Acid-Base Equilibria: Acids and Bases
What makes an Acid an Acid?
•An acid possess a sour taste
•An acid dissolves active metals magnesium
•An acid causes certain vegetable dyes to
turn characteristic colors
What makes a Base a Base?
•A bases possess a bitter taste
•A base feels slippery to the touch
•A base causes certain vegetable dues to
turn a characteristic color
7 strong acids and 8 strong
bases
• Acids - HI, HBr, HCl, HClO3,
HClO4, H2SO4, HNO3
• Bases – LiOH, NaOH, KOH,
RbOH, CsOH, Ca(OH)2,
Sr(OH)2, Ba(OH)2
The Arrhenius Definition of
an Acid and a Base
An acid is a substance that
produces H+ ions in water
solutions
HCl  H+ + ClA base is a substance that produces
OH- ions in a water solution
NaOH  Na+ + OH-
acid dissociation equations
HC6H5O3  C6H5O3
1- +
+
H
Fe(H2O)63+  Fe(H2O)5(OH)2+ + H+
CH3CH2NH31+  CH3CH2NH2 + H+
The Proton in Water
When HCl dissolves in water we write:
HCl(g)  H+(aq) + Cl-(aq)
Reality for the Hydronium ion
+
H5O2
H9O4
+
Acidic solutions are formed by a
chemical reaction in which
and acid transfers a proton (H+)
to water, so we can write them
either way.
+
HCl(aq) + H2O(aq)  H3O (aq) + Cl (aq)
or
HCl(aq)  H+(aq) + Cl-(aq)
Nitrogen compounds are
Bronsted acids when they are
protonated.
NH4Cl
+
+
NH4  NH3 + H
+
+
CH3)2NH2  (CH3)2NH + H
The Bronsted-Lowry definition
for Acids and Bases
Acids may be defined as a substance
that is capable of donating protons
Bases may be defined as substance that
accepts protons.
HCl + NH3
acid
base

NH4+ +
conjugate
acid
Clconjugate
base
Is Water an Acid?
+
NH3(aq) +H2O(aq)  NH4 (aq)
+
OH (aq)
Is Water a Base?
HC2H3O2(aq) + H2O(aq)  H3O+(aq) +C2H3O21-(aq)
The auto ionization of water
The reaction occurs to a very small extent;
about 1 in 108 molecules is ionized at any
given moment
H
O
+
H
H
H
 H
O
H
O
+
+
H
..
:O..
H
Dissociation of Water, pH Scale
H2O(l)  H+(aq) + OH-(aq)
K =[H+] [OH-]
[H2O]
since water is a liquid and its
concentration is therefore constant,
this expression may be written as:
Kw = [H+] [OH-] = 1.0 x 10-14
[H+] = [OH-] = 1.0 x 10-7 M
Sample exercise: Indicate whether
each of the following solutions is
neutral, acidic, or basic:
a) [H+] = 2 x 10-5 M
b) [OH-] = 0.010 M
c) [OH-] = 1.0 x 10-7 M
Calculate the concentration of H+(aq) in
(a)a solution in which the [OH-] is
0.020M
(b)a solution in which the
[OH-] = 2.5 x 10-6 M.
Indicate whether the solution is acidic or
basic
The pH Scale
pH = -log [H+]
If [H+] = 2. 5 x 105 the pH is?
-5
pH = -log [2. 5 x 10 ] = 4.6
+
H
If pH is 3.8 the
concentration is
Antilog -3.8= 1.58 x 10-4 M
In a sample of lemon juice, [H+] = 3.8 x
10-4 M. What is the pH?
A commonly available window cleaner
has a [H+] = 5.3 x 10-9 M. What is the
pH?
In a sample of freshly pressed apple juice
has a pH of 3.76. Calculate the [H+]
What if we took the –log of the Kw
expression
Kw = [H+] [OH-] = 1.0 x 10-14
pKw = pH + pOH = 14
What is the pH, [H+], [OH-], of
a solution with a pOH of 2.5?
Is the solution acidic or basic?
Major species
HCl(aq) + H2O(aq)  H3O+(aq) + Cl-(aq)
or
HCl(aq)  H+(aq) + Cl-(aq)
HC2H3O2(aq) + H2O(aq)  H3O+(aq) +C2H3O21-(aq)
Pb(NO3)2 + NaCl →NaNO3 + PbCl2
Indicators
What is the pH of 0.010 M
solution of HCl?
If it ionizes completely which is what
strong means then take the negative
log of the concentration.
HCl(aq)  H+(aq) + Cl-(aq)
.01M .01M
.01M
pH = 2
What is the pH of a solution
made from 20mL of 2.0M HCl
and 35mL of 3.2M HNO3?
What about H2SO4
+
H +
H2SO4 →
HSO4
1+
2HSO4 → H + SO4 (Weak)
1- (Strong)
What about weak acids
HX(aq) 
Ka =
+
H (aq)
+ X (aq),
+
[H ][X ]
then
[HX]
The smaller the value of the acid
dissociation constant Ka, the
weaker the acid
What is the Ka of a 0.10 M solution of
formic acid (HCHO2) which has a
pH = 2.38?
I
C
E
HCHO2  H+
+
0.10
0
.00417
.00417
.0958
.00417
CHO210
.00417
.00417
Ka = (.00417)2 = 1.8 x 10-4
.0958
What is the concentration of H+ ions
in a 0.10 M solution of HC2H3O2 (Ka
-5
= 1.8 x 10 )? pH? % ionization?
HC2H3O2 
I
.10
C
-X
E
.10 – X
1.8 x 10-5 =
X2
.10 – X
H+ + C2H3O210
0
+X
+X
X
X
X = 1.3 x 10-3
pH = 2.87
percent dissociation
1.3 X
.10
-3
10
x 100 = 1.3%
What is the pH and percent
ionization of a 0.20 M solution of
HCN? Ka = 4.9 x 10-10
Acid-Base Equilibria: Strong Bases
The most common soluble strong Bases are the hydroxides
of group IA and Ca, Ba and Sr
What is the pH of a 0.010 M
solution of Ba(OH)2?
Anions of Weak Acids
Amines
Dealing with Weak Bases
The base dissociation constant Kb refers to the
equilibrium in which a base reacts with H2O to
form the conjugate acid and OHWeak base + H2O  conjugate acid + OHNH3 (aq) + H2O (l)  NH4 +(aq) + OH-(aq)
+] [OH-]
[NH
4
Kb =
[NH3]
Calculate the [OH-] in a 0.15 M solution of NH3.
NH3 + H2O  NH4+ + OHI
C
E
Polyprotic Acids
H2SO3(aq) 
Ka1 = 1.7 x
HSO3-(aq)  H+(aq) + SO32-(aq)
Ka2 = 6.4 x 10-8
+
H (aq)
+ HSO3 (aq)
-2
10
Calculate the pH of a .1M solution
Anions of Weak Acids
HC2H3O2(aq) + H2O(aq)  H3O+(aq) + C2H3O2- (aq)
Bronsted
acid
Bronsted
base
Conjugate
acid
Conjugate
base
A second class of weak base is composed of the anions of
weak acids Anions of weak acids can be incorporated into
salts.
NaC2H3O2  Na+(aq) + C2H3O2- (aq)
C2H3O2- + H2O  HC2H3O2 + OH- Kb = 5.6 x 1010
NaOH(aq)+HC2H3O2(aq→H2O+NaC2H3O2
NaC2H3O2  Na+(aq) + C2H3O2-(aq)
Na+(aq) + H2O ← NaOH(aq) + H+(aq)
C2H3O2-(aq)+H2O HC2H3O2 (aq)+OH- (aq)
NH4Cl → NH4
+
+
NH4  NH3 + H
++
Cl
NH4Cl
NaC2H3O2
NH4C2H3O2
Anions of Weak Acids
Calculate the pH of a 0.01 M solution of sodium
hypochlorite (NaClO)
ClO- + H2O  HClO
I
C
E
+ OH-
Now it’s you turn: the Kb for BrO- is 5.0 x 10-6.
Calculate the pH of a 0.050 M solution of NaBrO
Ka and Kb
NH4+(aq)
Ka =
 NH3(aq)+
NH3(aq)+ H2O NH4+ (aq)+ OH-(aq)
+
H (aq)
[NH4][OH- ]
[H+][NH3]
Kb =
[NH4+]
[NH3]
NH4+(aq)  NH3(aq) + H+ (aq)
NH3(aq) + H2O(l) NH4+(aq) + OH- (aq)
H2O  H+(aq) + OH-(aq)
When two reactions are added to give a third reaction, the
equilibrium constant for the third reaction reaction is given
by the product of the equilibrium constants for the two
added reactions
pKa + pKb = pKw
K x K =K
a
b
w
Calculate the (a) base-dissociation constant, Kb, for the
fluoride ion, is the pKa of HF = 3.17
pKa = -log Ka
3.17 = -log Ka
Antilog -3.17 = 6.76 x 10-4
Since Ka x Kb = Kw
(6.76 x 10-4)x Kb = 1.0 x 10-14
Kb = 1.0 x 10-14/ 6.76 x 10-4 = 1.5 x 10-11
Calculate the pKb for carbonic acid (Ka = 4.3 x 10-7)
Now it’s your turn
Acid-Base Properties of Salt Solutions
•Anions of weak acids, HX, are basic and will react with
H2O to produce OH-
X- (aq) + H2O (l)  HX(aq) + OH-(aq)
•Anions of strong acids, such as NO3-, exhibit no basicitiy,
these ions do not react with water and consequently do not
influence the pH
•Anions of polyprotic acids, such as HCO3-, that still have
ionizable protons are capable of acting as either proton
donors or acceptors depending upon the magnitudes of the
Ka or Kb
•Anions of polyprotic acids, such as HCO3-, that still have
ionizable protons are capable of acting as either proton
donors or acceptors depending upon the magnitudes of the
Ka or Kb
Predict whether the salt Na2HPO4 will form an acidic or
basic solution on dissolvingin water.
Na2HPO4  2Na+ (aq) + HPO4HPO4- acting like an acid
HPO4 (aq) + H2O 
+ PO4 (aq)
HPO4- acting like an base
HPO4- (aq) + H2O  H2PO42-(aq) + OH-(aq)
-
H3O+
3-
K3 = 4.2 x 10-13
So HPO- is the conjugate base of H2PO4-.
Since the K2 of H2PO4- = 6.2 x 10-8 then:
=
Kw
Ka
1.0 x 10-14
= Kb =
6.2 x 10-8
1.6 x 10-7
Since Kb is larger than Ka, HPO4- will act like a base
•Salt derived from a strong base and a strong acid
will have a pH of 7
•Salt derived from a strong base and a weak acid
will have a pH above 7
•Salt derived from a weak acid and a weak base
depends upon whether the dissolved ion acts as
an acid or a base as determined by the size of
the Ka or Kb
Acid-Base Character and Chemical Structure
two things to consider polarity difference and
strength of the bond
HF > HCl > HBr > HI
(most polar
least) Based on electronegativity
difference HF is the most polar but a weak acid
because the bond is so strong
Acid strength of oxyacids
The more oxygen's the stronger
the acid because of the oxygen
pulling the electrons towards
themselves.
HClO4> HClO3> HClO2> HOCl