ACID AND BASE CHEMISTRY
I. Definitions and Comparisons
Arrhenius
Acids can deliver H+
Bases can deliver HO-
H:A
MOH
H+ +
A:-
M+ + -OH
Lowry-Bronsted
Acids are H+ donors
Bases are H+ acceptors
H:A + :B-
H:B +
:A-
Lewis
Acids are e- pair acceptors
Bases are e- pair donors
A+ + :B-
A:B
Are there any contradictions in the definitions above? NO.
Without a doubt, the Lewis definition is the most comprehensive and most broadly
applicable.
But the chemistry is the same for all. The L-B definiton and the Arrhenius definition only
allow H+ (or a molecule that holds it) to be an acid but, like the Lewis acid, it is
ELECTRON DEFICIENT.
The Arrhenius definition of a base only allows for HO- to be a base, but it is electron
rich. The L-B definition of a base allows for a species to accept a proton. To do this
means the L-B base must be electron rich. From the Lewis perspective, a base
is electron rich and behaves as a base when it donates an e- pair to an electron
poor species.
The Lewis acid and base fit naturally into the notion of covalent bond formation bewteen
electron poor and electron rich groups.
II. Naked Protons?
When a H+ is the acid, generated by ionization of a compound in water, it is
more accurately represented as the HYDRONIUM ION, H3O+.
We will use the two interchangably when water is the reaction medium.
H:A
+
H3O+ + A:-
H2 O
H+ ! H3O+
III. Electrophiles and nucleophiles
The suffix -phile means "attraction to, love"
An ELECTROPHILE loves electrons. ACIDS are electrophiles. This satisfies
Lewis' definition of an acid.
A NUCLEOPHILE loves "nuclei", that is to say, a + charge. BASES are
nucleophiles. This satisfies Lewis' definition of a base.
Electrophile/Acid
E or E+
Nucleophile/Base
N: or N:-
So why the different names? The words electrophile and nucleophile are more
general and refer to species that will react with electron rich (electrophile) or
electron poor (nucleophile) compounds. Acids and bases react with acids and
bases. It's just a communication distinction.
However, in a reaction, conditions can be chosen to favor, for example, whether
a nucleophile acts like a base or a nucleophile. For example...
O
O
H
H
H
H
N
H
H
O
N:Well, N:-, what will you do? pull a H+ or make a
bond with the carbonyl C?
H
H
+
H:N
Reminder about bonds and orbitals: we don't draw orbitals all the time. It's too much
trouble and anyway, when you draw lines and arrows you understand orbitals
are involved. Here's an example:
sp3
sp3 at N and B
empty p
N:
E
chemical bond
an sp3 orbital on the nitrogen atom of ammonia holds a lone e- pair.
Boron compounds are classic Lewis acids (electrophiles). B only uses
6 electrons to "complete" it's valence shell. It is sp2 hybridized and has an
empty 2p orbital.
The chemical bond above forms using a MO made by the overlap of
an sp3 hybrid AO and a pure p AO. B becomes sp3 hybridized with the bonding
to ammonia.
IV. Conjugate Acids and Bases
H:A + :Bacid
base
H:B +
conjugate
acid
:Aconjugate
base
Let's look at an L-B acid/base reaction.
When the base accepts the proton, the compound formed is called the
CONJUGATE ACID (the base becomes an acid)
When the acid donates the proton, the compound formed is called the
CONJUGATE BASE (the acid becomes a base)
We will re-visit this idea a little later but remember this:
The more stable (WEAKER) the conjugate base, the STRONGER the acid it came from.
The acid/base reaction will favor the forward direction.
The less stable (STRONGER) the conjugate base, the WEAKER the acid it came
from. The acid/base reaction will favor the reverse direction.
V. Acid and Base Strength (pKa and equilibrium)
Water undergoes autoionization. That is, it reacts with itself in an acid/base reaction
that produces hydronium and hydroxide ions.
H2O
H3 O +
+ H2 O
+ HO-
K w = [H 3O + ][HO ! ] = 10 !14
The ionization constant is defined as above. In water, the product of the [H3O+] and
[HO-] concentrations always equals 10-14.
Why care about water so much? Because relative acid and base strengths are
generally (though it isn't the only way) measured in water.
Hydronium and hydroxide ion concentrations are commonly placed on a log
scale called the pH scale. A log scale is a convenient way to compare very small
or very large numbers.
pH = ! log10 [H 3O + ]
The pH scale is usually bracketed from 0 to 14. Values below 7 are acidic and
values above are basic. 7 is neutral, where [H3O+] = [HO-]
The ionization of an acid in water...
HA
+ H2 O
H3 O +
+ A-
[H 3O + ][A ! ]
K eq =
[HA][H 2O]
K a = K eq [H 2O] =
[H 3O + ][A ! ]
[HA]
The bigger Ka, the stronger
the acid
The acid dissociation constant, Ka, is defined as the product of the reaction's
equilibrium constant, Keq and the concentration of water, whose presence as
a solvent means its concentration is essentially constant.
pKa is the -log10 of the acid dissociation constant. It scales like pH. The smaller
the value of pKa, the stronger the acid.
pK a = ! log10 K a
This diagram shows the relative acidity of several families of organic compounds.
Look, for example at the high pKa value for alkanes. This means that the conjugate
base, a C atom with a - charge on it, is an incredibly strong base.
H
H
C
H
H:B
:B-
C
STRONG
CONJUGATE
BASE
H
VI. Organic Acids and Bases
O
O
H
R
O
1
+
carboxylate group
H2O
+ H3O+
R
2
O
Structure 1 is the structure of the family of organic compounds called carboxylic acids.
Remember acetic acid from Chem 1? C2H3O2 ? That ought to look goofy to you now.
Can you draw the structure of acetic acid?
Anyway, as a family, carboxylic acids are relatively weak acids compared to mineral
acids like HCl, but are THE acids in organic chemistry.
Their pKa values are around 5, which means in water only a few % of it is dissociated
to ions.
The conjugate base, 2, of a carboxylic acid is stabilized by resonance. The more
stable that conjugate base is, the stronger (more dissociated) the acid.
R
O
O
O
!
R
O
R
O
O
resonance hybrid
shares the charge equally
over the two O atoms
Resonance means STABILITY!
Notice the conjugate base is negative. Can you think of a way to modify the structure
so that we STABILIZE it even more?
How about if we disperse it (spread it out) even more by adding an electronegative
atom that likes drawing e- toward it?
Let's draw a structure for the R group and modify it
O
O
H
vs.
H
H
H
C
H
O
pKa= 4.75
C
O
Cl
pKa = 2.86
Replacing one of the H atoms on the carbon next to the carboxylate group (the
position is called the " position) with a Cl will INCREASE the STABILITY
of the structure and hence INCREASE THE ACIDITY OF ITS PARENT ACID.
The Cl stabilizes by an ELECTRONIC EFFECT which acts THROUGH THE
BONDS. THIS IS AN INDUCTIVE EFFECT.
The further from the carboxylate group, the less effective the stabilization.
Let's move the Cl one C atom away (the position # to the carboxylate group).
H
O
H
C
Cl
C
H2
O
This Cl is less effective at stabilizing the
charge. pKa of the acid is 4.05
VII. Alcohols as acids and bases
Alcohols participate in a number of reactions. For now, let's focus on the hydroxyl
group as a L-B acid. The conjugate base, generically called alkoxides, are
important bases and nucleophiles in organic reactions.
The pKa of alcohols is around 16-17 so they are WEAK acids. The conjugate
base, the alkoxide, is a strong base.
R
O
+
O
B:-
+
R
H
H:B
alkoxide
alcohol
The alkoxide ion is not stabilized by resonance. However, electron withdrawing
groups in the R group will stabilize the ion and hence make the alcohol
a stronger acid.
This is stabilized by the Cl due to the electron withdrawing inductive
effect.
CH2 O
Cl
Alkyl groups tend to be weakly electron releasing. Hence, they will
destabilize the base by intensifying the - charge. This means the alkoxide
is a stronger base and its conjugate acid (alcohol) is a weaker acid.
n-butyl
O
O
O
sec-butyl
t-butyl
increasing base strength
decreasing nucleophilic strength
The bulky, STERICALLY hindered t-butyl group makes it hard for t-butoxide to
react like a nucleophile with a Lewis acid. But it is pretty easy for it to "pluck" out
a proton from a compound. In that case, it is acting as a L-B base.