Basicity of amines

Basicity of amines

In amines, there is a lone pair of electrons on nitrogen. The tendency of nitrogen to share these electrons with acids is responsible for the basic character of amines.

For comparison of the basic character of amines, the equilibrium constant of following reaction is a measure of their basic character.

Since [H₂O] is constant

The basically constant K_b = K_eq [H₂O]

Aliphatic amines with K_b between 10^-3 and 10^-4 are stronger bases then ammonia (K_b = 1.8 x 10^-5).

Like ammonia, amines are strong bases and react with mineral acids to form ammonium salts from which they can be liberated by treatment with a strong base like NaOH.

That alkylamines are stronger bases then ammonia can be expressed in terms of electron releasing inductive effect of alkyl groups. Alkyl groups by their electron releasing effect, concentrate electron density on nitrogen and hence make the lone pair of nitrogen more easily available for sharing with acids. Also the electron releasing effect of alkyl groups stabilizes the alkyl ammonium ion formed and hence shifts the equilibrium in forward direction making the alkylamines stronger bases than ammonia.

Thus the basic character of aliphatic amines should increase with increase of alkyl substitution. In the gasphase, the basic trend in nature is as expected i.e.,

tertiary amine > secondary amine > primary amine > ammonia.

However, the order of basicity in aqueous solution does not follow the expected trend and gets altered as revealed by their K_b values.

The basic strength in aqueous solution depends not only upon electron releasing effect but also upon steric effect and hydration effect.

a) Steric effect

Steric effect refers to the crowding of alkyl groups around N atom which hinders the attack of proton on the amine molecule and this decreases its basic strength. Since crowding of alkyl groups around N atom increases from 1^o to 3^o amines, the basic strength of amine should decrease in the order 1^o > 2^o> 3^o.

b) Hydration effect

Hydration effect refers to the stabilization of the protonated amine by water molecules. The water molecules from H - bonds with the protonated amine and release energy called hydration energy.

Greater the extent of H - bonding in protonated amine more will be its stabilization and consequently greater will be the basic strength of the corresponding amine.

Hydration due to hydrogen-bonding is maximum in monoalkyl ammonium ion (protonated cation of 1^oamine), it is less in dialkyl ammonia ion and still less in trialkyl ammonium ion. Thus basic strength should decrease from 1^o > 2^o > 3^o.

In 3^o amine, hydration is least, steric hinderance is maximum so least basic inspite of maximum inductive effect.

In 1^o amine, steric hinderance is least, hydration is maximum and inductive effect is minimum. So its basic strength is more than 3^o amine.

The resultant of all factors cause 2^o amine to be still more basic than 1^oamine.

So overall basic strength varies as 2^o > 1^o > 3^o.

Aromatic amines are weaker bases (Kb ~ 10^-9) than ammonia and aliphatic amines.

The less basic character of aniline can be explained on the following basis:

(a) Delocalization of one pair of electrons on the nitrogen atom

Amiline is a hybrid of the following resonating structures.

Thus the lone pair of electrons on nitrogen is withdrawn from it and is being partially shared with the benzene ring. Thus, in aniline the electron donating capacity of nitrogen for protonation is considerably decreased as compared to that of ammonia and aliphatic amines. Hence aniline is a weaker base than aliphatic amines and ammonia.

(b) Lower stability of the anilinium ion

Anilinium ion forward by aniline by accepting a proton is not stabilized by resonance.

Anilinium ion

does not form resonating structures similar to I, II, III and IV (in the case of aniline) 

due to absence of lone pair on N atom. Hence anilinium ion is less stable as compared to aniline. Therefore, aniline has less tendency to accept proton to form anilinium ion. This accounts for the lower basic strength of aniline.

An electron releasing group present in an aromatic amine ring especially in ortho / para position will stabilize the ammonium cation formed after the protonation of amine and hence increases the basic strength of aromatic amine.

The electron releasing groups (like -OCH₃, -CH₃, -NH₂ etc) enhance the availability of unshared electrons on nitrogen and increases the basic strength.

Electrons withdrawing groups like (-NO₂, -CN, -X etc) affect the stability of an aromatic ammonium cation and decreases the basic strength of parent aromatic amine.

The electron withdrawing groups decrease the availability of unshared electrons on nitrogen and thereby decrease the basic strength of aromatic amines.

Basicity of Amines

• A review of basic acid base concepts should be helpful to the following discussion. 
• Like ammonia, most amines are Brønsted and Lewis bases, but their base strength can be changed enormously by substituents.
•  It is common to compare basicity's quantitatively by using the pkas of their conjugate bases rather than their pK_b's. Since pK_a + pK_b = 14, the higher the pK_a the stronger the base, in contrast to the usual inverse relationship of pK_a with acidity. 
• Most simple alkyl amines have pK_a'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 following table, including ammonia, fall into that category.
• The last five compounds (colored cells) are significantly weaker bases as a consequence of three factors.
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• The first of these is the hybridization of the nitrogen. 
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• In pyridine the nitrogen is sp² hybridized, and in nitriles (last entry) an sp hybrid nitrogen is part of the triple bond. 
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• In each of these compounds (shaded red) 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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• 
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• to bond to a proton.
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• Secondly, aniline and p-nitroaniline (first two green shaded structures) are weaker bases due to delocalization of the nitrogen non-bonding elect
• Secondly, aniline and p-nitroaniline (first two green shaded structures) are weaker bases due to delocalization of the nitrogen non-bonding electron pair into the aromatic ring (and the nitro substituent). 
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•  The  electron pair delocalization is accompanied by a degree of rehybridization of the amino nitrogen atom, but the electron pair delocalization is probably the major factor in the reduced basicity of these compounds.
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•  A similar electron pair delocalization is responsible for the very low basicity (and nucleophilic reactivity) of amide nitrogen atoms (last green shaded structure). 
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• 
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•  Although 4-dimethylaminopyridine (DMAP) might appear to be a base similar in strength to pyridine or N,N-dimethylaniline, it is actually more than ten thousand times stronger, thanks to charge delocalization in its conjugate acid.
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•  The structure in the gray box shows the locations over which positive charge (colored red) is delocalized in the conjugate acid. This compound is often used as a catalyst for acyl transfer reactions.
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• Finally, the very low basicity of pyrrole (shaded blue) reflects the exceptional delocalization of the nitrogen electron pair associated with its incorporation in an aromayic ring . 
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• Indole (pK_a = -2) and imidazole (pK_a = 7.0), see above , also have similar heterocyclic aromatic rings.
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•  Imidazole 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.
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• Although resonance delocalization generally reduces the basicity of amines, a dramatic example of the reverse effect is found in the compound guanidine (pK_a = 13.6).
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•  Here, as shown below, resonance stabilization of the base is small, due to charge separation, while the conjugate acid is stabilized strongly by charge delocalization.
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• guanidine is exceptionally 
• Guanidine is protonated in physiological conditions. This conjugate acid is called the guanidinium cation, [CH₆N₃]⁺. The guanidinium cation has a charge of +1. It is a highly stable cation in aqueous solution due to the efficient resonance stabilization of the charge and efficient solvation by water molecules. As a result, itspK_a is 13.6^[7] meaning that guanidine is a very strong base in water.
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•  Consequently, aqueous solutions of guanidine are nearly as basic as are solutions of sodium hydroxide.

1. 

Strong bases have weak conjugate acids, and weak bases have strong conjugate acids.

Basicity of Amines, Acidity of Ammonium Ions

                                                                                                                                               
                      Basicity of Amines, Acidity of Ammonium Ions                                                                                                                                                                                              

• N lone pair relatively easily protonated

• note that K_b x K_a = [H⁺] [OH^-] = Kw = 10^-14
  or pK_a + pK_b = 14
• recall that when pH = pK_a , there are equal concentrations of the conjugate acid and conjugate base present (i.e., RNH₂ and RNH₃⁺ )
• for typical aliphatic amines, pK_b = 3 - 4
  so pK_a = 10 - 11 for their ammonium ions
  so at around pH 10 - 11 , RNH₂ and RNH₃⁺ are both present
• for typical aromatic amines, pK_b = 9 - 10
  so pK_a = 4 - 5 for their ammonium ions
  so at around pH 4 - 5 , ArNH₂ and ArNH₃⁺ are both present
• water solubility of amines can be easily changed with pH
  aromatic amines are water-soluble (protonated) below pH 4
  aliphatic amines are water-soluble (protonated) below pH 9

Basicity Trends

Amines as Bases

• deciding the stronger of two acids, e.g. H-A₁ and H-A₂

• the stronger acid has the more stable anion

• the stronger acid corresponds to the more favorable "reaction", left to right

• deciding the stronger of two bases, e.g. B₁ and B₂

• there are no anions here, but the stronger base STILL corresponds to the more favorable reaction

Example

• measure basicity in terms of acidity (pKa) of conjugate acid  (we will not use the pKb scale)

• compare (and don't confuse!!)

Trends in Basicity ??

• no clear trends among the aliphatic amines.....

But

• (minor) resonance delocalization (i.e. partial bonding) stabilizes (lowers energy) of the non-bonding electrons on N, these electrons are less reactive

• aromatic amines are thus less basic than aliphatic amines

• clear decrease in basicity with decreasing energy of non-bonding electrons as the hybridization of the orbitals gains more s character and less p character

• the protonated nitrile is about the strongest acid that we discuss this entire course

why aromatic amines are less basic ?

• aromatic amines are less basic due to resonance delocalization of the N lone pair

important conclusion : conjugation reduces basic nature  .

• amides are nonbasic due to strong delocalization of the N lone pair

• electron withdrawing effects decrease basicity
  because the N lone pair is less available for bonding to a proton

ACIDS & BASES -------------------an attempt

 as the source of  -I effect on

ACIDITY OF ACIDS

               Acidity of Aromatic Acids

             Acidity of carboxylic acids depends on the stability of acid anion  a conjugate base thus.

            Electron withdrawing group stabilises anion, hence increases acidity

                                                  

            Similarly , electron donating group destabilises anion, hence decreases acidity.

                                                      

            The First member of aromatic acid is a benzoic acid which dissociates as follows

                                                

            Thus  acidity  of  benzoic acid  will  depends on the stability of benzoate anion .  .Greater the stability of carboxylate ion ,it is a weaker conjugate base and tendency to react with H+ decreases  resulting in the increase of acidic nature .

            Acidity of substituted Acids: Ortho substituted benzoic acid is always a strong acid than m- and    p-derivative due to the ortho effect.

            Case I :           When group is –M and –I group.

               at  ORTHO  POSITION           

            Thus , anion is stabilised by  – M and  – I  effect and  – I  power is maximum

             AT META POSITION                  

            Anion is stabilised only by –I effect of NO2 group .

                        AT PARA POSITION           

Anion is stabilised by –M and –I effect of NO2 group

                                    NO SUBSTITUENT    

                                   

            Thus decreasing order of the stability of these anions is follows:

                                    I > III > II > IV

            Ortho derivative  is  the  most acidic  due to ortho effect therefore decreasing  order of acidity of these acids is as follows :

  

Case II :          When group has +I effect.

                                                  Ortho derivative will be most acidic due to ortho effect. The anion is destabilised due to +I effect.As distance between + I group and carboxylate increases +I influence decreases .Thus methyl at para is less donating compared to methyl at meta .At ortho position eventhough the distance between CH3 group and carboxylate is very less DUE TO ORTHO EFFECT IT IS MOST ACIDIC.

              

Case III :   When group has +M and –I effect and - I dominates +M

                                              

 Case IV :   When group has +M > –I effect 

   
 ultimately

 

PSK CHAKRAVARTHY

                                  

                                   

Why is phenol acidic?

Compounds like alcohols and phenol which contain an -OH group attached to a hydrocarbon are very weak acids.

 Alcohols are so weakly acidic that, for normal lab purposes, their acidity can be virtually ignored.

 However, phenol is sufficiently acidic for it to have recognizably acidic properties - even if it is still a very weak acid.

 A hydrogen ion can break away from the -OH group and transfer to a base.
For example, in solution in water:

Phenol is a very weak acid and the position of equilibrium lies well to the left. Phenol can lose a hydrogen ion because the phenoxide ion formed is stabilised to some extent. 

The negative charge on the oxygen atom is delocalised around the ring. The more stable the ion is, the more likely it is to form. 

One of the lone pairs on the oxygen atom overlaps with the delocalised electrons on the benzene ring.

This overlap leads to a delocalization which extends from the ring out over the oxygen atom. 

As a result, the negative charge is no longer entirely localized on the oxygen, but is spread out around the whole ion.

Spreading the charge around makes the ion more stable than it would be if all the charge remained on the oxygen. 

However, oxygen is the most electronegative element in the ion and the delocalized electrons will be drawn towards it. 

That means that there will still be a lot of charge around the oxygen which will tend to attract the hydrogen ion back again. 

That is why phenol is only a very weak acid.
Why is phenol a much stronger acid than cyclohexanol? To answer this question we must evaluate the manner in which an oxygen substituent interacts with the benzene ring. 

 It was proposed that resonance delocalization of an oxygen non-bonded electron pair into the pi-electron system of the aromatic ring was responsible for this substituent effect. A similar set of resonance structures for the phenolate anion conjugate base appears below the phenol structures.