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  Properties of Amines Homework Help - K-12 Grade Level, College Level Chemistry

Introduction to Properties of Amines

Boiling Point and Water Solubility

It is useful to compare the water solubility and boiling points of amines with those of corresponding ethers and alcohols. The dominant issue here is hydrogen bonding and first table below documents the powerful intermolecular attraction that results from -O-H---O- hydrogen bonding in alcohols (light blue columns in the table). Subsequent -N-H---N- hydrogen bonding is weaker, like the lower boiling points of likewise sized amines (light green columns in the diagram) demonstrate. Alkanes provide reference compounds in which hydrogen bonding is not possible, and the increase in boiling point for equal 1º-amines is approximately half the increase observed for equivalent alcohols.

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The second table demonstrates variations associated with isomeric 1º, 2º & 3º-amines, similarly the effect of chain branching. Because 1º-amines have two hydrogens available for hydrogen bonding, we suppose them to have higher boiling points than isomeric 2º-amines, which sequentially should boil higher than isomeric 3º-amines (no hydrogen bonding). Indeed, 3º-amines have boiling points identical to equivalent sized ethers; and in all but the smallest compounds, subsequent ethers, 3º-amines and alkanes have same boiling points. In the examples displayed here, it is further illustrated that chain branching reduces boiling points by 10 to 15 ºC.

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Basicity of Amines

A review of basic acid-base concepts should be useful to the following discussion. Such as ammonia, most amines are Lewis and Brønsted bases, but their base strength can be changed enormously by substituents. It is usual to compare basicity's quantitatively by using the pKa's of their conjugate acids rather than their pKb's. Because pKa + pKb = 14, the higher the pKa the stronger the base, in contrast to the usual inverse relationship of pKa with acidity. Various simple alkyl amines have pKa's in the range 9.5 to 11.0, and their water solutions are basic (have a pH of 11 to 12, depending on concentration). The first four compounds in the table, including ammonia, fall into that category.
The last five compounds (colored cells) are considerably weaker bases as a consequence of three factors. The first of these is the hybridization of the nitrogen. The nitrogen is sp2 hybridized, in pyridine, and an sp hybrid nitrogen is part of the triple bond, in nitriles (last entry). In each of these compounds (shaded red in the diagram) the non-bonding electron pair is localized on the nitrogen atom, but increasing s-character brings it closer to the nitrogen nucleus, reducing its tendency to bond to a proton.

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Acidity of Amines

We generally think of amines as bases, but it must be kept in mind that 1º and 2º-amines are also very weak acids (ammonia has a pKa = 34). In this way it should be noted that pKa is being employed as a measure of the acidity of the amine itself rather than its conjugate acid, as in the earlier section. This is expressed by the following hypothetical equation for ammonia:

NH3   +   H2O   ____>   NH2(-)   +   H2O-H(+)

Similar issues that decreased the basicity of amines increase their acidity. This is depicted by the following instances, which are displayed in order of increasing acidity. It should be noted that the first four illustrations have similar order and degree of increased acidity same as they exhibited decreased basicity in the previous table. The first compound is a usual 2º-amine, and the 3 next to it are characterized by changing degrees of nitrogen electron pair delocalization. The last two compounds (shaded blue in the diagram) display the effect of adjacent sulfonyl and carbonyl groups on N-H acidity. From earlier discussion it should be clear that the basicity of these nitrogens is respectively reduced.

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The acids displayed here may be converted to their conjugate bases by reaction with bases derived from weaker acids (stronger bases). Three illustrations of such type of reactions are displayed below, with the acidic hydrogen colored red in each example. For complete conversion to the conjugate base, as displayed, a reagent base approximately 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

Important Reagent Bases

The importance of all these acid-base relationships to practical organic chemistry lies in the requirement for organic bases of changing strength, as reagents tailored to the requirements of purticular reactions. The common base sodium hydroxide is not soluble in several organic solvents, and so is not extensively used as a reagent in organic reactions. Several base reagents are alkoxide salts, amines or amide salts. Because alcohols are much stronger acids than amines, the conjugated base of them are weaker than amide bases, and fill the gap in base strength between amide and amines salts. In the table, pKa again consider to the conjugate acid of the base drawn above it.

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Pyridine is generally employed as an acid scavenger in reactions that produce mineral acid co-products. Its nucleophilicity and basicity may be altered by steric hindrance, as in the example of 2,6-dimethylpyridine (pKa=6.7), or resonance stabilization, as in the case of 4-dimethylaminopyridine (pKa=9.7). Hünig's base is comparatively non-nucleophilic (because of steric hindrance), and like DBU is frequently used as the base in E2 elimination reactions conducted in non-polar solvents. Barton's base is poorly-nucleophilic, strong, neutral base which serves in cases where electrophilic substitution of DBU or another amine bases is a problem. Alkoxides are stronger bases that are frequently employed in the subsequent alcohol as solvent, or for greater reactivity in DMSO. At last, the two amide bases see extensive use in generating enolate bases from carbonyl compounds and other weak carbon acids.

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