CHAPTER 9
Acids & Bases
General, Organic, & Biological
Chemistry
Janice Gorzynski Smith
CHAPTER 9: Acids & Bases
Learning Objectives:
 Define Acids & Bases: Bronsted-Lowry
 Understand the difference between strong & weak acids
 Hydronium ion formation
 Identify conjugate acids & bases
 Define & calculate Ka, Kb, Kw
 Water is both an acid & a base
 Titrations to calculate analyte concentration
 Buffers
 Weak acid + conjugate base
 Weak base + conjugate acid
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Smith. General Organic & Biologicial Chemistry 2nd Ed.
Acid/Base
Brønsted-Lowry Definition
The Brønsted–Lowry definition is more widely used:
•A Brønsted–Lowry acid is a proton (H+) donor.
•A Brønsted–Lowry base is a proton (H+) acceptor.
This proton is donated.
HCl(g) + H2O(l)
H3O+(aq) + Cl−(aq)
•HCl is a Brønsted–Lowry acid because it donates
a proton to the solvent water.
•H2O is a Brønsted–Lowry base because it accepts
a proton from HCl.
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Acid/Base
Brønsted-Lowry Definition
A Brønsted–Lowry acid must contain a hydrogen atom.
•A monoprotic acid contains one acidic proton.
HCl
•A diprotic acid contains two acidic protons.
H2SO4
•A triprotic acid contains three acidic protons.
H3PO4
•A Brønsted–Lowry acid may be neutral or it may
carry a net positive or negative charge.
HCl, H3O+, HSO4−
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Brønsted-Lowry Definition
Acid/Base
A base must contain a lone pair of electrons that
can be used to form a new bond to the proton.
This e− pair forms a new
bond to a H from H2O.
H
N
H
+
H2O(l)
H
H
N
H
+ OH− (aq)
H
Lone pairs make these
neutral compounds bases.
NH3
ammonia
+
H
H2O
water
Smith. General Organic & Biologicial Chemistry 2nd Ed.
The OH− is the base in
each metal salt.
NaOH
sodium
hydroxide
KOH
potassium
hydroxide
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Acid/Base
This e− pair
stays on A.
H A
acid
Brønsted-Lowry Definition
This e− pair forms
a new bond to H+.
gain of H+
+
B
base
A
−
+
H
B+
loss of H+
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Conjugate Acids & Bases
Acid/Base
gain of H+
H A
acid
+
B
base
A − +
H B+
conjugate conjugate
acid
base
loss of H+
•The product formed by loss of a proton from an acid is called its
conjugate base.
•The product formed by gain of a proton by a base is called
its conjugate acid.
H Br
acid
+
H2O
base
Smith. General Organic & Biologicial Chemistry 2nd Ed.
Br− +
H3O+
conjugate conjugate
acid
base
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Water: An Acid & Base
Acid/Base
Amphoteric compound: A compound that contains
both a hydrogen atom and a lone pair of e−; it can
be either an acid or a base.
+
H
H
O
H
add H+
H2O as a base
H
O
H
H2O as an acid
Smith. General Organic & Biologicial Chemistry 2nd Ed.
H
O
H
conjugate acid
remove
H+
−
H
O
conjugate base
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Acid/Base
Strong Acids
•When a covalent acid dissolves in water, the proton
transfer that forms H3O+ is called dissociation.
•When a strong acid dissolves in water, 100% of
the acid dissociates into ions.
HCl(g) + H2O(l)
H3O+(aq) + Cl−(aq)
•A single reaction arrow is used, because the
product is greatly favored at equilibrium.
•Common strong acids are HI, HBr, HCl, H2SO4,
and HNO3.
Smith. General Organic & Biologicial Chemistry 2nd Ed.
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Acid/Base
Weak Acids
•When a weak acid dissolves in water, only a
small fraction of the acid dissociates into ions.
•Unequal reaction arrows are used, because the
reactants are usually favored at equilibrium.
CH3COOH(l) + H2O(l)
H3O+(aq) + CH3COO−(aq)
•Common weak acids are H3PO4, HF, H2CO3, and
HCN.
Smith. General Organic & Biologicial Chemistry 2nd Ed.
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Acid/Base
Strong & Weak Bases
•When a strong base dissolves in water, 100% of
the base dissociates into ions.
NaOH(s) + H2O(l)
Na+(aq) +
−OH(aq)
•Common strong bases are NaOH and KOH.
•When a weak base dissolves in water, only a
small fraction of the base dissociates into ions.
NH3(g) + H2O(l)
Smith. General Organic & Biologicial Chemistry 2nd Ed.
NH4+(aq) + −OH(aq)
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Acid/Base
Strong/Weak Acids/Bases
•A strong acid readily donates a proton, forming
a weak conjugate base.
HCl
strong acid
Cl−
weak conjugate base
•A strong base readily accepts a proton, forming
a weak conjugate acid.
OH−
strong base
Smith. General Organic & Biologicial Chemistry 2nd Ed.
H2O
weak conjugate acid
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Acid/Base
Strong/Weak Acids/Bases
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Ka & Kb
Acid/Base
•A Brønsted–Lowry acid–base reaction represents
an equilibrium.
H A
acid
+
B
base
A − +
H B+
conjugate conjugate
acid
base
•The position of the equilibrium depends upon the
strengths of the acids and bases.
•The stronger acid reacts with the stronger base
to form the weaker acid and the weaker base.
Smith. General Organic & Biologicial Chemistry 2nd Ed.
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Acid/Base
Ka & Kb
HA(g) + H2O(l)
H3O+(aq) + A
− (aq)
[H3O+][ A −]
Ka =
[HA]
acid dissociation
constant
•The stronger the acid, the larger the Ka value.
•Equilibrium favors formation of the weaker acid,
the acid with the smaller Ka value.
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Acid/Base
Ka & Kb
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Acid/Base
Ka & Kb
OH- (aq) + BH+ (aq)
B (g) + H2O(l)
Kb
=
[OH- ][BH+]
[ B]
Base dissociation
constant
•The stronger the base, the larger the Kb value.
Smith. General Organic & Biologicial Chemistry 2nd Ed.
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Kw
Acid/Base
loss of H+
H
O
H + H
acid
O
H
base
+
H
−
O
H
+
conjugate
base
H
O
H
conjugate
acid
gain of H+
ion-product
constant
Kw = [H3O+][OH−]
Kw is a constant for
all aqueous
solutions at 25 oC.
Kw
=
(1.0 x 10−7) x (1.0 x 10−7)
Kw
=
1.0 x 10−14
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Smith. General Organic & Biologicial Chemistry 2nd Ed.
Acid/Base
pH
The lower the pH, the higher the concentration of H3O+:
•Acidic solution:
pH < 7  [H3O+] > 1 x 10−7
•Neutral solution:
pH = 7  [H3O+] = 1 x 10−7
•Basic solution:
pH > 7  [H3O+] < 1 x 10−7
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Acid/Base
pH
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Acid/Base
Common Acid/Base Reactions
Neutralization reaction: An acid-base reaction that
produces a salt and water as products.
HA(aq) + MOH(aq)
base
acid
H
OH(l) + MA(aq)
water
salt
•The acid HA donates a proton (H+) to the OH− base
to form H2O.
•The anion A− from the acid combines with the
cation M+ from the base to form the salt MA.
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Acid/Base
Common Acid/Base Reactions
A net ionic equation contains only the species involved
in a reaction.
HCl(aq) + NaOH(aq)
H—OH(l) + NaCl(aq)
•Written as individual ions:
H+(aq) + Cl−(aq) + Na+(aq) + OH− (aq)
H—OH(l) + Na+(aq) + Cl−(aq)
•Omit the spectator ions, Na+ and Cl–.
•What remains is the net ionic equation:
H+(aq) + OH− (aq)
H—OH(l)
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Smith. General Organic & Biologicial Chemistry 2nd Ed.
Acid/Base
Common Acid/Base Reactions
•A bicarbonate base, HCO3−, reacts with one H+ to
form carbonic acid, H2CO3.
H+(aq) + HCO3−(aq)
H2CO3(aq)
H2O(l) + CO2(g)
•Carbonic acid then decomposes into H2O and CO2.
•For example:
HCl(aq) + NaHCO3(aq)
NaCl(aq) + H2CO3(aq)
H2O(l) + CO2(g)
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Acid/Base
Common Acid/Base Reactions
•A carbonate base, CO32–, reacts with two H+ to
form carbonic acid, H2CO3.
2 H+(aq) + CO32–(aq)
H2CO3(aq)
H2O(l) + CO2(g)
•For example:
2 HCl(aq) + Na2CO3(aq)
2 NaCl(aq) + H2CO3(aq)
H2O(l) + CO2(g)
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Acid/Base
Titration
Initial
Midway
Final
AH + B  A- + BH+
Acid + Base  Conjugate Base + Conjugate Acid
HCl + NaOH  NaCl + H2O
H+ (acid) + OH- (base)  H2O (neutral water)
Calculations:
Know: the concentration of Base & the Volume of Base added.
1. VBase (L) x MBase (moles/L) = moles Base added to the acid
2. Mole : Mole ratio of Acid : Base is 1:1 (if monoprotic), therefore
moles Base added = moles of Acid reacted
3. Determine the Molarity of the Acid solution initially by dividing the #
moles that reacted with the initial volume of the solution (L).
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Acid/Base
Titration Caclulations
mole–mole
conversion
factor
Moles of
base
M (mol/L)
conversion
factor
[1]
Volume of
base
Smith. General Organic & Biologicial Chemistry 2nd Ed.
[2]
Moles of
acid
[3]
M (mol/L)
conversion
factor
Volume of
acid
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Acid/Base
Buffers
A buffer is a solution whose pH changes very little
when acid or base is added.
Most buffers are solutions composed of roughly
equal amounts of:
•A weak acid
•The salt of its conjugate base
The buffer resists change in pH because
•Added base, OH−, reacts with the weak acid
•Added acid, H3O+, reacts with the conjugate base
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Acid/Base
Buffers
If an acid is added to the following buffer equilibrium,
then the excess acid reacts with the conjugate base,
so the overall pH does not change much.
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Acid/Base
Buffers
If a base is added to the following buffer equilibrium,
then the excess base reacts with the conjugate acid,
so the overall pH does not change much.
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Acid/Base
Buffers
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Acid/Base
Buffers
The effective pH range of a buffer depends on its Ka.
H3O+(aq) + A −(aq)
HA(aq) + H2O(l)
[H3O+][ A −]
[HA]
Ka =
Rearranging this expression to solve for [H3O+]:
[H3
O+]
=
Ka
x
[HA]
[ A −]
determines the
buffer pH
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Acid/Base
Buffers
•Normal blood pH is between 7.35 and 7.45.
•The principle buffer in the blood is carbonic acid/
bicarbonate (H2CO3/HCO3−).
CO2(g) + H2O(l)
H2CO3(aq)
H2O
H3O+(aq) + HCO3−(aq)
•CO2 is constantly produced by metabolic
processes in the body.
•The amount of CO2 is related to the pH of the blood.
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