohols. If CH

OH is off-

cially known as methanol

, then CH

is methanamine. Figure 6.31 gives two

examples.

Still another method names amines as substituted alkanes. Now CH

becomes aminomethane. Figure 6.32 gives some examples of amines named

this way.

Methanamine

Cyclohexanamine

FIGURE 6.31 Two systematically

named amines.

minomethane

Aminocyclohexan

1-Aminopropane

2-Aminopropane

FIGURE 6.32 Still another method for naming amines.

The amino group is just below the alcohol group in the naming protocol prior-

ity system. Figure 6.33 gives some correct and incorrect names of polyfunctional

compounds containing amino groups.

2-Methyl-1-propanamine

(1-amino-2-methylpropane)

3-Chloro-2-butanamine

(not 3-amino-2-chlorobutane)

3-Amino-2-butanol

(or 3-amino-2-hydroxybutane,

not 2-amino-3-butanol

or 2-amino-3-hydroxybutane)

3-Bromocyclopentanamine

(neither 3-aminobromocyclopentane

nor 4-bromoaminocyclopentane)

3-Aminocyclopentanol

(not 3-hydroxycyclopentylamine)

FIGURE 6.33 Some polyfunctional

amines and their names.

242 CHAPTER 6 Alkyl Halides, Alcohols, Amines, Ethers, and Their Sulfur-Containing Relatives

Simple, cyclic amino compounds are also widely seen in nature, and are almost

invariably known by their common names. These compounds are often named as

“aza” analogues of an all-carbon system. Thus aziridine becomes azacyclopropane.

Methylamine could be called azaethane, although in point of fact this system is

almost never used for acyclic compounds.Unless there is an oxygen atom in the ring,

the numbering scheme places the nitrogen at 1 and proceeds so as to minimize the

numbering of any substituents. Figure 6.34 gives a few important examples. Once

again there is no way out other than to learn these names.

WEB 3D

H H

Aziridine

(azacyclopropane)

Morpholine 3-Methylpiperidine

(

not 5-meth

eridine

)

3-Bromopyrrolidine

(

not 4-bromo

rrolidine

)

4-Benzylmorpholine

(

N-benz

lmor

holine

)

Azetidine

(azacyclobutane)

Pyrrolidine

(azacyclopentane)

Piperidine

(azacyclohexane)

Pyridine

(azabenzene)

Pyrrole

FIGURE 6.34 Cyclic amines

and their names.

Despite this seeming excess of nomenclature, another aspect of naming amines

must be mentioned.Substituted amines can also be named as N-substituted amines.

In this system, the substituent on nitrogen is designated with the locating prefix N.

Some examples are shown in Figure 6.35.

-Benzylaziridine N-Ethylpiperidine N-Methylazacyclooctane

(heptamethylenemethylamine)

FIGURE 6.35 Still another naming

protocol for amines.

WEB 3D

Tetraethylammonium bromide Butylethylmethylpropylammonium iodide

N(CH

)

–

Ethylmethylammonium chloride

–

FIGURE 6.36 Names for three ammonium ions.

Ammonium ions are named by alphabetically attaching the names of the substituents.

Don’t forget to append the name of the negatively charged counterion (Fig. 6.36).

6.7 Amines 243

6.7b Structure and Physical Properties of Amines In simple amines, the

carbon–nitrogen bond distance is a little shorter than the corresponding

carbon–carbon bond distance in alkanes.This phenomenon is similar to that found

in the alcohols and ethers (Section 6.8) (Fig. 6.37).

COH

Alcohol Amine Alkane

CNH

CCH

1.43 A

⬚

1.47 A

⬚

1.54 A

⬚

FIGURE 6.37 The carbon–nitrogen

bond in amines is shorter than the

carbon–carbon bond in alkanes.

PROBLEM 6.11 Why is the carbon–nitrogen bond distance in aniline slightly

shorter than that in methylamine?

1.47 A

⬚

1.40 A

⬚

Aniline

Methylamine

Simple amines are hybridized approximately sp

at nitrogen and thus are

pyramidal. The bond angles are close to the tetrahedral angle of 109.5° but, of

course, cannot be exactly the tetrahedral angle.

PROBLEM 6.12 Why not?

For amines in which all three R groups are identical, there will be only a single

angle,but for less symmetrical amines,the different angles

cannot be the same (Fig. 6.38).

The most important structural feature of amines comes from the presence of the

pair of nonbonding electrons acting as the fourth substituent on the pyramid. In

alkanes, where there are four substituents, there is nothing more to consider once

the pyramid is described. In amines there is, because the pyramid can undergo an

“umbrella flip,”amine inversion, forming a new, mirror-image pyramid (Fig. 6.39).

108.4⬚

112.9⬚

105.9⬚

107.3⬚

FIGURE 6.38 Some bond angles in

simple amines.

A carbon-based pyramid is static

The nitrogen-based pyramid can invert

Inversion produces the

mirror-image pyramid

Mirror

.. ..

FIGURE 6.39 Amine inversion

interconverts enantiomeric amines.

This inversion process is sometimes confusing when first encountered,so let’s work

through it, making comparisons with the alkanes as we go along. At the beginning

of the process,the pyramid begins to flatten and the bond angles increase.R

244 CHAPTER 6 Alkyl Halides, Alcohols, Amines, Ethers, and Their Sulfur-Containing Relatives

At the halfway point along the path to the inverted pyramid, the transition state is

reached. Here the molecule is planar, with 120° bond angles. At this

point, the hybridization of the central nitrogen is sp

! The inversion process main-

tains a mirror plane of symmetry. In the alkanes, we could certainly imagine the

beginning of the flattening process,but there is no way to force the fourth substituent

through the molecule to the other side (Fig. 6.40).

Inversion progress

This flattening

will cost energy

This process can't

go further

Planar transition state

N is

FIGURE 6.40 The transition state for

amine inversion is planar.

One enantiomer The other enantiomer

ΔG

†

~ 5 kcal/mol

FIGURE 6.41 Rapid inversion

racemizes optically active amines.

PROBLEM 6.13 In general terms, what is the hybridization of nitrogen at some

point between the starting pyramidal amine and the transition state?

An important concern is the rate of this inversion process. If it is slow at room

temperature, then we should be able to isolate appropriately substituted optically

active amines. If inversion is fast, however, racemization will also be fast as the

inversion process interconverts enantiomers. In practice, this is what happens.

Barriers to inversion in simple amines are very low, about 5–6 kcal/mol, and

isolation of one enantiomer, even at very low temperature, is difficult. Inversion

is simply too facile, and the rapid inversion interconverts enantiomeric amines

(Fig. 6.41).

PROBLEM 6.14 In contrast to simple amines, aziridines (three-membered

rings containing a nitrogen) can often be separated into enantiomers.

For example, the activation energy (ΔG

‡

) for the inversion of 1,2,2-

trimethylaziridine is about 18.5 kcal/mol,much higher than that for simple amines.

Explain.

(continued)

6.7 Amines 245

When the methyl on nitrogen is replaced

by phenyl, ΔG

†

decreases to 11.2 kcal/mol

Ph =

The barrier to this inversion is unusually high, ΔG

†

= 18.5 kcal/mol

PROBLEM 6.15 If a phenyl group (Ph in the figure above) is attached to the

nitrogen atom of the aziridine, the barrier to inversion decreases (see figure in

Problem 6.14). Explain.

Like alcohols, amines are much higher boiling than hydrocarbons of similar

molecular weight, although the effect is not as great as with alcohols (Table 6.8;

compare to Table 6.3). The reasons for the increased boiling points are the same as

TABLE 6.8 Some Physical Properties of Amines and a Few Related Alkanes

Compound bp (°C) mp (°C)

Ammonia

–33.4 –77.7

–6.3 –93.5

–88.6 –183.3

Ethane

16.6 –81

Density (g/cm

)

0.77

0.70

0.57

0.68

Ethylamine

Methylamine

–189.7

–93

–43.1

7.4

Dimethylamine

(CH

)

N 2.9 –117.2

184 –6.3

0.59

0.68

0.64

1.02

Propane

Trimethylamine

Aniline

(CH

)

WEB 3D

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 increas-

ing the boiling point (Fig. 6.42).

Association because of favorable

dipole alignment

Oligomer formation through

dro

en bondin

FIGURE 6.42 The effective molecular

weight of amines is increased through

association and hydrogen bonding.

246 CHAPTER 6 Alkyl Halides, Alcohols, Amines, Ethers, and Their Sulfur-Containing Relatives

When hydrogen bonding is impossible, as with tertiary amines, boiling points

decrease (Table 6.8).

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.

PROBLEM 6.16 Many people use lemon when eating fish. This custom is a carry-

over from the days when it was difficult to preserve fish, and the lemon acted to

diminish the unpleasant odor (if not the decomposition). Lemons contain 5–8%

citric acid, and this acidity contributes to their sour taste. Explain why lemon juice

is effective at reducing the odor of fish.

6.7c Acid and Base Properties of Amines Like water and alcohols, amines

are Brønsted bases (proton acceptors). Primary and secondary amines are weak

Brønsted acids (proton donors) as well. Brønsted basicity results in the formation

of ammonium ions through proton transfer (Fig. 6.43).

–

Ammonium

ion

mmonia

FIGURE 6.43 Ammonia acting as a

Brønsted base in proton transfer.

A weakly basic amine is a less effective competitor for

and produces an ammonium ion that is a relatively

strong Brønsted acid (low pK

)

–

… an ammonium ion that is a relatively

weak Brønsted acid (high pK

)

A strongly basic amine is a good

competitor for H

and produces…

FIGURE 6.44 Ammonium ions with

high pK

values are related to

strongly basic amines, and

ammonium ions with low pK

values

are related to weakly basic amines.

The better competitor an amine is in the proton-transfer reaction, the stronger

is its Brønsted basicity. A way of turning this statement around is to focus not on

the Brønsted basicity of the amine, but on the Brønsted acidity of its conjugate acid,

the ammonium ion. If the ammonium ion is a strong Brønsted acid, the related

amine must be a weak Brønsted 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

values. So the pK

of the ammonium ion

has come to be a common measure of the basicity of the related amine.Ammonium

ions with high pK

values (weak acids) are related to strongly basic amines, and

ammonium ions with low pK

values (strong acids) are related to weakly basic

amines (Fig. 6.44).

6.7 Amines 247

Table 6.9 relates the basicity of ammonia and some very simple alkylamines by

tabulating the pK

values of the related ammonium ions.

TABLE 6.9 Some pK

Values for Simple Ammonium Ions in Solution

Amine Formula Ammonium Ion pK

(in aqueous solution)

Ammonia NH



9.24

Methylamine CH

10.63

Dimethylamine (CH

)

NH 10.78

Trimethylamine (CH

)

N9.80(CH

)

(CH

)

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. 138): 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 ammo-

nium ion, the less readily it loses a proton and the higher its pK

. 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

units) than between methylamine and dimethylamine (only 0.15 pK

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 almost as strong as the ammonium ion itself,

which implies that trimethylamine is a weaker base than dimethylamine or methyl-

amine. 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 sub-

stitution, and the acidity of ammonium ions decreases with substitution (Fig. 6.45).

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 elec-

trostatic stabilization or of partial covalent bond formation, as in hydrogen bonding

(Fig. 6.46).

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! There is

little steric interference with solvation when we replace one hydrogen of the ammo-

nium ion with a methyl group, and the pK

change is more than a full pK

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.

Order of increasing gas-phase pK

Lowest pK

Highest pK

NRH

FIGURE 6.45 The gas-phase acidity

of ammonium ions.

Stabilization through

dipole–dipole

interaction with a

polar solvent

Stabilization

through hydrogen

bonding

FIGURE 6.46 Stabilization of amines

in solution.

248 CHAPTER 6 Alkyl Halides, Alcohols, Amines, Ethers, and Their Sulfur-Containing Relatives

PROBLEM 6.17 Amines can be stabilized by other factors. Explain the following

data.

4.6pK

10.6

It must be admitted that this traditional discussion of base strength of amines

in terms of the pK

of the related ammonium ion is indirect. There is another,

less common but undeniably more direct method. That is to frame the discus-

sion in terms of the pK

(log K

), where K

is the basicity constant of the amine

(Fig. 6.47).

For

RNH

[RNH

] [HO

–

]

[RNH

]

–

FIGURE 6.47 The expression for pK

The higher the pK

, the weaker is the base, just as the higher the pK

, the weaker

is the acid. Typical amines have pK

values of about 4, which makes them quite

strong bases.

Amines are much stronger bases than alcohols are. For example, it is much

harder to deprotonate



than



.The pK

value for the ammonium ion (9.2)

is much higher than that for the oxonium ion (–1.7) (Fig. 6.48).

= 9.2

N(CH

)

Ammonium ion

Tetramethylammonium

chloride, a stable solid,

easily isolable

Trimethyloxonium

fluoborate, not easily

isolated, and very reactiv

= –1.7

Oxonium ion

–

O(CH

)

–

FIGURE 6.48 Ammonium ions are

weaker acids than oxonium ions.

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. 6.48).

In addition, as we progress along the series



NRH



, there are fewer hydrogens available for hydrogen bonding and therefore less

stabilization of the ammonium ion in solution.