1

Chapter 2

Acids & Bases

2

Brnsted-Lowry Acid : Any species that can donates a

proton.

Base : Any species that can accept a proton.

Most important principle: Conjugate acid-base pairs.

CH3 C

O

OH + CH3 NH2 CH3 C

O

O-

+ CH3 NH3+

acid conjugate base

base conjugate

acid

3

Lewis Acids and Bases

Acids accept electron pairs => electrophile

Bases donate electron pairs => nucleophile

4

Any reaction in which one species donates an electron pair to another species is a Lewis acid-base reaction.

In a Lewis acid-base reaction, a Lewis base donates an electron pair to a Lewis acid.

Lewis acid-base reactions illustrate a general pattern in organic chemistry. Electron-rich species react with

electron-poor species.

In the simplest Lewis acid-base reaction one bond is formed and no bonds are broken. This is illustrated in

the reaction of BF3 with H2O. H2O donates an electron

pair to BF3 to form a new bond.

5

Acid and Base Strength

Acid dissociation constant, Ka

Base dissociation constant, Kb

For conjugate pairs, (Ka)(Kb) = Kw

Spontaneous acid-base reactions proceed

from stronger to weaker.

CH3 C

O

OH + CH3 NH2 CH3 C

O

O-

+ CH3 NH3+

pKa 4.74 pKb 3.36 pKb 9.26 pKa 10.64

6

Structural Effects on Acidity

Electronegativity

Size

Resonance stabilization of conjugate base

Hybridization

Inductive effect

7

Electronegativity

As the bond to H becomes more polarized, H

becomes more positive and the bond is easier

to break.

8

Size

As size increases, the H is more loosely held and the bond is easier to break.

A larger size also stabilizes the anion.

9

Resonance Stabilization

Delocalization of the negative charge on the conjugate base will stabilize the anion, so the substance is a stronger acid.

More resonance structures usually mean greater stabilization.

CH3CH2OH < CH3C

O

OH < CH3 S

O

O

OH

Acidity increases

10

Inductive Effects

When electron density is pulled away from the negative charge through bonds by very electronegative atoms, it is

referred to as an electron withdrawing inductive effect.

More electronegative atoms stabilize regions of high electron density by an electron withdrawing inductive effect.

The more electronegative the atom and the closer it is to the site of the negative charge, the greater the effect.

The acidity of HA increases with the presence of electron withdrawing groups in A.

11

Hybridization Effects

The final factor affecting the acidity of HA is the hybridization of A.

The higher the percent of s-character of the hybrid orbital, the closer the lone pair is held to the nucleus,

and the more stable the conjugate base.

Let us consider the relative acidities of three different

compounds containing CH bonds.

12

It should be noted that:

Strong bases have weak conjugate acids with high pKa values, usually > 12.

Strong bases have a net negative charge, but not all negatively charged species are strong bases. For example,

none of the halides F, Cl, Br, or I, is a strong base.

Carbanions, negatively charged carbon atoms, are especially strong bases. A common example is butyllithium.

Two other weaker organic bases are triethylamine and pyridine.

Commonly Used Bases in Organic Chemistry

(a) The following compound can become protonated on any

of the nitrogen atoms. However, one of these nitrogen

atoms is much more basic than the others.

(i) Draw the important resonance forms of the products of

protonation on EACH of the three nitrogen atoms (include

the formal charge when necessary)

(ii) Determine which nitrogen atom is the most basic. Give

reason for your choice.

13

CH3 NH CNH

NH2

....

:

14

Rank the following protons in order of increasing acidity. Explain

your answer.

H

H

H

H

O

1

2

3