General Chemistry II
2302102
Acid and Base Equilibria
Lecture 1
i.fraser@rmit.edu.au
Ian.Fraser@sci.monash.edu.au
Acids and Bases
- 3 Lectures
Outline - 5 Subtopics
• Autoionization of Water and pH
• Defining Acids and Bases
• Interaction of Acids and Bases with Water
• Buffer Solutions
• Acid-Base Titrations
Acids and Bases
Objectives - Lecture 1
By the end of this lecture AND completion of the set
problems, you should be able to:
• Understand strong and weak electrolytes, Kw and the
autoionization equilibrium in water, definition of the pH scale
and its relationship to [H3O+] and [OH-].
• Know the Arrhenius and Brønsted definitions of acids and
bases.
• Understand and know examples of conjugate acid-base pairs.
• Be familiar with the ionization reactions for strong and weak
acids, know examples of typical monoprotic, diprotic and
triprotic acids.
• Understand the common ion effect in acid ionization.
Acids & Bases
Would you evacuate?
Gold mining frequently uses the base, cyanide (CN-), in the
extraction process. Tailings dams sometimes have high
concentrations of CN-. The conjugate acid of CN-, HCN, is
extremely toxic.
“Acid mine drainage” (AMD) is naturally produced by the
exposure of sulfide ores to water. This process is greatly
exacerbated by mining the sulfide ores (for Cu, Pb, Zn etc).
AMD results in streams with elevated H+ concentrations and
pH values in the range 1-3.
A gold mining tailings dam leaks into an AMD-affected
stream - would you evacuate?
Acids & Bases - 2 Lectures
Outline
Introduction to acids and bases
Strong & Weak Acids
Conjugate acid-base pairs
Common Ion Effect
Bases
Buffers
Indicators
Titrations
• Strong Acids
• Weak Acids
Acids & Bases - Lecture 1
Objectives
By the end of this lecture AND completion of the set
problems, you should be able to:
•
•
•
•
Define BrØnsted acids and bases
Calculate [H+], [OH-] and pH
Distinguish between strong & weak acids
Calculate equilibrium concentrations of acids &
bases using acidity constants
• Identify conjugate acid-bases pairs
• Determine the effect of adding a common ion on
equilibrium concentrations
Acids & Bases - Why are we interested?
Large Range of Industrial Processes:
Acids - solvents (dissolve other materials)
- ores, food (stomach contains HCl)
Bases - solvents - cleaners like “Draino”, bleach
Foods:
Acids - wine, beer, citrus fruits, vinegar, coffee
Physiology:
Blood - pH 7.3-7.5 falls below 6.8 will be fatal (acidosis)
Most bodily functions under ‘circumneutral’ conditions
Environment:
‘Natural’ Erosion of Limestone Caves
“acid rain” - dissolved H2SO4 & HNO3 in upper atmosphere
- rainfall runoff into acid lakes - devoid of life
Acids & Bases
BrØnsted-Lowry Definitions:
1. ACID
Species which can donate a hydrogen ion (H+).
2. BASE
Species which can accept a hydrogen ion.
3. AMPHOTERIC SPECIES
Species which can act as both an acid and a
base. e.g. HCO3- (CO32-, H2CO3), HSO4-
What is the Hydrogen Ion?
• Hydrogen Atom
Atomic weight = 1
1 proton + 1 electron
• Hydrogen Ion
Atomic weight = 1
1 proton
• Hydrogen Ion in
Water
Represented as:
H+(aq) or H3O+
Acid-Base Reaction
In fact a proton-transfer reaction in which the
proton is transferred from the acid to the base.
Acidity - pH
Acidity measure of the H3O+ concentration
pH convenient measure of H3O+ concentration


H 3O   10 pH
concentration of H 3O   10 pH M
e.g.
pH = 8.5


H 3O   10  8.5
 3.16  10  9
concentration of H 3O   3.16  10  9 M
AQUEOUS SYSTEMS
2H2O(l)
H3O+(aq) + OH-(aq)
Equilibrium between H3O+ and OH-:
at equilibrium [H3O+][OH-] = Kw
at 25 °C
Kw = 1.00 x 10-14
(c.f. Ksp = [Ag+][Cl-])
AQUEOUS SYSTEMS
H3O+(aq) + OH-(aq)
2H2O(l)
Kw = [H3O+][OH-]
at 25 °C
Kw = 1.00 x 10-14
Pure water at 25 °C. If 2z mol/L of water react:
[H3O+] = z
and
[OH-] = z
z2 = 1.00 x 10-14 \ z = 1.00 x 10-7
and pH = 7.00
DEFINITIONS
“acid”
pH < 7.00
[H3O+(aq)] > [OH-(aq)]
“basic”
pH > 7.00
[H3O+(aq)] < [OH-(aq)]
“neutral”
pH = 7.00
[H3O+(aq)] = [OH-(aq)]
Acid-Base Reaction: Proton-Transfer
between an Acid and a Base (Water)
• According to the Brönsted-Lowry concept, when HCl
gas is dissolved in water to form the solution of
hydrochloric acid, a proton-transfer reaction occurs:
H3O+ (aq) + Cl- (aq)
H2O (l) + HCl (g)
H O: + H
Cl :
H
O
H
+
+ :Cl:
H
H
Acid
(Proton
Base
Hydronium
(Proton Donor)
Ion
Acceptor)
A hydrated proton = H+ (aq) = H3O+
-
ACIDS
ACIDS
Reaction between acids and water
H3O+(aq) + A-(aq)
H2O(l) + HA(aq)
Equilibrium constant Ka
At equilibrium
A




( aq) H 3O ( aq)
 Ka
HA( aq)
Note HA may be a molecule
anion
or cation
HCl( aq), HCOOH( aq)
HCO 3 ( aq)
NH 4 ( aq)
ACIDS
Reaction between acids and water
H3O+(aq) + A-(aq)
H2O(l) + HA(aq)
Equilibrium constant Ka
Usually tabulated as pKa:
pK a   log10 K a 
so that
K a  10
pK
a
Note - this is exactly the same relationship as
between [H+] and pH (pH = - log10 [H+])
Strength of Acids
H2O(l) + HA(aq)
H3O+(aq) + A-(aq)
• Strong versus Weak acids
• The strength of an acid is related to the
position of the equilibrium above
• It is NOT related to the corrosive ability (this
often causes confusion)
• As we shall see, HF is a weak acid, yet it is
one of the most corrosive acids known.
STRONG ACID
Reaction between acids and water
H2O(l) + HA(aq)
H3O+(aq) + A-(aq)
Position of equilibrium almost completely to the
right (acid is almost totally ionized)
e.g. 0.1 M HCl(aq)
[H3O+]
[Cl-]
[HCl]
pKa
= 0.1
= 0.1
not known ( < 10-10)
< -10
Acids are Electrolytes
Non
Strong
Weak
No ions
in solution
Many ions
in solution Strong Acids
Few ions
in solution Weak Acids
WEAK ACIDS
Reaction between acids and water
H2O(l) + HA(aq)
1.
H3O+(aq) + A-(aq)
Definition of weak acid
Position of equilibrium lies to the left. Only
a small fraction of the acid reacts with the water
and is ionized. i.e. Acid is weakly ionized
Ka small < 10-3
pKa large > 3
WEAK ACIDS
2.
Dissociation in water
H2O(l) + HA(aq)
H3O+(aq) + A-(aq)
eg. 0.5 M HF(aq) Ka = 6.8 x 10-4
If x mol/L react then we have
H2O(l) + HF(aq)
initial
0.5 M
equilibrium (0.5 - x) M
H3O+(aq) + F-(aq)
0
xM
0
xM
WEAK ACIDS
H2O(l) + HF(aq)
initial
0.5 M
equilibrium (0.5 - x) M
H3O+(aq) + F-(aq)
0
xM
0
xM
At equilibrium:
F  (aq)H 3O  (aq)  6.8  104
HF (aq)
x2
 6.8  104
0.5  x 
x  1.8  10 2
[ F  ]  1.8  10 2 M

[ H 3O ]  1.8  10
2
M
[ HF ]  0.48M
So HF is ca. 4% ionized
Diprotic & Triprotic Acids
• Acids that we’ve considered thus far have been
monoprotic (donate 1 proton) e.g. HCl, HNO3
• Other acids can donate 2 or 3 protons diprotic or triprotic acids e.g. H2SO4, H3PO4
H2O(l) + H2A(aq)
H3O+(aq) + HA-(aq)
[H 3 O ] [HA  ]
K a1 
[H 2 A]
H2O(l) + HA-(aq)
K a2 
H3O+(aq) + A2-(aq)
[ H 3 O  ] [ A 2 ]
[HA ]
Diprotic & Triprotic Acids
[H 3 O ] [HA  ]
K a1 
[H 2 A]
Acid
Formula
Ka1
Sulfuric
H2SO4
ca. 100
Oxalic
H2C2O4
5.9 x 10
Phosphoric H3PO4
7.52 x 10
K a2 
pKa1
-2
-3
Ka2
[HA ]
pKa2
ca. -2 1.2 x 10
-2
6.4 x 10
-5
1.23
[ H 3 O  ] [ A 2 ]
2.12 6.23 x 10
-8
Ka3
pKa3
1.92
4.19
-13
7.21 2.2 x 10
Successive pKas increase in magnitude
- removal of successive protons results in weaker
acids
12.67
Conjugate Acid-Base Pairs
Consider the proton-transfer reaction:
– Forward reaction
H2O (l) + NH3 (aq)
NH4+ (aq) + OH- (aq)
Acid
(Proton
Donor)
Base
(Proton
Acceptor)
- Reverse reaction
H2O (l) + NH3 (aq)
NH4+ (aq) + OH- (aq)
Acid
(Proton
Donor)
Base
(Proton
Acceptor)
Conjugate Acid-Base Pairs
Conjugate AcidBase Pair
H2O (l)
Acid 1
+
NH3 (aq)
Base 1
NH4+ (aq)
Acid 2
Conjugate AcidBase Pair
+ OH- (aq)
Base 2
Conjugate Acid-Base Pairs
For the general equation
Base
(Proton
Acceptor)
B
+
Acid
(Proton
Donor)
HA
HB+
Acid
(Proton
Donor)
+
A-
Base
(Proton
Acceptor)
The acid HB+ results from the base B gaining a
proton. HB+-B is a conjugate acid-base pair.
The base A- results from the acid HA losing a
proton. HA-A- is a conjugate acid-base pair, too.
Conjugate Acid-Base Pairs
• Any two substances that differ by one proton
are a conjugate acid-base pair
• Can write the formula of the conjugate base of
any acid simply by removing the proton:
If HCO3- is the acid, CO32- is the conjugate base
If HCO3- is the base, H2CO3 is the conjugate acid
(HCO3- is amphoteric)
Conjugate Acid-Base Pairs
• The conjugate (base) of a strong acid is a
weak base
• The conjugate (base) of a weak acid is a
strong base
• The conjugate (acid) of a strong base is a
weak acid
• The conjugate (acid) of a weak base is a
strong acid
Direction of Acid-Base Reactions
• Acid-Base Reactions proceed spontaneously
with the strongest acid and strongest base
forming the weakest acid and the weakest
base
Returning to:
H2O (l)
+
NH3 (aq)
NH4+ (aq)
+ OH- (aq)
NH4+ is a stronger acid than H2O, and
OH- is a stronger base than NH3
So reaction proceeds spontaneously to the left
The Common Ion Effect
Recall that with sparingly soluble salts:The
presence of an ion in solution which is common
to the electrolyte will decrease the solubility
(adding KF to a solution of PbF2 reduced the Pb2+ concentration)
The presence of an ion in solution which is
common to the weak acid will suppress it’s
ionization (decrease the concentration of the free
ion)
The Common Ion Effect
H2O(l) + HF(aq)
initial
0.5 M
equilibrium (0.5 - x) M
H3O+(aq) + F-(aq)
0
xM
0
xM
At equilibrium, [F-] = [H3O+] = 1.8 x 10-2 M
so pH = 1.74
What happens if we add 0.05 M NaF?
H2O(l) + HF(aq)
initial
0.5 M
equilibrium (0.5 - x) M
H3O+(aq) + F-(aq)
0
xM
0.05
0.05 + x M
The Common Ion Effect
H2O(l) + HF(aq)
initial
0.5 M
equilibrium (0.5 - x) M
H3O+(aq) + F-(aq)
0
xM
0.05
0.05 + x M
F  (aq)H 3O  (aq)  6.8  104
HF (aq)
So (0.05 + x) x = (0.5 - x) 6.8 x 10-4
x = 6.00 x 10-3 M So [H3O+] = 6.00 x 10-3 M
Thus pH = 2.22 c.f. 1.74 without the NaF
Acids and Bases - End of Lecture 1
Objectives Covered in Lecture 1
After studying this lecture and the set problems,
you should be able to:
• Understand strong and weak electrolytes, Kw and the
autoionization equilibrium in water, definition of the pH
scale and its relationship to [H3O+] and [OH-].
• Know the Arrhenius and Brønsted definitions of acids and
bases.
• Understand and know examples of conjugate acid-base
pairs.
• Be familiar with the ionization reactions for strong and weak
acids, know examples of typical monoprotic, diprotic and
triprotic acids.
• Understand the common ion effect in acid ionization.