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Goals
After completing this section you should be able to do so.
- explain the basicity and nucleophilicity of amines.
- explain why amines are more basic than amides and better nucleophiles.
- describe how an amine can be extracted from a mixture that also contains neutral compounds, and illustrate the reactions that occur with the appropriate equations.
- explain why primary and secondary (but not tertiary) amines can be considered very weak acids, and illustrate the synthetic utility of the strong bases that can be formed from these weak acids.
key terms
Make sure you can define the key term below and use it in context.
- In between
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 previous chapters that amines react with electrophiles in various 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 pkAof about 10 or 11. However, these simple amines are all more basic (that is, they have a higher pkA) than ammonia. Why? Remember that alkyl groups are electron donating compared to hydrogen, and the presence of an electron donating group stabilizes ions that carry 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 p.kAlike the last one.
Basicity of Nitrogen Groups
In this section, we consider the relative basicity of amines. In assessing the basicity of a nitrogen-containing organic functional group, we must ask the central question: how reactive (and therefore how basic and nucleophilic) is the lone pair of electrons on nitrogen? In other words, how much does this lone pair want to break away from the nitrogen nucleus and form a new bond with a hydrogen? The lone pairs of electrons make the nitrogen in the amines electron-dense, represented by a red color on the electrostatic potential map in the lower left corner. Amines are basic and readily react with hydrogen from electron-deficient acids, as seen below.
Amines are one of the few neutral functional groups considered basic, a consequence of the presence of lone pairs of electrons in nitrogen. During an acid/base reaction, lone pairs of electrons attack an acidic hydrogen to form an 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 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= -logKB
Just as the acidity of a carboxylic acid can be measured by defining an acidity constant, KA(Section 2-8) the basic strength of an amine can be measured by defining an analogous basicity constant KB. The higher the value of KBand the smaller the pK valueB, the more favorable the proton transfer equilibrium and the stronger the base.
However, Kb values are not often used to discuss the relative basicity of amines. It is common to compare the basicity of amines usingkA's of their conjugate acids, which are the corresponding ammonium ion. Fortunately, K.Ae kBValues 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 gives the equation for the autoionization of water:
\[\cancel{\ce{RNH3+}(aq)}+\ce{H2O}(l)+\cancel{\ce{RNH2}(aq)}+\ce{H2O}(l)⇌\ce{H3O+ }(aq)+\cancel{\ce{RNH2}(aq)}+\ce{OH-}(aq)+\cancel{\ce{RNH3+}(aq)}\]
\[\ce{2H2O}(l)⇌\ce{H3O+}(aq)+\ce{OH-}(aq)\]
SincekThe expression for a chemical equation formed by adding two or more other equations is the mathematical product of the input equations.kconstants.
kAXKB= {2 hours2OH3Ö+}{OH-} = kW
\[K_\ce{a}=\dfrac{K_\ce{w}}{K_\ce{b}}\]
pKA+ pKB=14
So if the KAfor an ammonium ion the K is knownBfor the corresponding amine can be calculated using the equation KB= Kc/KA. This relationship shows that as an ammonium ion becomes more acidic (KAincreases /pKAdecreases) the corresponding base becomes weaker (KBpias / pKBincreases)
weaker base= K majorAand lower pKAdes Ammonium-Ions
stronger base= Minor KAand higher pKAdes Ammonium-Ions
Like ammonia, most amines are Brønsted-Lowry and Lewis bases, but their base strength can be altered enormously by substituents. Most simple alkylamines have pKAranges from 9.5 to 11.0, and its aqueous solutions are basic (having a pH of 11 to 12, depending on concentration).
Heterocyclic aromatic amines (such as pyrimidine, pyridine, imidazole, pyrrole) are significantly weaker bases due to three factors. The first of these is nitrogen hybridization. The heterocyclic nitrogen is in each case sp2hybridized. The increasing character of s brings it closer to the nitrogen nucleus, reducing its propensity to bind a proton compared to sp3hybridized nitrogens. The very low basicity of pyrrole reflects the exceptional nitrogen pair delocalization associated with its incorporation into an aromatic ring. Imidazole (pKA= 6.95) is more than a million times more basic than pyrrole because sp2Nitrogen, which forms part of a double bond, is structurally similar to pyridine and has comparable basicity.
Basicity of common amines (pKAof conjugated ammonium ions)
Inductive effects on nitrogen basicity
Alkyl groups donate electrons to the more electronegative nitrogen. The inductive effect makes the electron density in alkylamine nitrogen greater than in ammonia nitrogen. The small amount of extra negative charge accumulated on the nitrogen atom makes the lone pair of electrons even more attractive to hydrogen ions. Consequently, primary, secondary, and tertiary alkylamines are more basic than ammonia.
| Connection | pKA the conjugate acid |
| NH3 | 9.3 |
| CH3NH2 | 10.66 |
| (CH3)2NH | 10.74 |
| (CH3)3N | 9.81 |
Comparison of basicity of alkylamines with amides
The nitrogen atom is strongly basic when in an amine, butnodistinctly basic when part of an amide group. Whereas the lone pair of electrons in an amine nitrogen is located in one site, the lone pair in an amide nitrogen is delocalized by resonance. The electron density - in the form of a lone pair of electrons - is stabilized by resonance delocalization, even though no negative charge is involved. Here's another way to think about it: the lone pair of electrons in an amide nitrogen are not available for bonding with a proton - these two electrons are very stable as they are 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 oxygen. This greatly reduces the basicity of the nitrogen pairs in an amide.
Comparison of amines and amides to rationalize pKAValues of their conjugate acids
Amine extraction in the laboratory
Extraction is often used in organic chemistry to purify compounds. Liquid-liquid extractions take advantage of the differential solubility of a substance in two immiscible liquids (eg, 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 extraction solvent. The two immiscible liquids can then be easily separated with a separatory funnel. The basicity of amines can be exploited using the protonated salt (RNH2+Kl−), which is soluble in water. The salt is extracted into the aqueous phase, leaving neutral compounds in the non-aqueous phase. The aqueous layer is then treated with base (NaOH) to regenerate the amine and NaCl. A second extraction separation is then performed to isolate the amine in the non-aqueous layer and leave NaCl in the aqueous layer.
Important reagent bases
The importance of all these acid-base relationships in practical organic chemistry lies in the need for organic bases of varying strengths as reagents tailored to the needs of particular 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 of the basic reagents are alkoxide salts, amines or amide salts. Because alcohols are much stronger acids than amines, their conjugate bases are weaker than amine bases, bridging the gap in base strength between amines and amide salts.
| base name | pyridine | Trietil Amen | Base Hünigs | Bartons Base | Potassium t-butóxido | Natrium-HMDS | LDA |
|---|---|---|---|---|---|---|---|
| Formula |  | (C2H5)3N |  |  | (CH3)3CO(–)k(+) | [(CH3)3E]2N(–)Already(+) | [(CH3)2CH]2N(–)li(+) |
| pKA conjugate acid | 5.3 | 10.7 | 11.4 | 14 | 19 | 26 | 35,7 |
Basicity of common amines (pKAof conjugated ammonium ions)
Pyridine is commonly used as an acid scavenger in reactions that produce mineral acid by-products. Its basicity and nucleophilicity can 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 commonly used as a base in E2 elimination reactions performed in nonpolar solvents. Barton's base is a strong, weakly nucleophilic, neutral base that serves in cases where electrophilic substitution of other amine bases is a problem. Alkoxides are stronger bases that are often used in the corresponding alcohol as a solvent or in DMSO for greater reactivity. Finally, the two amide bases are commonly used to generate enolate bases from carbonyl compounds and other weak carboxylic acids.
In addition to the basic function, 1Öe 2ÖAmines can act as very weak acids. Their N-H proton can be removed if they are treated with a strong enough base. An example is the formation of lithium diisopropylamide (LDA, LiN[CH(CH3)2]2) per reaction N-Butyllithium with diisopropylamine (pKA 36)(Section 22-5). LDA is a very strong base and is commonly used to generate enolate ions by deprotonating an alpha hydrogen from carbonyl compounds.(Section 22-7).
exercises
Credits and Attributions
dr Dietmar KennepohlFCIC (Professor of Chemistry,Athabasca University)
Prof. Steven Farmer (sonoma state university)
William Reusch, Professor Emeritus (michigan state u.),Organic chemistry virtual book
(Video) Ranking by Basicity of AminesOrganic Chemistry with a Biological FocusvonTim Soderberg(University of Minnesota, Morris)