The concepts of acid and base are central to chemistry as a whole. A vast majority of organic reactions are either promoted by acids or bases or involve the participation of acids or bases at some stage of the reaction pathway. Moreover, the application of acid-base chemistry in or- ganic chemistry is wide reaching; unlike the other subdisciplines of chemistry, an over- whelming majority of organic reactions are carried out in media other than water. This means that the concepts of acids and bases must be applied to a much broader range of conditions than those that applied when most of us learned about acids and bases for the first time. How- ever, let us begin by reviewing the definitions of acids and bases in aqueous solution.
Arrhenius Acids and Bases
The earliest operational definitions of acids and bases were related to the neutralization of acids by bases given by the relatively simple equation that follows:
Acid + Base → Salt + Water (Eq. 1.1)
This simple equation defines neutralization: the destruction of an acid by an equivalent amount of base to give a salt. The formation of the water itself comes from the neutralization process, a situation first recognized by the great Swedish chemist, Svante Arrhenius. As part of his dissociation theory of electrolytes (that when ionic compounds dissolve in water the ions dissociate and become free to move within the solution), Arrhenius looked at electrolytes that were originally nonconductors and proposed the first definition of an acid.
An acid is a substance that functions as a source of hydrogen ions in aqueous solution.
Once acids had been defined in this manner, the extension of this definition to bases was relatively straightforward. The Arrhenius definition of a base is stated below.
A base is a substance that functions as a source of hydroxide ions in aqueous solution.
The strength of an acid is related to the ease with which it functions as a hydrogen ion (or, more accurately, hydronium ion) source in aqueous solution. The functional defini- tions of strong and weak acids are still based largely on the Arrhenius definition. Strong acids ionize (dissociate) completely in aqueous solution. Many are familiar to most stu- dents in chemistry; sulfuric acid (H2SO4), nitric acid (HNO3), hydrochloric acid (HCl), hydrobromic acid (HBr), hydriodic acid (HI), and perchloric acid (HClO4) are among the more familiar strong acids. All these acids are covalent when pure but dissolve in water to give a strongly conducting solution containing hydronium (H3O+) ions. The pH of 1 M solutions of these acids is typically quite low—close to zero. Weak acids, on the other hand, do not dissolve to give complete dissociation. Instead, much of a weak acid remains unionized in aqueous solution. Typical weak acids are compounds such as acetic acid (CH3COOH), which ionizes only to the extent of a few percent in a 1 M solution. The pH of a 1 M solution of a weak acid is much more likely to be close to 2 or 3, indicating a hydronium ion concentration only a few percent of that of a strong acid.
In a similar vein, the strength of a base is related to the ease with which it functions as a source of hydroxide ions in aqueous solution; strong bases are good sources of hydroxide ions in aqueous solution and weak bases are not. Most of the strong bases commonly used in aqueous solution are ionic hydroxides that already have the hydroxide ion present, and so need only dissolve to dissociate and liberate their hydroxide ion. The ionic hydroxides of the Group IA and heavier Group IIA metals, especially those of the Group IA metals, dissolve freely in water to give solutions containing high concentrations of hydroxide ions.
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The substances LiOH, NaOH, and KOH are all widely used as strong bases. Ammonia, on the other hand, dissolves extremely well in water, but it does so without extensive ioniza- tion. Ammonia solutions, often labeled as “ammonium hydroxide,” are actually best de- scribed as ammonia molecules dissolved in water, with only a few percent being converted to ammonium and hydroxide ions. Like the aqueous solutions of weak acids, aqueous solutions of weak bases such as ammonia are relatively poor electrolytes.
There are two structural features that one can use to predict the probable strength of an acid: the atom to which the acidic hydrogen is bonded and the number of multiple bonds between the central atom and oxygen in oxyacids. In general, as the electronegativity of the element bonded to the acidic hydrogen atom increases, acid strength increases. For exam- ple, O–H hydrogens are more acidic than N–H hydrogens, which are in turn more acidic than C–H hydrogens. Acid strength also increases dramatically as the element to which the hydrogen is bonded becomes larger. Although hydrogen fluoride is a relatively weak acid (pKa ≈ 2), hydrogen iodide is the strongest binary (two-element) acid known. In addition, the effects of multiple bonds between the central atom and oxygen atoms appear to be cu- mulative. For example, hypochlorous acid (HClO), which has no chlorine-oxygen double bonds, is a very weak acid, and perchloric acid (HClO4), which has three oxygen atoms doubly bonded to the central chlorine atom, is one of the strongest acids known.
Lowry-Brứnsted Acids and Bases
Not all chemistry is carried out in aqueous solution, and not all acid-base reactions result in the formation of water. There are, in fact, many cases where an acid-base reaction has obviously occurred but where no water is formed. For example, hydrogen chloride gas reacts with ammonia gas to form ammonium chloride, a salt. Because hydroxide ion is never involved in this reaction, it does not fit well with the Arrhenius definition of an acid-base reaction. And yet this reaction clearly leads to the neutralization of the hydrogen chloride, the acid, and the ammonia, the base, leading to the formation of the salt. A more widely applicable definition of acids and bases is that due to Lowry19 and Brứnsted,20 who extended the Arrhenius definition of acids and bases to cover just such situations.
In the Lowry-Brứnsted formalism, an acid is a hydrogen ion donor; a base is a hydrogen ion acceptor.
A simple example of an acid-base reaction by the Lowry-Brứnsted definition is the protonation of the methanol molecule by hydrogen chloride to give the methyloxonium ion (Figure 1.2). The oxygen atom of the methanol molecule is functioning as the proton
19. Thomas Martin Lowry (1874-1936) took his DSc under H. E. Armstrong in 1899, and he remained with him until 1913. His career took him through positions at Westminster Training College and Guy’s Hospital Medical School, to the inaugural Chair in Physical Chemistry at Cambridge. Lowry, who wrote several books, was elected a Fellow of the Royal Society in 1914. For more biographical information, see: Pope, W.J. Obit. No- tices Fellows Roy. Soc. 1938, 2, 287.
20. Johannes Nicholaus Brứnsted (1879-1947) was educated at Copenhagen (PhD, 1908), where he remained as Professor of Physical and Inorganic Chemistry. His principal research was in thermodynamics, especially the strengths and catalytic properties of acids and bases. Throughout World War II, he firmly opposed the Nazis, a stand that won him election to the Danish Parliament in 1947. Unfortunately, by the time of his election, he was too ill to take his seat.
H3C O H H Cl
H3CO H
H Cl
base
acid
conjugate acid
conjugate base Figure 1.2 The transfer of a
proton to a proton acceptor is an acid-base reaction under the Lowry-Brứnsted definition. Here, the proton is transferred from a mole- cule of hydrogen chloride to a molecule of methanol to generate its conjugate acid, methyloxonium ion.
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acceptor, and is therefore the base. The hydrogen chloride molecule is the proton donor, and is therefore the acid. When a proton is removed from an acid, HA, the anion formed, A–, is termed the conjugate base of the acid. When a proton is added to a base, the prod- uct of the reaction is termed the conjugate acid of the base. Thus, during the course of the reaction in Figure 1.2, the acid (hydrogen chloride) is converted to its conjugate base (chloride ion), while the base (methanol) is converted to its conjugate acid (methyloxo- nium ion).
Because of the frequency and importance of proton-transfer reactions in organic chemistry, some familiarity with the acidity constant, Ka, is required. The acidity constant of a Lowry-Brứnsted acid was originally defined in terms of its ability to form hydronium ion, H3O+, from water and is given in Equation 1.3:
HA + H2O → H3O+ + A– (Eq. 1.2)
Ka [H O ][A ]3
5 [HA]
1 2
(Eq. 1.3)
However, with the extension of acid-base chemistry into nonaqueous reaction sol- vents, this definition is rather too narrow, and we are now in a position where Equation 1.1 is too limiting. Nevertheless, we have retained the concept of the acidity constant, Ka, and its negative log10, the pKa, as useful measures of the relative acid strength of a species in solution. Many of these values for organic compounds, however, have been measured indirectly—by competition reactions in a nonaqueous solvent, for example.
As the acid (or base) becomes stronger, its conjugate base (or acid) becomes weaker.
Thus, for a strong acid such as hydrogen chloride (or hydrochloric acid), the conjugate base (chloride anion) is weak, whereas for a weak acid, such as water, the conjugate base (hydroxide anion) is strong. All acid-base reactions proceed to give the weaker of the two conjugate acids and the weaker of the two conjugate bases as the products of the reaction.
Let us examine the case of a simple proton-transfer reaction between an acid, HA, and a base, B–:
HA + B– → HB + A– (Eq. 1.4)
The equilibrium constant for this reaction is given by Equation 1.5.
Keq [HB][A ] [HA][B ] 5
2
2 (Eq. 1.5)
By multiplying both the numerator and the denominator of this expression by [H3O+], Equation 1.6 is obtained.
Keq 3
3
[HB][A ][H O ] [HA][B ][H O ] 5
2 1
2 1
(Eq. 1.6)
Reorganization of Equation 1.6 gives the particularly useful Equation 1.7.
K K
eq 3 K
3
aHA aHB
[H O ][A ] [HA]
[HB]
[H O ][B ] 5
1 2
1 2
• =
(Eq. 1.7)
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What Equation 1.7 tells us is (1) that if HA is a stronger acid than HB (i.e., if the Ka of HA is greater than the Ka of HB), then the equilibrium for the reaction of HA with B– will lie to the right, and (2) that the equilibrium constant for the reaction will be given by the ratios of the Ka values for the two acids. The other measure of acidity, the pKa, which was defined earlier as –log10(Ka), is a much more commonly used measure of acid strength. Stronger acids have an algebraically smaller value of the pKa than weaker acids. Using pKa values for the two acids, Equation 1.7 can be rewritten as the useful Equations 1.8 and 1.9:
log eq) p aHB p
aHA
( K 5 K 2 K (Eq. 1.8)
Keq510(pKaHB2pKaHA) (Eq. 1.9)
As pointed out earlier, acid the strength of a covalent compound varies with the ele- ment to which the acidic hydrogen is bonded. As the covalent radius of that element in- creases, the acid strength increases, and as the electronegativity of the element bonded to the acidic hydrogen increases, acid strength increases. This allows us to observe that Se—H bonds are more acidic than S—H bonds; that S—H bonds are, in turn, more acidic than O—H bonds; and that O—H bonds are more acidic than N—H bonds, which are, in turn, more acidic than C—H bonds. One consequence of this is that we can expand our view of acids and, more particularly, bases. Bases such as amide anion, NH2–, the conju- gate base of ammonia (NH3, pKa ≈ 35) and methyllithium, CH3Li, the conjugate base of methane (CH4, pKa ≈ 45–60), are much stronger bases than hydroxide ion, the conjugate base of water. It is worth repeating here that although the pKa scale was originally defined for Lowry-Brứnsted acids in aqueous solution, we will use it for a wide variety of acids and bases which are much too strong to exist in aqueous solution without reacting with the water. Some typical pKa ranges are given in Table 1.1.
Let us now examine some examples of acid-base reactions in organic chemistry and see just how to approach the problems associated with proton transfer chemistry.
Example 1. CH3Li + H2O →
In this reaction, as in all potential acid-base reactions, we must first identify the acid and the base. While doing so, we must keep in mind that there is one acid and one base on each side of the arrow. In this example, the two species reacting are methyllithium, CH3Li,
Table 1.1 Typical Ranges of pKa Values for Representative Acids
Functional Group Acidic Bond Acid Examples Conjugate
Base Examples pKa
Range
Alcohol O–H RO–H (CH3)3C–OH, H2O RO– (CH3)3C–O–, OH– 15–19
Carboxylic acid COO–H RCOO–H HCO–OH, CH3CO–OH RCO–O– HCO–O–, CH3CO–O– 4–6 Amine N–H R2N–H (CH3)2CH–NH2, NH3 R2N– NH2–, (CH3)2CH–NH– 33–38
Amide CON–H RCONR–H CH3CO–NH2 RCONR– CH3CO–NH– 15–17
Alkane C–H R–H CH4, (CH2)6 R– CH3–, [(CH2)5CH]– 45–60
Alkyne C≡C–H RC≡C–H CH3–C≡C–H RC≡C– CH3–C≡C– 25–30
Aldehyde, ketone COC–H RCOCR2–H CH3COCH3 RCOCR2– CH3COCH2– 25–30
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and water, H2O. The methyllithium is a compound that contains a Group IA metal, lithium, and we are usually safe in considering that such compounds will react as ionic compounds containing Li+ cations. This means that meth- yllithium also contains the CH3– anion. This ion is capable of accepting a proton to become the neutral molecule, methane, CH3—H (or CH4); it is the conjugate base of methane. More importantly, it is highly unlikely that CH3–
will readily lose a proton to become CH22–. Therefore, the water is almost certain to be the source of protons in this reaction. At this point, we can fill out the table of reacting species in Table 1.2.
This now allows us to write the balanced equation for this reaction:
CH3– + Li + H2O → CH4 + Li+ OH–
From the data in Table 1.1, we find that the pKa of H2O is in the of 15 to 18 and that the pKa of CH4 is in the range of 45 to 60. On substituting these values in Equation 1.9, we obtain:
Keq510(45 18)2 51027 and Keq510(60 15)2 51045
as the minimum and maximum values for the equilibrium constant for this reaction.
These are both huge numbers, so we expect that methyllithium will react completely with water to give methane gas and lithium hydroxide. It does—methyllithium is rapidly (ex- plosively!) and quantitatively destroyed by water.
Example 2. LiN(CH3)2 + (CH2)5 →
In this reaction, the first step is also to identify the acid and the base. In this example, the two species reacting are lithium dimethylamide, LiN(CH3)2, and cyclopentane, (CH2)5. Again, we may consider that the lithium dimethylamide will react as an ionic lithium compound, which means that it contains the (CH3)2N– anion. This ion is capable of accept- ing a proton to become the neutral molecule, dimethylamine, (CH3)2N—H; it is the conju- gate base of dimethylamine. So the cyclopentane, which cannot accept a proton, is forced to be the source of protons in this reaction. At this point, we can fill out the table of react- ing species (Table 1.3).
This now allows us to write the balanced equation for this reaction:
N H3C H3C
Li + Li + N
H3C H3C H H
From the data in Table 1.1, we find that the pKa of (CH2)5 is in the range of 45 to 60 and the pKa of (CH3)2NH is in the range of 33 to 38. On substituting these values in Equation 1.9, we obtain:
Keq510(45 38)2 51025 and Keq510(60 33)2 510227
as the maximum and minimum values for the equilibrium constant for this reaction. These are both very small numbers, so we expect that lithium dime- thylamide will not react with cyclopentane. In fact, these numbers suggest that the reverse reaction (between cyclopentyllithium and dimethylamine to give cyclopentane and lithium dimethylamide) should be the one that actually occurs, and this is observed.
Table 1.2 Reacting Species
Reacting Species Conjugate
Acid H2O OH–
Base CH3– CH4
Table 1.3 Reacting Species
Reacting Species Conjugate Acid (CH2)5 [(CH2)4CH]– Base (CH3)2N– (CH3)2NH
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Worked Problem
1-3 Complete the following equations using the approximate pKa values given in Table 1.1. If no reaction should occur, write “N/R” and give the reason.
(a) CH3CH2CH2CH2Li + CH3OH → (b) NaNH2 + HCºCCH2CH3 → (c) LiN(CH3)2 + (CH3)4C →
§Answer below.
Leveling Effect
There is one very important consequence of Equations 1.6 and 1.7, and this is known as the leveling effect. Simply stated, the leveling effect states that one cannot use an acid or base in a solvent that has a stronger conjugate base or acid than the acid one is trying to use. In other words, the solvent limits both the strongest acid you can use (the conjugate acid of the solvent) and the strongest base (the conjugate base of the solvent). In water, for example, this means that any attempt to use an acid stronger than hydronium ion, H3O+, is doomed to failure, as is any attempt to use a base stronger than hydroxide ion, OH–, because the stronger acid will simply react with the water to give hydronium ion, and the stronger base will, likewise, react with the water to give hydroxide ion. As an illustration, because meth- ane (CH4) has a pKa of 45 to 60, one cannot use its conjugate base, methyllithium (CH3Li) in any solvent with a smaller pKa. Using the equation of Worked Problem 1-3 (a) above, we see that if one tries to use methyllithium in water, the solvent reacts with the base to give lithium hydroxide (the conjugate base of water) and methane. For similar reasons, one cannot use concentrated sulfuric acid (pKa ≈ –7) in aqueous solution, because it reacts with the water to give the conjugate acid of water, hydronium ion, H3O+. Beginning students at- tempting to use H3O+ and OH– in the same solution frequently commit this error (usually when answering a mechanism question). Of course, these two react to give water. Similar errors are frequently made using other incompatible acid-base combinations.
Problem
1-3 The reactant, product, and intermediate in the two-step reaction below are all correct, but the mechanism as written is wrong. What is the error in the mecha- nism as written? What changes are required to correct it (assume that the reac- tion occurs in D2SO4/D2O)?
D D
OD
D OD
§ Answers for Worked Problems:
(a) CH3CH2CH2CH2Li + CH3OH → CH3CH2CH2CH3 + Li+OCH3– (b) NaNH2 + HC≡CCH2CH3 → Na+ –C≡CCH2CH3 + NH3
Note that in this reaction, only the hydrogen directly attached to the triple bond is attacked. Its pKa from Table 1.1 is approximately 25 to 30, making it a stronger acid than ammonia; the approximate pKa values of the other hydrogens are predicted to be in the 45 to 60 range.
(c) LiN(CH2CH3)2 + (CH3)4C → N/R
The conjugate acid of LiN(CH2CH3)2 is H–N(CH2CH3)2, which is a stronger acid than (CH3)4C. Using Equa- tion 1.9 and the pKa values in Table 1.1, we can calculate that the equilibrium constant in favor of products will be less than 10–5, corresponding to no reaction; the equilibrium strongly favors the weaker acid and the weaker base.
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