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21.4: Acidity and Basicity of Amines

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  • Objectives

    After completing this section, you should be able to

    1. account for the basicity and nucleophilicity of amines.
    2. explain why amines are more basic than amides, and better nucleophiles.
    3. describe how an amine can be extracted from a mixture that also contains neutral compounds illustrating the reactions which take place with appropriate equations.
    4. explain why primary and secondary (but not tertiary) amines may be regarded as very weak acids, and illustrate the synthetic usefulness of the strong bases that can be formed from these weak acids.

    Key Terms

    Make certain that you can define, and use in context, the key term below.

    • amide

    Study Notes

    The lone pair of electrons on the nitrogen atom of amines makes these compounds not only basic, but also good nucleophiles. Indeed, we have seen in past chapters that amines react with electrophiles in several polar reactions (see for example the nucleophilic addition of amines in the formation of imines and enamines in Section 19.8).

    The ammonium ions of most simple aliphatic amines have a pKa of about 10 or 11. However, these simple amines are all more basic (i.e., have a higher pKa) than ammonia. Why? Remember that, relative to hydrogen, alkyl groups are electron releasing, and that the presence of an electron‑releasing group stabilizes ions carrying a positive charge. Thus, the free energy difference between an alkylamine and an alkylammonium ion is less than the free energy difference between ammonia and an ammonium ion; consequently, an alkylamine is more easily protonated than ammonia, and therefore the former has a higher pKa than the latter.

    Basicity of nitrogen groups

    In this section we consider the relative basicity of  amines. When evaluating the basicity of a nitrogen-containing organic functional group, the central question we need to ask ourselves is: how reactive (and thus how basic and nucleophilic) is the lone pair on the nitrogen? In other words, how much does that lone pair want to break away from the nitrogen nucleus and form a new bond with a hydrogen. The lone pair electrons makes the nitrogen in amines electron dense, which is represents by a red color in the electrostatic potential map present below left. Amine are basic and easily react with the hydrogen of acids which are electron poor as seen below.  


    Amines are one of the only neutral functional groups which are considered basis which is a consequence of the presence of the lone pair electrons on the nitrogen. During an acid/base reaction the lone pair electrons attack an acidic hydrogen to form a N-H bond.  This gives the nitrogen in the resulting ammonium salt four single bonds and a positive charge. 

    Amines react with water to establish an equilibrium where a proton is transferred to the amine to produce an ammonium salt and the hydroxide ion, as shown in the following general equation:

    \[RNH2_{(aq)}+H_2O_{(l)} \rightleftharpoons RNH3^+_{(aq)}+OH^−_{(aq)} \label{16.5.4}\]

    The equilibrium constant for this reaction is the base ionization constant (Kb), also called the base dissociation constant:

    \[K_b=\dfrac{[RNH3^+][OH^−]}{[NH2]} \label{16.5.5}\]

    pKb = -log Kb

    Just as the acid strength of a carboxylic acid can be measured by defining an acidity constant Ka (Section 2-8), the base strength of an amine can be mea­sured by defining an analogous basicity constant Kb. The larger the value of Kb and the smaller the value of pKb, the more favorable the proton-transfer equi­librium and the stronger the base.

    However, Kb values are often not used to discuss relative basicity of amines. It is common to compare basicity's of amines by using the Ka's of their conjugate acids, which is the corresponding ammonium ion. Fortunately, the Ka and Kb values for amines are directly related. 

    Consider the reactions for a conjugate acid-base pair, RNH3+ − RNH2:

    \[\ce{RNH3+}(aq)+\ce{H2O}(l)⇌\ce{RNH2}(aq)+\ce{H3O+}(aq) \hspace{20px} K_\ce{a}=\ce{\dfrac{[RNH2][H3O]}{[RNH3+]}}\]

    \[\ce{RNH2}(aq)+\ce{H2O}(l)⇌\ce{RNH3+}(aq)+\ce{OH-}(aq) \hspace{20px} K_\ce{b}=\ce{\dfrac{[RNH3+][OH-]}{[RNH2]}}\]

    Adding these two chemical equations together yields the equation for the autoionization for water:



    Given that the K expression for a chemical equation formed from adding two or more other equations is the mathematical product of the input equations’ K constants.

    Ka X Kb = {2 H2O} / (H3O+}{OH-} = Kw


    pKa + pKb =14

    Thus if the Ka for an ammonium ion is know the Kb for the corresponding amine can be calculated using the equation Kb = Kw / Ka.  This relationship shows that as an ammonium ion becomes more acidic (Ka increases / pKa decreases) the correspond base becomes weaker (Kb decreases / pKb increases)

    Weaker Base = Larger Ka and Smaller pKa of the Ammonium ion

    Stronger Base = Smaller Ka and Larger pKa of the Ammonium ion

    Like ammonia, most amines are Brønsted-Lowry and Lewis bases, but their base strength can be changed enormously by substituents. Most simple alkyl amines have pKa's in the range 9.5 to 11.0, and their aqueous solutions are basic (have a pH of 11 to 12, depending on concentration).

    Aromatic herterocyclic amines (such as pyrimidine, pyridine, imidazole, pyrrole) are significantly weaker bases as a consequence of three factors. The first of these is the hybridization of the nitrogen. In each case the heterocyclic nitrogen is sp2 hybridized. The increasing s-character brings it closer to the nitrogen nucleus, reducing its tendency to bond to a proton compared to sp3 hybridized nitrogens. The very low basicity of pyrrole reflects the exceptional delocalization of the nitrogen electron pair associated with its incorporation in an aromatic ring. Imidazole (pKa = 6.95) is over a million times more basic than pyrrole because the sp2 nitrogen that is part of one double bond is structurally similar to pyridine, and has a comparable basicity.

    Basicity of common amines (pKa of the conjugate ammonium ions)

    Inductive Effects in Nitrogen Basicity

    Alkyl groups donate electrons to the more electronegative nitrogen. The inductive effect makes the electron density on the alkylamine's nitrogen greater than the nitrogen of ammonia. The small amount of extra negative charge built up on the nitrogen atom makes the lone pair even more attractive towards hydrogen ions. Correspondingly, primary, secondary, and tertiary alkyl amines are more basic than ammonia.

    Compound pKa
    NH3 9.3
    CH3NH2 10.66
    (CH3)2NH 10.74
    (CH3)3N 9.81

    Comparing the Basicity of Alkylamines to Amides

    The nitrogen atom is strongly basic when it is in an amine, but not significantly basic when it is part of an amide group. While the electron lone pair of an amine nitrogen is localized in one place, the lone pair on an amide nitrogen is delocalized by resonance. The electron density – in the form of a lone pair – is stabilized by resonance delocalization, even though there is not a negative charge involved. Here’s another way to think about it: the lone pair on an amide nitrogen is not as available for bonding with a proton – these two electrons are too stable being part of the delocalized pi-bonding system. The electrostatic potential map shows the effect of resonance on the basicity of an amide. The map shows that the electron density, shown in red, is almost completely shifted towards the oxygen. This greatly decreases the basicity of the lone pair electrons on the nitrogen in an amide.


    Amine Extraction in the Laboratory

    Extraction is often employed in organic chemistry to purify compounds. Liquid-liquid extractions take advantage of the difference in solubility of a substance in two immiscible liquids (e.g. ether and water). The two immiscible liquids used in an extraction process are (1) the solvent in which the solids are dissolved, and (2) the extracting solvent. The two immiscible liquids are then easily separated using a separatory funnel. For amines one can take advantage of their basicity by forming the protonated salt (RNH2+Cl), which is soluble in water. The salt will extract into the aqueous phase leaving behind neutral compounds in the non-aqueous phase. The aqueous layer is then treated with a base (NaOH) to regenerate the amine and NaCl. A second extraction-separation is then done to isolate the amine in the non-aqueous layer and leave behind NaCl in the aqueous layer.


    Important Reagent Bases

    The significance of all these acid-base relationships to practical organic chemistry lies in the need for organic bases of varying strength, as reagents tailored to the requirements of specific reactions. The common base sodium hydroxide is not soluble in many organic solvents, and is therefore not widely used as a reagent in organic reactions. Most base reagents are alkoxide salts, amines or amide salts. Since alcohols are much stronger acids than amines, their conjugate bases are weaker than amine bases, and fill the gap in base strength between amines and amide salts. 

    Base Name Pyridine Triethyl
    Hünig's Base Barton's
    Sodium HMDS LDA
    Formula (C2H5)3N (CH3)3CO(–) K(+) [(CH3)3Si]2N(–) Na(+) [(CH3)2CH]2N(–) Li(+)
    pKa of conjugate acid 5.3 10.7 11.4 14 19 26 35.7

    Basicity of common amines (pKa of the conjugate ammonium ions)

    Pyridine is commonly used as an acid scavenger in reactions that produce mineral acid co-products. Its basicity and nucleophilicity may be modified by steric hindrance, as in the case of 2,6-dimethylpyridine (pKa=6.7), or resonance stabilization, as in the case of 4-dimethylaminopyridine (pKa=9.7). Hünig's base is relatively non-nucleophilic (due to steric hindrance), and is often used as the base in E2 elimination reactions conducted in non-polar solvents. Barton's base is a strong, poorly-nucleophilic, neutral base that serves in cases where electrophilic substitution of other amine bases is a problem. The alkoxides are stronger bases that are often used in the corresponding alcohol as solvent, or for greater reactivity in DMSO. Finally, the two amide bases see widespread use in generating enolate bases from carbonyl compounds and other weak carbon acids.

    In addition to acting as a base, 1o and 2o amines can act as very weak acids. Their N-H proton can be removed if they are reacted with a strong enough base. An example is the formation of lithium diisopropylamide (LDA, LiN[CH(CH3)2]2) by reacting n-butyllithium with diisopropylamine (pKa 36) (Section 22-5). LDA is a very strong base and is commonly used to create enolate ions by deprotonating an alpha-hydrogen from carbonyl compounds (Section 22-7).







    Select the more basic amine from each of the following pairs of compounds.





    The 4-methylbenzylammonium ion has a pKa of 9.51, and the butylammonium ion has a pKa of 10.59. Which is more basic? What's the pKb for each compound?







    The butylammonium is more basic. The pKb for butylammonium is 3.41, the pKb for 4-methylbenzylammonium is 4.49.

    Contributors and Attributions


    After completing this section, you should be able to

    1. use the concept of resonance to explain why arylamines are less basic than their aliphatic counterparts.
    2. arrange a given series of arylamines in order of increasing or decreasing basicity.
    3. discuss, in terms of inductive and resonance effects, why a given arylamine is more or less basic than aniline.

    Study Notes

    With reference to the discussion of base strength, the traditional explanation for the base‑strengthening effect of electron‑releasing (I) substituents is that such substituents help to stabilize the positive charge on an arylammonium ion more than they stabilize the unprotonated compound, thereby lowering ΔG°.

    The electron‑withdrawing (i.e., deactivating) substituents decrease the stability of a positively charged arylammonium ion.

    Note that the arylammonium ion derived from aniline, PhNH3+, is commonly referred to as the anilinium ion.

    Basicity of Aniline

    Aniline is substantially less basic than methylamine, as is evident by looking at the pKa values for their respective ammonium conjugate acids (remember that the lower the pKa of the conjugate acid, the weaker the base).


    This difference is basicity can be explained by the observation that, in aniline, the lone pair of electrons on the nitrogen are delocalized by the aromatic p system, making it less available for bonding to H+ and thus less basic. The lone pair electrons of aniline are involved in four resonance forms making them more stable and therefore less reactive relative to alkylamines. 

    The effect of delocalization can be seen when viewing the electrostatic potential maps of aniline an methyl amine. The nitrogen of methyl amine has a significant amount of electron density on its nitrogen, shown as a red color, which accounts for it basicity compared to aniline. While the electron density of aniline's nitrogen is delocalized in the aromatic ring making it less basic. 


    Basicity of Substituted Arylamines

    The addition of substituents onto the aromatic ring can can make arylamines more or less basic. Substituents which are electron-withdrawing (-Cl, -CF3, -CN, -NO2) decrease the electron density in the aromatic ring and on the amine making the arylamine less basic. In particular, the nitro group of para-nitroaniline allows for an additional resonance form to be drawn, which further stabilizes the lone pair electrons from the nitrogen, making the substituted arylamine less basic than aniline. This effect is analogous to the one discussed for the acidity of substituted phenols in Section 17.2.


    Substituents which are electron-donating (-CH3, -OCH3, -NH2) increase the electron density in the aromatic ring and on the amine making the arylamine more basic. In the case of para-methoxyaniline, the lone pair on the methoxy group donates electron density to the aromatic system, and a resonance contributor can be drawn in which a negative charge is placed on the carbon adjacent to the nitrogen, which makes the substituted arylamine more basic than aniline.


    Increased Basicity of para-Methoxyaniline due to Electron-Donation

    The shifting electron density of aniline, p-nitroaniline, and p-methoxyaniline are seen in their relative electrostatic potential maps.  For p-Nitroaniline virtually all of the electron density, shown as a red/yellow color. is pulled toward the electron-withdrawing nitro group.  In p-methoxyaninline the electron donating methoxy group donates electron density into the ring. The amine in p-methoxyaniline is shown to have more electron density, shown as a yellow color, when compared to the amine in aniline.  


    Exercise s\(\PageIndex{1}\)

    1) Using the knowledge of the electron donating or withdrawing effects of subsituents gained in Section 16.6, rank the following compound in order of decreasing basicity.
    a) p-Nitroaniline, methyl p-aminobenzoate, p-chloroaniline
    b) p-Bromoaniline, p-Aminobenzonitrile, p-ethylaniline
    c) p-(Trifluoromethyl)aniline, p-methoxyaniline, p-methylaniline


    a) p-Chloroaniline, methyl p-aminobenzoate, p-nitroaniline
    b) p-Ethylaniline, p-Bromoaniline, p-aminobenzonitrile
    c) p-Methoxyaniline, p-methylaniline, p-(trifluoromethyl)aniline


    Contributors and Attributions

    Acidity of AminesEdit section

    We normally think of amines as bases, but it must be remembered that 1º and 2º-amines are also very weak acids (ammonia has a pKa = 34). In this respect it should be noted that pKa is being used as a measure of the acidity of the amine itself rather than its conjugate acid, as in the previous section. For ammonia this is expressed by the following hypothetical equation:

    NH3 + H2O ____> NH2(–) + H2O-H(+)

    The same factors that decreased the basicity of amines increase their acidity. This is illustrated by the following examples, which are shown in order of increasing acidity. It should be noted that the first four examples have the same order and degree of increased acidity as they exhibited decreased basicity in the previous table. The first compound is a typical 2º-amine, and the three next to it are characterized by varying degrees of nitrogen electron pair delocalization. The last two compounds (shaded blue) show the influence of adjacent sulfonyl and carbonyl groups on N-H acidity. From previous discussion it should be clear that the basicity of these nitrogens is correspondingly reduced.

    Compound C6H5SO2NH2
    pKa 33 27 19 15 10 9.6

    The acids shown here may be converted to their conjugate bases by reaction with bases derived from weaker acids (stronger bases). Three examples of such reactions are shown below, with the acidic hydrogen colored red in each case. For complete conversion to the conjugate base, as shown, a reagent base roughly a million times stronger is required.

    C6H5SO2NH2 + KOH C6H5SO2NH(–) K(+) + H2O a sulfonamide base
    (CH3)3COH + NaH (CH3)3CO(–) Na(+) + H2 an alkoxide base
    (C2H5)2NH + C4H9Li (C2H5)2N(–) Li(+) + C4H10 an amide base
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