Chapter 14.
Acids and Bases
Early attributes of acids and bases (1600's)
Acids
Taste sour
Turn litmus red
React with metals
React with carbonates
Bases
Taste bitter
Turn litmus blue
Feel slippery
React with fats
Arrhenius Acids and Bases
Definitions of Svante Arrhenius, 1884
First working theory about acids and bases
Acids contain hydrogen and produce H1+
ions in water.
Bases contain hydroxide ions (OH1-) and
are soluble in water.
Acids and Bases
Arrhenius Acids and Bases
Acids are molecular compounds; a covalent
bond attaches the hydrogen ion to the adjacent atom.
Ionization, the separation of the molecule
into ions, occurs when the molecule dissolves in water.
Arrhenius Acids and Bases
Bases are ionic compounds; the hydroxide
ion exists in the crystal structure of the
solid compound.
Dissociation occurs when the ionic solid
dissolves in water, releasing the ions to
move about.
Arrhenius Acids and Bases
Common Acids:
HCl(aq), H2SO4, H3PO4, HNO3
HC2H3O2 = CH3COOH = acetic acid
Common Bases:
NaOH, KOH
Bronsted-Lowry
Acids and Bases
The Arrhenius definition has some problems:
It's restricted to water.
It doesn't explain why solutions of some
molecular compounds (NH3) and salts
(Na2CO3) are basic.
It doesn't explain why some salt solutions
are acidic (aqueous Al3+, Fe3+ solutions).
Bronsted-Lowry
Acids and Bases
Definitions of Brønsted and Lowry, 1923
widely used theory of acids and bases
Most
Acids are proton donors.
Bases are proton acceptors.
Reactions:
HCl(aq) + H2O(l)  H3O1+(aq) + Cl1-(aq)
H3O1+(aq) + NH3(aq)  NH41+(aq) + H2O(l)
Bronsted-Lowry
Acids and Bases
Bronsted-Lowry Acids and Bases
Formation of water by the transfer of protons
from H3O1+ ions to OH1 ions.
Bronsted-Lowry
Acids and Bases
Works in solvents other than water
Solves the base problem:
NH3(aq) + H3O1+(aq)  NH41+(aq) + H2O(l)
CO32-(aq) + H3O1+(aq)  HCO31-(aq) + H2O(l)
Doesn't solve the acid problem;
What is it with Al3+(aq) and Fe3+(aq)?
Lewis Acids and Bases
Definitions of Gilbert Lewis, 1923
Most general theory of acids and bases
Acids are electron pair acceptors.
Bases are electron pair donors.
Bronsted-Lowry
Acids and Bases
Conjugate acid-base pairs:
A reaction between and acid and a base
produces a conjugate acid and a conjugate base
HCl(aq) + H2O(l)  H3O1+(aq) + Cl1(aq)
Acid
Base
Conjugate
Acid
Conjugate
Base
H3O1+(aq) + NH3(aq)  H2O(l) + NH41+(aq)
Acid
Base
Conj.
Base
Conj.
Acid
Bronsted-Lowry
Acids and Bases
Choose the acid, base, conjugate acid, and
conjugate base:
HCOOH(aq) + NH3(aq)  HCOO1(aq) + NH41+ (aq)
H2PO41(aq) + H2O(l)  HPO42(aq) + H3O1+ (aq)
H2O(l) + HPO42(aq)  + H3O1+ (aq) + PO43(aq)
Bronsted-Lowry
Acids and Bases
Amphoteric substances can act as both
acids and bases:
HCOOH(aq) + H2O(l)  HCOO1(aq) + H3O1+ (aq)
NH3 (aq) + H2O(l)  NH41+ (aq) + OH1(aq)
HPO42(aq) + OH1(aq)  PO43(aq) + H2O(l)
HPO42(aq) + H3O1+ (aq)  H2PO42(aq) + H2O(l)
Mono-, Di- and Triprotic Acids
Monoprotic acids can transfer one proton
CH3COOH + H2O  CH3COO1 + H3O1+
Diprotic acids can transfer two protons
H2CO3 + H2O  HCO31 + H3O1+
HCO31 + H2O  CO32 + H3O1+
The first proton transfer is complete before the
second one starts.
Mono-, Di- and Triprotic Acids
Triprotic acids can transfer three protons
H3PO4 + H2O  H2PO41 + H3O1+
H2PO41 + H2O  HPO42 + H3O1+
HPO42 + H2O  PO43 + H3O1+
The first proton transfer is complete before the
second one starts. The second proton transfer
is complete before the third one starts.
Strengths of Acids and Bases
Acids differ in the extent of ionization when
they are put in solution
Strong acids ionize completely. There are
only a few strong acids.
Weak acids do not ionize completely. Most
acids are weak acids.
The equilibrium constant, Ka, is a measure of
the strength of an acid.
The Strong Acids
Formula
HCl(aq)
HBr(aq)
HI(aq)
HNO3
HClO4
HClO3
H2SO4
Name
Hydrochloric acid
Hydrobromic acid
Hydriodic acid
Nitric acid
Perchloric acid
Chloric acid
Sulfuric acid*
*first proton
Some Weak Acids
Formula
HSO41
C9H8O4
HCOOH
HC3H5O3
CH3COOH
H2CO3
H2S(aq)
HCN(aq)
C6H5OH
Name
Hydrogen sulfate
Acetylsalicylic acid
Formic acid
Lactic acid
Acetic acid
Carbonic acid
Hydrosulfuric acid
Hydrocyanic acid
Phenol
Ka
1.2 x 102
3.0 x 104
1.8 x 104
1.4 x 104
1.8 x 105
4.3 x 107
1.0 x 107
4.9 x 1010
1.3 x 1010
A comparison of the
number of H3O1+
ions present in
strong acid and
weak acid solutions of equal
concentration.
The Strong Bases
Soluble compounds that contain OH1Group 1A Hydroxides
LiOH
NaOH
KOH
RbOH
CsOH
Group 2A Hydroxides
Ca(OH)2
Sr(OH)2
Ba(OH)2
The Weak Bases
Ammonia:
NH3 + H2O  NH41+ + OH1
Kb = 1.8 x 105
NH41+ + H2O  NH3 + H3O1+
Ka = 5.6 x 1010
The Weak Bases
Anions from weak acids:
CH3COO1 + H2O  CH3COOH + OH1
CO32- + H2O  HCO31- + OH1
Salts
A salt is a compound containing a metal or
polyatomic cation, and a nonmetal or
polyatomic anion (except OH1).
NaCl, NH4Cl, BaSO4, CaCO3, Al2(SO4)3
Neutralization reactions between an acid
and a base produce a salt and water.
HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq)
2 Al(OH)3(s) + 3 H2SO4(aq) 
6 H2O + Al2(SO4)3 (aq)
Hydrolysis of Salts
Hydrolysis is a reaction of a substance with
water. Salts may hydrolyze to form H3O1+
or OH1 along with other products.
Hydrolysis of Salts
The salt of a weak acid and a strong base
gives a weakly basic aqueous solution.
NaOH + HC2H3O2  NaC2H3O2 + H2O
NaC2H3O2 + H2O  HC2H3O2 + OH1 + Na1+
Reestablishes equilibrium between acetate
anion and acetic acid.
Hydrolysis of Salts
The salt of a weak base and a strong acid
gives a weakly acidic aqueous solution.
NH3 + HCl  NH4Cl
NH4Cl + H2O  NH3 + H3O1+ + Cl1
Hydrolysis of Salts
The salt of a weak base and a weak acid can
give a weakly acidic, neutral, or weakly basic aqueous solution, depending on acid
strengths.
NH4C2H3O2 + H2O  HC2H3O2 + NH3
The salt of a strong acid and a strong base
give a neutral solution.
NaCl + H2O  Na1+ + Cl1 + H2O
Hydrolysis of Salts
Some metal ions, if they're small and have a
high charge, give acidic solutions.
Al3+ + 2 H2O  AlOH2+ + H3O1+
Keq = 1.4 x 105
Fe3+ + 2 H2O  FeOH2+ + H3O1+
Keq = 6.3 x 103
Cr3+ + 2 H2O  CrOH2+ + H3O1+
Keq = 1.6 x 104
Net Ionic Equations
2 Al(OH)3(s) + 3 H2SO4(aq) 
6 H2O(l) + Al2(SO4)3(aq)
2 HCl(aq) + CaCO3(s) 
2 CaCl2(aq) + CO2(g) + H2O(l)
Ionic Equations show dissolved ionic substances as ions rather than as compounds.
Net Ionic Equations show only the participating species. "Spectator" ions are not shown.
Net Ionic Equations
Ionic Equations
2 Al(OH)3(s) + 6 H1+(aq) + 3 SO42(aq) 
6 H2O(l) + 2 Al3+(aq) + 3 SO42(aq)
2 H1+(aq) + 2 Cl2(aq) + CaCO3(s) 
Ca2+(aq) + 2 Cl2(aq) + CO2(g) + H2O(l)
Net Ionic Equations (NIE's)
2 Al(OH)3(s) + 6 H1+  6 H2O(l) + 2 Al3+
2 H1+ + CaCO3(s)  Ca2+ + CO2(g) + H2O(l)
Self-Ionization of Water
The self-ionization of water is an acid-base
reaction in which one water molecule transfers a proton to another.
2 H2O  H3O1+ + OH1
Kw = 1.0 x 1014
Self-Ionization of Water
2 H2O  H3O1+ + OH1
Kw = 1.0 x 1014
Kw = ion product constant for water
Kw = 1.0 x 1014 = [H3O1+] [OH1]
In pure water, [H3O1+] = [OH1] = 1.0 x 107M
The relationship between [H3O1+] and [OH1] in
aqueous solution is an inverse proportion;
when [H3O1+] is increased, [OH1] decreases,
and vice versa.
Self-Ionization of Water
An acidic solution has [H3O1+] > 1.0 x 107 M
[OH1] < 1.0 x 107 M
A basic solution has [OH1] > 1.0 x 107 M
[H3O1+] < 1.0 x 107 M
A neutral solution has
[H3O1+] = [OH1] = 1.0 x 107 M
Self-Ionization of Water
Examples:
In a 0.015 M solution of HCl, what is the
concentration of OH1?
Is the solution acidic or basic?
[OH1] is 4.0 x 105. What is [H3O1+]?
Is the solution acidic or basic?
The pH Scale
[H3O1+] can vary over a wide range, and is
often low. Often, you need scientific
notation to express it. This isn't always
convenient.
A simpler way to write [H3O1+] is pH
pH = log [H3O1+]
[H3O1+] = 10pH
Common (base 10) Logarigthms
A logarithm is the power to which a base, such
as 10, must be raised to produce a given
number.
Number
Logarithm
0.010 = 1.0 x 10-2
1.0 = 1.0 x 100
10 = 1.0 x 101
-2.00
0.00
1.00
Coefficient Exponent
Characteristic Mantissa
Power of 10
1.E+06
9.E+05
8.E+05
7.E+05
6.E+05
5.E+05
4.E+05
3.E+05
2.E+05
1.E+05
0.E+00
Base 10 logarithm
Powers of 10 and their logarithms
8
7
6
5
4
3
2
1
0
Common (base 10) Logarigthms
What happens if the coefficient of the number
isn’t 1.0?
Number
0.050 = 5.0 x 10-2
5.0 = 5.0 x 100
50 = 5.0 x 101
Coefficient Exponent
Logarithm
-1.30
0.70
1.70
Characteristic Mantissa
Integers and their logarithms
1
0.8
0.7
0.6
0.5
0.4
0.3
0.2
0.1
Integer
10
9
8
7
6
5
4
3
2
1
0
0
Base 10 logarithm
0.9
Common (base 10) Logarigthms
How logarithms simplify mathematics:
2.594 x 103 x 6.022 x 1023 = 1.562 x 1027
log(2.594e3) + log(6.022 e23) = log(1.562e27)
3.4140 + 23.7797 = 27.1937
antilog(27.1937) = 1027.1937 = 1.562 x 1027
Common (base 10) Logarigthms
How logarithms simplify mathematics:
Slide rules use logarithmic scales for
multiplication and division.
The pH Scale
An acidic solution has [H3O1+] > 1.0 x 107 M
pH < 7.0
A basic solution has [OH1] > 1.0 x 107 M
pH > 7.0
A neutral solution has
[H3O1+] = [OH1] = 1.0 x 107 M
pH = 7.0
A pH meter is used to measure pH values. The
pH of vinegar is 2.32 (left). The pH of milk of
magnesia in water is 9.39 (right).
The pH Scale
Give the pH for
[H3O1+] =
=
=
=
=
0.010 M
4.2 x 103 M
1.0 x 107 M
6.8 x 1010 M
1.0 x 1012 M
Are the solutions acidic or basic?
The pH Scale
Give [H3O1+] for
pH =
=
=
=
=
3.00
4.50
6.85
7.00
10.75
Are the solutions acidic or basic?
Buffers
A buffer is a solution that resists major changes
in pH when acids or bases are added to it.
A buffer contains
A weak acid to react with added base
A weak base to react with added acid
Most often, the acid and base are conjugate pairs
Buffers
A buffer made of equimolar amounts of a
weak acid and its conjugate base will
have a pH equal to log Ka.
log Ka = pKa
Adding a acid will shift the pH of the buffer
down, adding base will shift the pH of
the buffer up.
Buffers
Examples:
What is the pH of a buffer made of 0.10
mole of CH3COOH and 0.10 mole of
NaCH3COO?
What is the pH of a buffer made with 1.0
mole of ammonia and 1.0 mole of
ammonium chloride?
Buffers
A buffer is made with 0.10 mole of CH3COOH
and 0.10 mole of NaCH3COO.
What component of the buffer reacts with
added H3O1+?
What component of the buffer reacts with
added OH1?
Buffers
A buffer is made with 1.0 mole of NH3 and
1.0 mole of NH4Cl.
What component of the buffer reacts
with added H3O1+?
What component of the buffer reacts
with added OH1?
Acid-Base Titrations
In an acid-base titration, a measured volume
of an acid or base of known concentration is
reacted with a measured volume of a base or
acid of unknown concentration.
The reaction is conducted in a way that exactly
equimolar amounts of H3O1+ and OH1 are
combined.
Acid-Base Titrations
A student titrates 2.00 mL of vinegar (acetic
acid in water) with 15.85 mL of 0.1048 M
NaOH. What is the concentration of acetic
acid in the vinegar? Calculate both molarity
and mass percent. The density of vinegar is
1.05 g/mL.
CH3COOH(aq) + NaOH(aq) 
CH3COONa(aq) + H2O(l)