Thus far, we have dealt with Brứnsted-Lowry bases that have only one site that can act as a proton acceptor in an acid-base reaction. Many organic compounds have two or more such sites. In the following discussion, we restrict our attention to compounds containing a carbonyl group in which the carbonyl carbon is bonded to either an oxygen or a nitrogen. The principle we develop here is an extremely important one and applies to other types of molecules as well. The more stable charged species is the one in which the charge is more delocalized. Relative charge delocalization can often be understood by considering resonance (Section 1.8).
Let us consider fi rst the potential sites for proton transfer to an oxygen atom of a carboxylic acid such as acetic acid. Proton transfer to the carbonyl oxygen gives cation A, and proton transfer to the hydroxyl oxygen gives cation B.
O O
CH3 C H H2SO4 CH3 O H or H
C O
HSO4
A
(protonation on the carbonyl oxygen)
O CH3 C H
B
(protonation on the hydroxyl oxygen)
H O
We now examine each cation and determine which is the more stable (lower in energy). For cation A, we can write three contributing structures. Two of these place the positive charge on oxygen, and one places it on carbon.
O
CH3 H
H C O
A-1 (C and O have complete octets)
A-2 (C has incomplete
octet)
A-3 (C and O have complete octets)
O
CH3 H
H C O
O
CH3 H
H C O
Of these three structures, A-1 and A-3 make the greater contributions to the hybrid because all atoms in each have complete octets; A-2 makes a lesser contribution because its carbonyl carbon has an incomplete octet. Thus, on protonation of the carbonyl oxygen, the positive charge is delocalized over three atoms with the greater share of it being on the two equivalent oxygen atoms. (The two oxygens were not equivalent before protonation, but they are now.)
Protonation on the hydroxyl oxygen gives cation B for which we can write two resonance contributing structures.
O O
B-1 B-2
(charge separation and adjacent positive charges) O
CH3 C H H
O H
CH3 C H
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158 Chapter 4 Acids and Bases
Of these, B-2 makes at best only a minor contribution to the hybrid because of the adjacent positive charges; therefore, the charge on this cation is, in effect, localized on the hydroxyl oxygen.
From this analysis of cations A and B, we see that protonation of a carboxyl group occurs preferentially on the carbonyl oxygen because this cation has greater delocalization of the positive charge.
Example 4.3
The functional group created when the !OH of a carboxyl group is replaced by an NH2 group is called an amide (Section 1.3F). Draw the structural formula of acet- amide, which is derived from acetic acid, and determine if proton transfer to the amide group from HCl occurs preferentially on the amide oxygen or the amide nitrogen.
Solution
Following is a Lewis structure for acetamide and its two possible protonated forms.
Acetamide (an amide)
A
(protonation on the amide oxygen)
B
(protonation on the amide nitrogen) N
CH3 C H HCl or
H
N
CH3 C H
H H
H O
Cl N
CH3 C H
H
O O
Of the three contributing structures that can be drawn for cation A, structures A-1 and A-3 make the greater contributions to the hybrid because all atoms in each have complete octets; of these two contributors, A-3 has the positive charge on the less electronegative atom and, therefore, makes a greater contribution than A-1.
The result is that the positive charge in cation A is delocalized over three atoms, the greater share of it being on nitrogen and oxygen.
A-1 (complete octets;
positive charge on more electronegative atom)
A-2 (incomplete
octet for C) N
CH3 C H H H
O
N CH3 C H
H H
O
A-3 (complete octets;
positive charge on less electronegative atom)
N CH3 C H
H H
O
Only two contributing structures can be drawn for cation B. Of these, B-2 requires creation and separation of unlike charges and places positive charges on adjacent atoms. It therefore makes little contribution to the hybrid. Thus, the positive charge in cation B is essentially localized on the amide nitrogen.
B-1 (protonation on the
amide nitrogen)
B-2
(adjacent plus charges;
charge separation) CH3 C H
H N H O
CH3 H H N H
C
O
From this analysis, we conclude that as in the acetic acid example, proton transfer to the carbonyl oxygen of the amide group gives the more delocalized, and hence, the more stable cation, A.
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4.2 Brứnsted-Lowry Acids and Bases 159 Problem 4.3
Following is a structural formula for guanidine, the compound by which migratory birds excrete excess metabolic nitrogen. The hydrochloride salt of this compound is a white crystalline powder, freely soluble in water and ethanol.
(a) Write a Lewis structure for guanidine showing all valence electrons.
(b) Does proton transfer to guanidine occur preferentially to one of its !NH2 groups (cation A) or to its "NH group (cation B)? Explain.
NH
H2N9C9NH2 1 HCl Guanidine
NH NH21
H2N9C9NH31 or H2N9C9NH2 1 Cl2
A B
C. p Electrons as Brứnsted-Lowry Bases
Thus far we have considered proton transfer to atoms having a nonbonding pair of elec- trons. Proton-transfer reactions also occur with compounds having p electrons (e.g., the p electrons of carbon-carbon double and triple bonds). The p electrons of the carbon- carbon double bond of 2-butene, for example, react with strong acids such as H2SO4, H3PO4, HCl, HBr, and HI by proton transfer to form a new carbon-hydrogen bond.
CH39CH"CH9CH31 H9Br
sec -Butyl cation (a 2° carbocation) 2-Butene
CH3 CH3
H H
C H
1 1
C Br–
The result of this proton-transfer reaction is formation of a carbocation, a species in which one of its carbons has only six electrons in its valence shell and carries a charge of 11. Because the carbon bearing the positive charge in the sec-butyl cation has only two other carbons bonded to it, it is classifi ed as a secondary (2°) carbocation. We will study the formation, structure, and reactions of carbocations in detail in Chapter 6.
Example 4.4
The acid-base reaction between 2-methyl-2-butene and HI can in principle form two carbocations. Write chemical equations for the formation of each carbocation and use curved arrows to show the proton transfer in each reaction.
CH39C"CH9CH3 2-Methyl-2-butene
CH3
Solution
Proton transfer to carbon 3 of this alkene gives a tertiary (3°) carbocation. Proton transfer to carbon 2 gives a secondary (2°) carbocation.
CH39C"CH9CH3
H3C9C9C9CH3
CH3
H
H I
H A 3° carbocation
H3C
I
H3C9C9C9CH3 H
H A 2° carbocation
H3C
I
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160 Chapter 4 Acids and Bases
Problem 4.4
Write an equation to show the proton transfer between each alkene or cycloalkene and HCl. Where two carbocations are possible, show each.
(a) CH3CH2CH"CHCH3 (b)
Cyclohexene 2-Pentene
4.3 Acid Dissociation Constants, pKa, and the Relative Strengths of Acids and Bases
Any quantitative measure of the acidity of organic acids or bases involves mea- suring the equilibrium concentrations of the various components in an acid-base equilibrium. The strength of an acid is then expressed by an equilibrium constant.
The dissociation (ionization) of acetic acid in water is given by the following equation.
CH3COH
Acetic acid Water Acetate ion
O
H2O CH3CO– O
Hydronium ion H3O
We can write an equilibrium expression for the dissociation of this or any other uncharged acid in a more general form; dissociation of the acid, HA, in water gives an anion, A2, and the hydronium ion, H3O1. The equilibrium constant for this ionization is
HA1H2O mA21H3O1 Keq5 3H3O14 3A24
3HA4 3H2O4
Because water is the solvent for this reaction and its concentration changes very little when HA is added to it, we can treat the concentration of water as a constant equal to 1000 g/L or approximately 55.6 mol/L. We can then combine these two constants (K eq and the concentration of water) to defi ne a new constant called an acid dissociation constant, given the symbol K a.
Ka5Keq 3H2O45 3H3O14 3A24
3HA4
Because dissociation constants for most acids, including organic acids, are num- bers with negative exponents, acid strengths are often expressed as pK a (2log10Ka).
The pK a for acetic acid is 4.76, which means that acetic acid is a weak acid be- cause the major species present in aqueous solution is undissociated CH3COOH.
Table 4.1 gives names, molecular formulas, and values of pK a for some organic and inorganic acids. As you study the information in this table, note the following relationships:
pK a 5 2log10K a The larger the value of p
• K a, the weaker the acid.
The smaller the value of p
• K a, the stronger the acid.
The weaker the acid, the stronger its conjugate base.
•
The stronger the acid, the weaker its conjugate base.
•
Acid dissociation constant Equal to the equilibrium constant (Keq) for an acid dissociation reaction multiplied by the concentration of water [H2O].
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4.3 Acid Dissociation Constants, pKa, and the Relative Strengths of Acids and Bases 161
Example 4.5
For each value of pK a, calculate the corresponding value of K a. Which compound is the stronger acid?
(a) Ethanol, pK a 5 15.9 (b) Carbonic acid, pK a 5 6.36 Solution
Recall that pK a 5 2log10K a so calculating K a 5 102pKa values gives:
(a) For ethanol, K a 5 1.3 3 10216 (b) For carbonic acid, K a 5 4.4 3 1027 Because the value of pK a for carbonic acid is smaller than that for ethanol, carbonic acid is the stronger acid, and ethanol is the weaker acid.
Problem 4.5
For each value of K a, calculate the corresponding value of pK a. Which compound is the stronger acid?
(a) Acetic acid, K a 5 1.74 3 1025 (b) Chloroacetic acid, K a 5 1.38 3 1023 The pKa values can also be used to readily estimate the equilibrium constants (Ka). For example, if the pKa of an acid is near zero, then the equilibrium constant for the reaction of that acid protonating water is near 1. Negative pKa values cor- relate to acids with equilibrium constants greater than 1, while positive pKa values are for acids with equilibrium constants less than 1. Each single unit difference between pKa values represents a tenfold increase or decrease in the strength of the acids being compared.
Table 4.1 pKa Values for Some Organic and Inorganic Acids
Acid Formula pK a Conjugate Base
Ethane CH3CH3 51 CH3CH22
Ethylene CH2"CH2 44 CH2”CH2
Ammonia NH3 38 NH22
Hydrogen H2 35 H2
Acetylene HC#CH 25 HC#C2
Ethanol CH3CH2OH 15.9 CH3CH2O2
Water H2O 15.7 HO2
Methylammonium ion CH3NH31 10.64 CH3NH2
Bicarbonate ion HCO32 10.33 CO322
Phenol C6H5OH 9.95 C6H5O2
Ammonium ion NH41 9.24 NH3
Hydrogen sulfi de H2S 7.04 HS2
Carbonic acid H2CO3 6.36 HCO32
Benzoic acid C6H5COOH 4.19 C6H5COO2
Hydrogen fl uoride HF 3.2 F2
Phosphoric acid H3PO4 2.1 H2PO42
p-Toluenesulfonic acid CH3C6H4SO3H 0.7 CH3C6 H4SO32
Nitric acid HNO3 21.5 NO32
Hydronium ion H3O1 21.74 H2O
Sulfuric acid H2SO4 25.2 HSO42
Hydrogen chloride HCl 27 Cl2
Hydrogen bromide HBr 28 Br2
Hydrogen iodide HI 29.9 I2
Stronger conjugate base
Weaker conjugate base Weaker
acid
Stronger acid
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162 Chapter 4 Acids and Bases
Values of pKa in aqueous solution in the range 2 to 12 can be measured quite accurately. Values of pKa smaller than 2 are less accurate because very strong acids, such as HCl, HBr, and HI, are completely ionized in water, and the only acid present in solutions of these acids is H3O1. For acids too strong to be measured accurately in water, less basic solvents such as acetic acid or mixtures of water and sulfuric acid are used. Although none of the halogen acids, for example, is completely ionized in acetic acid, HI shows a greater degree of ionization in this solvent than either HBr or HCl; therefore, HI is the strongest acid of the three. Values of pKa greater than 12 are also less precise. For bases too strong to be measured in aqueous solution, more basic solvents such as liquid ammonia and dimethyl sulfoxide are used. Because different solvent systems are used to measure relative strengths at either end of the acidity scale, pKa values smaller than 2 and greater than 12 should be used only in a qualita- tive way when comparing them with values in the middle of the scale.
4.4 The Position of Equilibrium in Acid-Base Reactions
We know from the value of pKa for an acid whether an aqueous solution of the acid contains more molecules of the undissociated acid or its ions. A negative pKa value indicates that the majority of molecules of the acid are dissociated in water, while a positive value indicates that most acid molecules remain undissociated in water.
HCl, for example, a strong acid with a pKa of 27, is almost completely dissociated at equilibrium in aqueous solution, and the major species present are H3O1 and Cl2.
HCl 1 H2O 2F H3O1 1 Cl2 pK a 5 27
For acetic acid on the other hand, which is a weak acid with a pK a of 4.76, the major species present at equilibrium in aqueous solution are CH3COOH molecules.
O