Chapter
Acids & Bases
8
COURSE NAME: CHEMISTRY 101
COURSE CODE: 402101-4
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
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1-Arrhenius Concept of Acids and Bases
Acid : is a substance that, when dissolved in
water, increases the concentration of
hydronium ion (H3O+).
A base: in the Arrhenius concept, is a substance that, when
dissolved in water, increases the concentration of hydroxide
ion, OH-(aq).
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2- Brønsted-Lowry Concept of Acids
and Bases
• Acid: is the species donating the proton in
a proton-transfer reaction.
A base : is the species accepting the proton 
in a proton-transfer reaction.
– In any reversible acid-base reaction, both forward
and reverse reactions involve proton transfer.
2
Consider the reaction of NH3 and H2O.
base
acid
NH 3 (aq )  H 2O( l )

NH 4 (aq )  OH  (aq )
H+
– In the forward reaction, NH3 accepts a proton
from H2O. Thus, NH3 is a base and H2O is an
acid.
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3- Lewis Concept of Acids and
Bases
The Lewis concept defines an acid as an electron pair
acceptor and a base as an electron pair donor.
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3- Lewis Concept of Acids and Bases
B
+ :N H
:F :
H
:F
:F :
H
B
N
:F :
H
:
H
:
:F :
: :
: :
:F
:
:
The reaction of boron trifluoride with ammonia is an
example.
H
– Boron trifluoride accepts the electron pair, so it is a
Lewis acid. Ammonia donates the electron pair,
so it is the Lewis base.
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Self-ionization of Water
Self-ionization is a reaction in which two like
molecules react to give ions.
– In the case of water, the following equilibrium is
established.
H 2O ( l )  H 2O ( l )
H 3O  (aq )  OH  (aq )
– The equilibrium-constant expression for this
system is:


[ H 3O ][OH ]
Kc 
2
[ H 2O ]
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Self-ionization is a reaction in which two like
molecules react to give ions.
– The concentration of ions is extremely
small, so the concentration of H2O remains
essentially constant. This gives:


[ H 2O ] K c  [ H 3O ][OH ]
2
constant
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– Because we often write H3O+ as H+, the ionproduct constant expression for water can be
written:


K w  [ H ][OH ]
– Using Kw you can calculate the concentrations of
H+ and OH- ions in pure water.
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Self-ionization of Water
These ions are produced in equal numbers in pure 
water, so if we let x = [H+] = [OH-]
1.0  10
14
 ( x )( x )
x  1.0  10
14
o
at 25 C
 1.0  10
7
– Thus, the concentrations of H+ and OH- in pure
water are both 1.0 x 10-7 M.
– If you add acid or base to water they are no longer
equal but the Kw expression still holds.
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Solutions of Strong Acid or Base
By dissolving substances in water, you can alter the
concentrations of H+(aq) and OH-(aq).
– In a neutral solution, the concentrations of H+(aq)
and OH-(aq) are equal, as they are in pure water.
– In an acidic solution, the concentration of H+(aq) is
greater than that of OH-(aq).
– In a basic solution, the concentration of OH-(aq) is
greater than that of H+(aq).
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At 25°C, you observe the following conditions.
– In an acidic solution, [H+] > 1.0 x 10-7 M.
– In a neutral solution, [H+] = 1.0 x 10-7 M.
– In a basic solution, [H+] < 1.0 x 10-7 M.
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1
The pH of a Solution
Although you can quantitatively describe the acidity of a
solution by its [H+], it is often more convenient to give
acidity in terms of pH.
– The pH of a solution is defined as the negative
logarithm of the molar hydrogen-ion concentration.

pH   log[ H ]
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For a solution in which the hydrogen-ion concentration
is 1.0 x 10-3, the pH is:
3
pH   log(1.0  10 )  3.00
– Note that the number of decimal places in
the pH equals the number of significant
figures in the hydrogen-ion concentration.
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The pH of a Solution
In a neutral solution, whose hydrogen-ion concentration 
is 1.0 x 10-7, the pH = 7.00.
• For acidic solutions, the hydrogen-ion
concentration is greater than 1.0 x 10-7, so the
pH is less than 7.00.
• Similarly, a basic solution has a pH greater
than 7.00.
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Figure : The pH Scale
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A sample of orange juice has a hydrogen-ion
concentration of 2.9 x 10-4 M. What is the pH?

pH   log[ H ]
4
pH   log( 2.9  10 )
pH  3.54
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A measurement of the hydroxide ion concentration, 
similar to pH, is the pOH.
– The pOH of a solution is defined as the
negative logarithm of the molar hydroxideion concentration.

pOH   log[OH ]
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A measurement of the hydroxide ion concentration,
similar to pH, is the pOH.
– Then because Kw = [H+][OH-] = 1.0 x 10-14
at 25 oC, you can show that
pH  pOH  14.00
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Example : An ammonia solution has a hydroxide-ion
concentration of 1.9 x 10-3 M. What is the pH of the
solution?
You first calculate the pOH:
3
pOH   log(1.9  10 )  2.72
Then the pH is:
pH  14.00  2.72  11 .28
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1
Weak acid and base calculations
pH of strong acids and bases
HA
H+ + A- complete dissociation of an acid
c(HA) = [H+] = [A-]
pH = -log a(H+)
a – activity
a(H+) = γ±·c(HA)
γ± - mean activity coefficient
pH = -log[H+]
In very diluted solutions: γ± = 1!
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pH of strong acids and bases
pOH = -log[OH-]
BOH
B+ + OH-
complete dissociation of a base
pH = 14 - pOH = 14 + log [OH-]
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pH of weak acids and bases
Dissociation of weak acids (Ka < 10-4)
c-x
HA + H2O
x
x
A- + H3O+
[A-][H3O+]
x2
x2
Ka =
=
=
[HA]
c-x
c
c-x = concentration of an acid at equilibrium
x = concentration of products at equilibrium
c = concentration of an acid at the beginning
pKa = -logKa
pH = -log[H3O+]
c >> x
for diluted weak acids
[H3O+] = x = (Ka c)1/2 / log
pH = -log [H3O+] = ½ [pKa – log(c)]
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pH of weak acids and bases
Dissociation of weak bases
c-x
x
BH+
B + H 2O
+
x
OH-
[BH+][OH-]
x2
Kb =
=
[B]
c-x
c-x = concentration of a base at equilibrium
x = concentration of products at equilibrium
c = concentration of a base at the beginning
pKb = -logKb
x2
=
c
c >> x for diluted weak bases
[OH-] = x = (Kb c)1/2 / log
pOH = -log[OH-]
pH = 14 - pOH
pH = 14 – pOH = 14 – ½ [pKb – log(c)]
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Salt hydrolysis
 When salts composed of ions of a strong electrolyte (acid
or base) and ions of a weak electrolyte are dissolved,
complete salt dissociation occurs because ions of a strong
electrolyte can exist only in ionized form
 Ions originating from a weak electrolyte react with water
producing their conjugated particle
 Examples: CH3COONa, KCN, NH4Cl, NH4NO3
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pH calculations of buffer solutions
Buffer consisting of a weak acid and its salt with a strong base
Or
a weak base and its salt with a strong acid
HA + H2O
A- + H3O+ Ka
Henderson – Hasselbalch equation
pH = pKa + log[A-]/[HA]
A- – conjugated base
B + H 2O
HA – weak acid
BH+ + OH-
pOH = pKb + log[BH+]/[B] B – weak base
BH+ - conjugated acid
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1. The solution with the lowest pH is
A. 1.0M HF
B. 1.0M HCN
C. 1.0M HCOOH
D. 1.0M CH3COOH
2. As the [H3O+] in a solution decreases, the [OH-]
A. increases and the pH increases.
B. increases and the pH decreases.
C. decreases and the pH increases.
D. decreases and the pH decreases.
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3. The value of pKw at 25°C is
A. 1.0 x 10-14
B. 1.0 x 10-7
C. 7.00
D. 14.00
4. What is the pOH of 0.1 M NaOH?
A. 1
B. 0.0032
C. 0.40
D. 13.60
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5. A 0.010M acid solution has a pH of 2.00. The acid could be
A. HNO3
B. H2SO3
C. HCOOH
D. CH3COOH
6. Which of the following salts dissolves to produce a basic
aqueous solution?
A. LiF
B. KClO4
C. NaHSO3
D. NH4NO3
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7. Which of the following gases is a contributor to
the formation of acid rain?
A. H2
B. O3
C. SO2
D. NH3
8. Which of the following describes the relationship
between [H3O+] and [OH-]?
A. [H3O+][OH-] = 14.00
B. [H3O+] + [OH-] = 14.00
C. [H3O+][OH-] = 1.0 x 10-14
D. [H3O+] + [OH-] = 1.0 x 10-14
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33. A solution of known concentration is the definition of a
A. buffer solution.
B. neutral solution.
C. standard solution.
D. saturated solution.
19. A 1.0 ´ 10-4 M solution has a pH of 10.00. The solute is a
A. weak acid.
B. weak base.
C. strong acid.
D. strong base.
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9. The ionization of water at room temperature is represented by
A. H2O = 2H+ + O2B. 2H2O = 2H2 + O2
C. 2H2O = H2 + 2OHD. 2H2O = H3O+ + OH-
10. Addition of HCl to water causes
A. both [H3O+] and [OH-] to increase.
B. both [H3O+] and [OH-] to decrease.
C. [H3O+] to increase and [OH-] to decrease.
D. [H3O+] to decrease and [OH-] to increase.
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11. Which of the following statements concerning Arrhenius acids and Arrhenius
bases is incorrect?
a)
In the pure state, Arrhenius acids are covalent compounds.
b)
In the pure state, Arrhenius bases are ionic compounds.
c)
Dissociation is the process by which Arrhenius acids produce H+ ions
in solution.
d)
Arrhenius bases are also called hydroxide bases.
12. According to the Bronsted-Lowry theory, a base is a(n)
a)
proton donor.
b)
proton acceptor.
c)
electron donor.
d)
electron acceptor.
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13. The pH of a solution for which [OH–] = 1.0 x 10–6 is
a)
1.00.
b)
8.00.
c)
6.00.
d)
–6.00.
the pH of 1.0 M acetic acid (Ka is 1.86 x10-5 at 20 °C.
a) 0.0043 b) 0.0034 c) 0.043 d) 0.034
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