Chapter Two
Acids and Bases
Key words








Bronsted Acid
Arrhenius Acid
Electronegativity
Resonance
Size
Conjugate acid
Conjugate base
Lewis acid and Lewis base
Arrhenius Acids and Bases
• In 1884, Svante Arrhenius proposed.
• Acid: A substance that dissolves in water to produce H3O+
(hydroniums) ions.
• Base: A substance that dissolves in water to produce OHions.
• This definition of an acid is a slight modification of the
original Arrhenius definition, which was that an acid
produces H+ in aqueous solution.
• This definition fails when considering acids in non-aqueous
solution.
• Today we know that H+ reacts immediately with a water
molecule to give a hydronium ion H3O+.
Arrhenius Acids and Bases……
• HCl dissolves in water, to give a hydronium ion and a chloride
ion.
H
H
O
2
H
O
2
H
O
3
H
O
3
C
l
H
C
l
• We use curved arrows to show the change in position of
electron pairs during this reaction.
H
H
H2O
H
Cl
O
H
Hydronium ion
Cl
Arrhenius Acids and Bases……
• With bases, the situation is slightly different.
• Many bases are metal hydroxides such as KOH, NaOH,
Mg(OH)2, and Ca(OH)2.
• These compounds are ionic solids and, when they dissolve in
water, their ions merely/hardly separate.
N
a
O
H
(
s
)
H
O
2
N
a
(
a
q
)
O
H
(
a
q
)
• Other bases are not hydroxides; these bases produce OH- by
reacting with water molecules.
N
H
3
(
a
q
)
H
O
2
(
l
)
N
H
4
(
a
q
)
O
H
(
a
q
)
Mechanism showing reaction of ammonia with water molecules.
N
H
3
(
a
q
)
H
O
2
(
l
)
N
H
4
(
a
q
)
O
H
(
a
q
)
H
H
N
H
H
O
H
H
H
N
H
OH(aq)
H
Base
Equillibrium arrows means the forward reaction is rapidly reversible
Brønsted-Lowry Acids & Bases
• Acid: Is a proton (H+) donor.
• An acidic hydrogen is an H that is bonded to an
electronegative atom (e.g. N, O, F, Cl).
• Because this bond has such a strong dipole moment
(separation of charge), the H is easily pulled off by a base as
H +.
• Base: Is a proton (H+) acceptor.
• To identify, look for an atom that has lone pair(s) of
electrons (··) and/or a negative charge (−)!
Conjugate Acids & Bases
 Conjugate base: The species formed from an acid when an
acid donates a proton to a base.
• Conjugate acid: The species formed from a base when the
base accepts a proton from an acid.
• Acid-base reaction: A proton-transfer reaction.
• Conjugate acid-base pair: Any pair of molecules or ions that
can be interconverted by transfer of a proton.
C
onjugateAcid-Base
pair
H
2O
l
H C
Base
Acid
C
onjugateAcid-Base
pair
H
3O
C
onjugateAcid
of H
2O
C
l
C
onjugatebase
of H
C
l
Reaction mechanism of ethanoic acid (acetic acid) and ammonia :C on jug ate A cid -B ase
pair
O
H 3C
O
H
A cid
E th an oic acid
H
N
H
O
H
H
B a se
H 3C
H
H
O
C on juga te b ase
o f etha no ic acid
C o njug ate A cid -B ase
pair
N
H
C on jug ate A cid
o f am m onia
Brønsted-Lowry Acids & Bases
List of Examples of Acids and their Conjugate bases
Brønsted-Lowry Acids & Bases
• Note the following about the conjugate acid-base pairs in
Table 2.1:
1. An acid can be positively charged, neutral, or negatively
charged; examples of each type are H3O+, H2CO3, and
H2PO4- .
2. A base can be negatively charged or neutral; examples are
OH-, Cl-, and NH3.
3. Acids are classified as monoprotic, diprotic, or triprotic
depending on the number of protons that each may give
up; examples are HCl, H2CO3, and H3PO4.
Brønsted-Lowry Acids & Bases
• Carbonic acid, for example, can give up one proton to
become bicarbonate ion, and then the second proton to
become carbonate ion.
H2 CO3 + H2 O
Carbonic
acid
-
HCO3 + H2 O
Bicarbonate
ion
-
+
2-
+
HCO3 + H3 O
Bicarbon ate
ion
CO3
+ H3 O
Carbonate
ion
4. Several molecules and ions appear in both the acid and
conjugate base columns; that is, each can function as both
an acid and as a base.
Brønsted-Lowry Acids & Bases
5. There is an inverse relationship between the strength of an
acid and the strength of its conjugate base.
• The stronger the acid, the weaker its conjugate base.
• HI, for example, is the strongest acid in Table 2.1 and its
conjugate base, I-, is the weakest base in the table.
• CH3COOH (acetic acid) is a stronger acid than H2CO3 (carbonic
acid); conversely, CH3COO- (acetate ion) is a weaker base that
HCO3- (bicarbonate ion).
Acid and Base Strength
• Strong acid: One that completely or almost completely ionizes in
water to form H3O+ ions.
• Strong base: One that completely or almost completely ionizes
in water to form OH- ions.
• Here are the six most common strong acids (on left) and the
four most common strong bases (on right).
Formula
HCl
HBr
HI
HNO3
H2 SO4
HClO4
N ame
Hydrochloric acid
Hydrob romic acid
Hydroiodic acid
N itric acid
Su lfu ric acid
Perch loric acid
Formula
LiOH
NaOH
KOH
Ba( OH) 2
N ame
Lith iu m h yd roxide
Sodiu m hydroxide
Potass iu m h yd roxide
Barium hydroxide
Acid and Base Strength
• Weak acid: A substance that only partially dissociates in water to
produce H3O+ ions.
CH3 COOH(aq) + H2 O( l)
Acetic acid
98.7 %
-
CH3 COO ( aq) + H3O+ ( aq)
Acetate ion
1.3 %
• Weak base: A substance that only partially dissociates in water to
produce OH- ions.
• ammonia, for example, is a weak base
NH3 (aq) + H2 O( l)
NH4 + (aq) + OH-( aq)
Acid and Base Reactions
• Acetic acid is incompletely ionized in aqueous solution.
O
O
H2 O
CH3 COH +
Acid
Base
(w eaker acid ) (w eaker base)
-
CH3 CO
+
+
H3 O
Conju gate b ase
of CH 3CO 2H
(stron ger bas e)
Conju gate acid
of H 2O
(s tronger acid)
• The equation for the ionization of a weak acid, HA, is
-
+
HA + H2 O
A
+ H3 O
Ka = Keq[ H2 O] =
[H3 O ] [A ]
+
[HA]
pKa = −log10Ka
-
pKa = – log Ka
for Ka = 10–20, pKa = 20
for Ka = 10–5, pKa = 5
for Ka = 107, pKa = –7
The stronger the acid, the larger the Ka value
The stronger the acid, the smaller the pKa value
**The lower the pKa, the more acidic!**
Helpful Ranking of pKa values:
pKa
-8
5
9
~16
25
25
38
50
Compound
HCl
<------- most acidic
H3CCOOH
NH4+
ROH
H3CCOOCH3
HCCH
NH3
CH4
<-------- most basic
Acid-Base Equilibria
• To determine the position of equilibrium in an acid-base reaction:
1. Identify the two acids in the equilibrium; one on the left and
one on the right.
2. Use the information in Table 2.2 to determine which is the
stronger acid and which is the weaker acid.
3. Remember that the stronger acid gives the weaker conjugate
base, and the weaker acid gives the stronger conjugate base.
4. The stronger acid reacts with the stronger base to give the
weaker acid and weaker base.
5. Equilibrium lies on the side of the weaker acid and the
weaker base.
Acid-Base Equilibrium………..
• Equilibrium in the following acid-base reaction lies to the right, on
the side of the weaker acid.
O
CH3 COH
O
+
NH3
Acetic acid
Ammonia
pK a 4.76
(stron ger bas e)
(s tronger acid)
-
CH3 CO
+
NH4
+
A cetate ion Ammon ium ion
(w eak er b ase)
pK a 9.24
(w eaker acid )
Reaction will shift from strong side → weak side!
Acid-Base Equilibrium………..
• Equilibrium in the following acid-base reaction lies to the left, on
the side of the weaker acid.
C
H
C
H
O
H
3
2
w
e
a
k
e
ra
c
id
p
K
a=1
5
.9
+
C
H
N
H
3
2
C
H
C
H
O
3
2
+
Reaction will shift from strong side → weak side!
C
H
N
H
3
3
s
tro
n
g
e
ra
c
id
p
K
a=1
0
.7
Structure and Acidity
• The most important factor in determining the relative acidity of an
organic acid is the relative stability of the anion, A-, formed when
the acid, HA, transfers a proton to a base.
• We consider these four factors:
1. The electronegativity of the atom bonded to H in HA.
2. Resonance stabilization of A-.
3. The inductive effect.
4. The size & delocalization of charge (“polarizability”) in A- .
Structure and Acidity
1. ELECTRONEGATIVITY of the atom bearing the negative charge
• Pertains to elements going across a period…
• The greater the electronegativity of the atom bearing the
negative charge, the more strongly its electrons are held.
• The more strongly the electrons are held, the more stable
the anion A-.
• The more stable the anion A- the greater the acidity of the acid HA.
Structure and Acidity
2. RESONANCE delocalization of the charge on A• Compare the acidity of a carboxylic acid and an alcohol, both
of which contain an -OH group.
• Carboxylic acids are weak acids. Values of pKa for most
unsubstituted carboxylic acids fall within the range of 4 to 5.
CH3
COOH +
H2 O
CH3
- +
COO
+
H3 O
pKa = 4.76
• Alcohols are very weak acids. Values of pKa for most alcohols
fall within the range of 15 to 18.
CH3 CH2 O-H + H2 O
An alcohol
-
CH3 CH2 O
An alkoxide
ion
+ H3 O+
pKa = 15.9
How do we account for the fact that carboxylic acids are
stronger acids than alcohols?
O
H3C
O
O
H
H
O
H
H
H3C
O
Resonance
stabilized
O
H3C
H3C
H
Base
Acid
Ethanoic acid
H2
C
O
O
O
CANNOT BE STABILIZED BY RESONANCE DELOCALIZATION!!
H
Rank the following compounds from the least acidic to most
acidic, giving reasons, using 1 as (most acid)…..etc.
O
a.
O
H
Br
O
O
H
O
O
H
O
O
H
Br
Br
Br
O
O
H
H
H
H
O
O
O
O
b.
c.
Structure and Acidity
3. INDUCTIVE EFFECT
• …which is the polarization of electron density transmitted
through covalent bonds by a nearby atom of higher
electronegativity
Structure and Acidity
3. Inductive effect continued………………
• Electronegative groups PULL electron density in the molecule
towards themselves…
• This weakens the bond between the electronegative atom and
the acidic hydrogen and delocalizes (a.k.a. “spreads out”) the
resulting negative charge (after the H is pulled off).
• In other words, electronegative groups help to STABILIZE the A−
through delocalization of the negative charge.
• The inductive effect DECREASES, or dies off, as the
electronegative group gets farther away from the negative
charge!!
INCREASING ELECTRONEGATIVITY OF HALOGEN ATOM
INCREASING ACIDITY
O
O
O
<
<
<
O
C
H
C
H
C
H
C
O
H
C
H
C
H
C
H
C
O
H
C
H
C
H
C
H
C
O
H
3
2
3
2
C
H
C
H
C
H
C
O
H
3
2
3
2
I
B
r
C
l
F
The inductive effect DECREASES, or dies off, as the
electronegative group gets farther away from the negative
charge!!
INCREASING DISTANCE FROM MOST ACIDIC HYDROGEN ATOM
DECREASING ACIDITY
>
>
>
C
H
C
H
C
H
C
H
O
H
C
H
C
H
C
H
C
H
O
H
C
H
C
H
C
H
C
H
O
H
3
2
2
C
H
C
H
C
H
C
H
O
H
3
2
2
2
2
2
2
3
2
2
C
l
C
l
C
l
C
l
Structure and Acidity
4. SIZE -- the more stable the anion A- (the more “polarizable”),
the greater the acidity of the acid HA.
• The larger the volume over which the charge on an anion (or
cation) is delocalized, the greater the stability of the anion (or
cation).
• When considering the relative acidities of the hydrogen
halides, (HI > HBr > HCl > HF), we need to consider the
relative stabilities of the halide ions.
• Recall from general chemistry that atomic size is a periodic
property.
• For main group elements, atomic radii increase going down a
group and increase going across a period from right to left.
Electronegativity
Electronegativity and size
Atomic Size (radius)
(Cesium is the largest)
Structure and Acidity
• For the halogens, iodine has the largest atomic radii, fluorine has
the smallest I > Br > Cl > F.
Anions are always larger than the atoms from which they are
derived. For anions, nuclear charge is unchanged but the added
electron(s) introduce new repulsions and the electron clouds
swell.
Among the halide ions, I- has the largest atomic radius, and F- has
the smallest atomic radius.
Thus, HI is the strongest acid because the negative charge on the
iodide ion is delocalized over a larger volume than the negative
charge on chloride, etc.
Note that the result of both resonance and the inductive effect is
due to the delocalization of charge.
Structure and Acidity
• Notice that the electronegativity and size trends contradict one
another!
• When comparing the relative acidities of different elements,
SIZE beats electronegativity!!
• When comparing the relative acidities of the same element in
different compounds, look at the following in this order of
importance:
1. Charge (greater (+) charge = more acidic)
2. Resonance (more resonance = more acidic)
3. Induction (nearby electronegative groups = more acidic)
Lewis Acids and Bases
• Lewis acid: Any molecule or ion that can form a new covalent
bond by accepting a pair of electrons.
• Lewis Base: Any molecule or ion that can form a new covalent
bond by donating a pair of electrons.
A
Lewis
acid
+
:B
Lewis
base
-
+
A B
new covalent bond
formed in this Lewis
acid-base reaction
Lewis Acids and Bases
When HCl dissolves in water, for example, the strongest available
Lewis base is an H2O molecule, and the following proton-transfer
reaction takes place.
H
O
H
H
H
Cl
O
H
H
Hydronium
ion
+
Cl
When HCl dissolves in methanol, the strongest available Lewis
base is a CH3OH molecule, and the following proton-transfer
reaction takes place.
H
O
H
C H 3
M e th a n o l
H
C l
O
H
3
C
+
H
O x o n iu m
io n
C l
Lewis Acids and Bases
Why is −NH2 a stronger base than NH3? (Or −OH stronger than H2O?)
Why is −NH2 a stronger base than −OH? (Or NH3 stronger than H2O?)
Lewis Acids and Bases
• Another type of Lewis acid we will encounter in later chapters is
an organic cation in which a carbon is bonded to only three
atoms and bears a positive formal charge.
• Such carbon cations are called carbocations
• Remember that a cation is DEFICIENT in electron(s);
therefore, it desires electrons!
-
Lewis defined acids as electron pair acceptors and bases as electron pair
donors. This broadens the definitions for acid and base: Any electron deficient
atom can act as a Lewis acid. For example, BF3 functions as an acid since
boron is electron deficient.
In BF3 the number of electrons around the central atom are 6 .It is electron deficient
and has a tendency to accept a pair of electrons to achieve stable octet state
configuration and thus behave as a Lewis acid
In AlCl3 the number of electrons around the central atom
are 6 .It is electron deficient and has a tendency to accept a pair
of electrons to achieve stable octet state configuration and thus
behave as a Lewis acid
Chapter Two
Acids and Bases
Key words








Bronsted Acid
Arrhenius Acid
Electronegativity
Resonance
Size
Conjugate acid
Conjugate base
Lewis acid and Lewis base