clayden ch08 acids and bases -...

Chapter 8

Acidity, Basicity and pKa

p182

In this reaction water is acting as a base, according to our definition above, by accepting a proton from HCl which in turn is acting as an acid by donating a proton. If we consider the reverse reaction, the chloride is acting as a base and the hydronium ion as an acid. The chloride ion is called the conjugate base of hydrochloric acid and the hydronium ion, H3O+, is the conjugate acid of water.

Every acid has a conjugate base associated with it and every base has a conjugate acid associated with it.

p185

The lower the pKa, the larger the equilibrium constant, Ka, is and hence the stronger the acid.

The pKa of the acid is the pH where it is exactly half dissociated.

At pHs above the pKa, the acid HA exists as A– in water; at pHs below the pKa, it exists as undissociated HA. €

pH = pKa + logA−[ ]HA[ ]

p187

Adjusting the pH alters the proportions of the acid form and of the conjugate base.

At low pH the compound exists entirely as AH and at high pH entirely as A–. At the pKa the concentration of each species, AH and A–, is the same.

P187-188

An acid’s pKa depends on the stability of its conjugate base

HCl is a much stronger acid than acetic acid: the pKa of HCl is around –7 compared

to 4.76 for acetic acid. This tells us that in solution Ka for hydrogen chloride is 107

mol dm–3 whilst for acetic acid it is only 10–4.76 = 1.74 × 10–5 mol dm–3. Why are the equilibria so different? Why does hydrogen chloride fully dissociate but acetic acid do so only partially?

The answer must have something to do with the conjugate base A– of each acid.

i.e. the chloride ion is fundamentally more stable than is the acetate ion.

p190

In water, our effective pKa range is only –1.74 to 15.74. This is known as the levelling

effect of the solvent. If we want to remove the proton from something with a high pKa, say 25–30, it would be impossible to do this in water since the strongest base we can use is hydroxide. If we do need a stronger base than OH–, we must use a different solvent system. So, no matter what base we dissolve in water, we will only at best get hydroxide ions, this is the best we could do in water.

Using the pKas of NH3 (ca. 33) and ethyne (25) we would predict an equilibrium

constant for this reaction of 108 (10–25/10–33)—well over to the right. Amide ions can be used to deprotonate alkynes.

The choice of solvent limits the pKa range we can use:

• Anything that stabilizes a conjugate base A: – makes the starting acid H – A more acidic.

Five factors affect the acidity of H – A: [1] Element effects [2] Resonance effects [3] Inductive effects [4] Hybridization effects [5] Solvent effects

No matter which factor is discussed, the same procedure is always followed. To compare the acidity of any two acids: • Always draw the conjugate bases. • Determine which conjugate base is more stable. • The more stable the conjugate base, the more acidic the acid.

Factors that affect the acidity of H – A

p208a

[1] Element Effects—Trends in the Periodic Table The most important factor determining the acidity of H – A is the location of A in the periodic table.

Across a row of the periodic table, the acidity of H – A increases as the electronegativity of A increases.

a-Comparing Elements in the Same Row of the Periodic Table:

The pKa values for second row hydrides CH

4, NH

3, H

2O, and HF are about 48,

33, 16, and 3, respectively.

p208a

Element Effects—Trends in the Periodic Table b-Comparing Elements Down a Column of the Periodic Table Positive or negative charge is stabilized when it is spread over a larger volume.

Because Br– is larger than F–, Br– is more stable than F–, and H – Br is a stronger acid than H–F. Down a column of the periodic table, the acidity of H – A increases as the size of A increases.

p182

[2] Resonance Effects A third factor that determines acidity is resonance. • The acidity of H – A increases when the conjugate base A:– is resonance stabilized.

p196a

Becoming aromatic

p197c

The pKa of ammonia is much greater than the pKa

of water (about 33 compared with 15.74). This is because oxygen is more electronegative than nitrogen and so can stabilize the negative charge better. The oxygen equivalent of an amide (a carboxylic acid) has a low pKa. Nevertheless, the carbonyl group of an amide does lower the pKa

from that of an amine (about 30) to around 17. It’s not surprising, therefore, that the two carbonyl groups in an imide lower the pKa

still further, as in the case of phthalimide.

Nitrogen acids:

p208a

A third factor affecting the acidity of H – A is the presence of electronegative atoms. [3] Inductive Effects

• Electron-withdrawing groups stabilize a conjugate base, making a carboxylic acid more acidic. • Electron-donating groups destabilize the conjugate base, making a carboxylic acid less acidic. • The larger the number of electronegative substituents, the stronger the acid. • The more electronegative the substituent, the stronger the acid. • The closer the electron-withdrawing group to the COOH, the stronger the acid.

Figure 4.8 shows the energies of the hybrid orbitals relative to the s and p orbitals from which they are formed.

The unshared pair of electrons of the conjugate base of ethane occupies an sp3 atomic orbital on the carbon. The unshared pair of electrons of the conjugate base of ethene occupies an sp2 atomic orbital. Because the sp2 orbital is lower in energy, the unshared electrons in this orbital in the conjugate base of ethene are more stable (and less basic) than the electrons in the sp3 orbital of the conjugate base of ethane. Thus, ethene is a stronger acid than ethane. Because the electrons of the conjugate base of ethyne are even lower in energy in an sp orbital, ethyne is an even stronger acid.

[4] Hybridization Effects

In this series, as the hybridization changes from sp3 in ethane to sp2 in ethene and to sp in ethyne, the acidity increases and the pKa

decreases. This is because of the relative stability of the unshared electrons in the conjugate bases of each of these compounds.


The higher the percent s-character of the hybrid orbital, the more stable the conjugate base. i.e. the acidity of H – A increases as the percent s-character of the A:– increases.


p182

[5] The solvent

Hydrochloric acid is a strong acid: the free energy ∆G° for its ionization equilibrium in water is –40 kJ mol–1.

Such a large negative ∆G° value means that the equilibrium lies well over to the right. In the gas phase, however, things are drastically different and ∆G° for the ionization is +1347 kJ mol–1.

This ∆G° value corresponds to 1 molecule of HCl in 10240 being dissociated! This means that HCl does not spontaneously ionize in the gas phase—it does not lose protons at all.

P182, 185

It should not be forgotten that the prime requirement of the solvent is that it should be capable of functioning as a base: the weaker the base, the smaller the dissociation of the acid. So HCl is a strong acid in methanol but almost wholly undissociated in toluene.

P182, 185

Why then is HCl such a strong acid in water? In aqueous solution the proton is strongly attached to a water molecule to give the very stable hydronium ion, H3O+, and the ions are no longer isolated but solvated. In fact water can solvate both cations and anions. The effect is particularly marked with anions for H-bonded type solvation can occur. Similar H-bonded type solvation cannot in general occur with cations except with H3O+.

pKa values of some compounds in different solvents at 25 oC

In some cases, solvent effects are invoked to explain small difference in pKa values. For

example, compare the acidity of tert-butanol and ethanol:

The pKa values indicate that tert-butanol is less acidic than ethanol by two orders of

magnitude. In other words, the conjugate base of tert-butanol is less stable than the conjugate base of ethanol. This difference in stability is best explained by considering the interactions between each conjugate base and the surrounding solvent molecules (Figure 3.7). Compare the way in which each conjugate base interacts with solvent molecules. The tert-butoxide ion is very bulky, or sterically hindered, and is less capable of interacting with the solvent. The ethoxide ion is not as sterically hindered so it can accommodate more solvent interactions. As a result, ethoxide is better solvated and is therefore more stable than tert-butoxide (Figure 3.7).

3.7 Solvating Effects

p208a

Substituted Benzoic Acids Rule [1]: Electron-donor groups destabilize a conjugate base, making an acid less acidic.

An electron-donor group destabilizes a conjugate base by donating electron density onto a negatively charged carboxylate anion. A benzoic acid substituted by an electron-donor group has a higher pK

a than benzoic acid (pK

a = 4.2).

p208a

Rule [2]: Electron-withdrawing groups stabilize a conjugate base, making an acid more acidic.

An electron-withdrawing group stabilizes a conjugate base by removing electron density from the negatively charged carboxylate anion. A benzoic acid substituted by an electron-withdrawing group has a lower pK

a than benzoic acid (pK

a = 4.2).

p208a

Figure 19.8 How common substituents affect the reactivity of a benzene ring towards electrophiles and the acidity of substituted benzoic acids

p208a

HOW TO Determine the Relative Acidity of Protons Step [1] Identify the atoms bonded to hydrogen, and use periodic trends to assign relative acidity. • The most common H – A bonds in organic compounds are C – H, N – H, and O – H. Because acidity increases left-to- right across a row, the relative acidity of these bonds is C – H < N – H < O – H. Therefore, H atoms bonded to C atoms are usually less acidic than H atoms bonded to any heteroatom.

Step [2] If the two H atoms in question are bonded to the same element, draw the conjugate bases and look for other points of difference. Ask three questions: • Is the conjugate base resonance stabilized? • Do electron-withdrawing groups stabilize the conjugate base? • How is the conjugate base hybridized?

p198b Basicity

For example, if we want to know which is the stronger base—formate anion or acetylide anion. We look up the pKas for their conjugate acids. We find that the pKa

for formic acid (HCO2H) is 3.7, whilst the pKa

for ethyne (acetylene) is around 25. This means that acetylide is much more basic than formate. If we want to know the basicity of ammonia, we must look up the pKa

of its conjugate acid, the ammonium cation, NH4+, protonated ammonia. Its pKa

is 9.24.

P198-199

There are two main factors that determine the strength of a neutral base: how accessible is the lone pair and to what extent can the resultant positive charge formed be stabilized either by delocalization or by the solvent. Why ammonia is 1010 times more basic than water:

Since oxygen is more electronegative than nitrogen. In other words, the oxygen atom in water wants to keep hold of its electrons more than the nitrogen in ammonia does and is therefore less likely to donate them to a proton. The pKaH

for ammonia (that is, the pKa for ammonium ion) is 9.24

whilst the pKaH for water (the pKa

for hydronium ion) is –1.74. We use pKaH

to mean the pKa of the conjugate acid.

p199

Any substituent that increases the electron density on the nitrogen makes it more basic (larger pKaH). Conversely, any substituent that withdraws electron density from the nitrogen makes it less basic (smaller pKaH).

A strange feature though is that, whilst substituting one hydrogen of ammonia increases the basicity by more than a factor of ten (one pKa unit), substituting two has less effect and in the trisubstituted amine the pKaH

is actually lower.

Ammonia is the simplest nitrogen base and has a pKaH of 9.24.

p199

The observed basicity therefore results from a combination of effects: (1) the increased availability of the lone pair and the stabilization of the resultant positive charge, which increases with successive replacement of hydrogen atoms by alkyl groups; and (2) the stabilization due to solvation, an important part of which is due to hydrogen bonding and this effect decreases with increasing numbers of alkyl groups.

p200b

We should compare these values with typical values of about 11 for simple primary and secondary amines.

p200c

If the lone pair itself is in an sp2 or an sp orbital, it is more tightly held and therefore much harder to protonate. This explains why the lone pair of the nitrile group is not at all basic and needs a strong acid to protonate it.

p201a

The low pKaH of aniline (PhNH2), 4.6, is partly due to the nitrogen being

attached to an sp2 carbon but also because the lone pair can be delocalized into the benzene ring. In order for the lone pair to be fully conjugated with the benzene ring, the nitrogen would have to be sp2 hybridized with the lone pair in the p orbital. This would mean that both hydrogens of the NH2

group would be in the same plane as the benzene ring but this is not found to be the case. Instead, the plane of the NH2

group is about 40° away from the plane of the ring. That the lone pair is partially conjugated into the ring is shown indirectly by NMR shifts and by the chemical reactions that aniline undergoes. Notice that, when protonated, the positive charge cannot be delocalized over the benzene ring.

Account for the Acidity of Protonated Aniline by Resonance Stabilization

p202a

An amidine is the nitrogen equivalent of an amide—a C=NH group replaces the carbonyl. Amidines are much more basic than amides, the pKaHs of amidines are larger than those of amides by about 13 so there is an enormous factor of 1013 in favour of amidines. In fact, they are among the strongest neutral bases.

p202b

An amidine has two nitrogen atoms that could be protonated—one is sp3 hybridized, the other sp2 hybridized. We might expect the sp3 nitrogen to be more basic but protonation occurs at the sp2 nitrogen atom. This happens because we have the same situation as with an amide: only if we protonate on the sp2 nitrogen can the positive charge be delocalized over both nitrogens. The electron density on the sp2 nitrogen in an amidine is increased through conjugation with the sp3 nitrogen.

p202d

Even more basic is guanidine, pKaH 13.6, nearly as strong a base as NaOH! On

protonation, the positive charge can be delocalized over three nitrogen atoms to give a very stable cation. All three nitrogen lone pairs cooperate to donate electrons but protonation occurs, as before, on the sp2 nitrogen atom.

Chapter 16 43

Pyridine

•  Pyridine is basic, with a pair non-bonding electrons available to abstract a proton.

•  The protonated pyridine (the pyridinium ion) is still aromatic.

Pyridine Is Aromatic

p203c

This contrasts to pyrrole in which the lone pair on the only nitrogen atom is needed to complete the six aromatic π electrons and is therefore delocalized around the ring. Protonation, if it occurs at all, occurs on carbon rather than on nitrogen since the cation is then delocalized. But the cation is no longer aromatic (there is a saturated CH2

group interrupting the conjugation) and so pyrrole is not at all basic (pKaH

about –4).

Oxygen bases in general are so much weaker than their nitrogen analogues that we don’t regard them as bases at all and strong acids are needed to protonate them.

clayden ch08 acids and bases -...

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