Acids and Bases
Chapter 15
Acids in Industry
Sulfuric acid, H2SO4, is the chemical
manufactured in greatest quantity in the U.S.
Eighty billion pounds of sulfuric acid are used
each year to manufacture:
-fertilizers
-pharmaceuticals
-detergents
-storage batteries
-plastics
-metals
-petroleum
Properties of Acids
Acids:
•
taste sour (citrus fruits & vinegar)
•
affect indicators (e.g. turn blue litmus red)
•
produce H+ ions in aqueous solution
•
corrosive to metals
•
pH < 7
Classifying Acids
Organic acids contain a carboxyl group or
-COOH -- HC2H3O2 & citric acid.
Inorganic acids -- HCl, H2SO4, HNO3.
Oxyacids -- acid proton attached to oxygen
-- H3PO4.
Monoprotic -- HCl & HC2H3O2
Diprotic -- H2SO4
Triprotic -- H3PO4
Properties of Bases
Bases:
•
taste bitter
•
feel slippery
•
affect indicators (e.g. turn red litmus blue)
•
produce OH- ions in aqueous solution
•
pH > 7
•
caustic
Models of Acids and Bases
Arrhenius Concept: Acids produce H+ in
solution, bases produce OH ion.
Brønsted-Lowry: Acids are proton (H+)
donors, bases are proton (H+) acceptors.
HCl + H2O  Cl + H3O+
acid base
Bronsted-Lowry Model
The Bronsted-Lowry Model is not limited to
aqueous solutions like the Arrhenius
Model.
NH3(g) + HCl(g) ----> NH4Cl(s)
This is an acid-base reaction according to
Bronsted-Lowry, but not according to
Arrhenius!
Hydronium Ion
Hydronium (H3O+) ion is a hydrated proton -H+ . H2O.
The H+ ion is simply a proton. It has a very
high charge density, so it is strongly
attracted to the very electronegative
oxygen of the polar water molecule.
Conjugate Acid/Base Pairs
HA(aq) + H2O(l)  H3O+(aq) + A(aq)
conj
acid 1
conj
base 2
conj
acid 2
conj
base 1
conjugate base: everything that remains of
the acid molecule after a proton is lost.
conjugate acid: formed when the proton is
transferred to the base.
Which is the stronger base--H2O or A-?
Conjugate Acid-Base Pairs
Conjugate Acid Substance Conjugate Base
OHHOH
H3O+
NH4+
NH3
NH2-
H2SO4
HSO4-
SO42-
H3PO4
H2PO4-
HPO42-
H2PO41
-
HPO4
2-
PO43-
Conjugate Acid-Base Pairs
Which of the following represent conjugate
acid-base pairs?
a) HF, F-
e) OH-, HNO3
b) NH4+, NH3
f) H2O, H3O+
c) HCl, H2O
g) H2SO4, SO42-
d) HC2H3O2, C2H3O2-
h) HClO4, ClO4-
Conjugate Bases
Write the conjugate base for each of the
following:
a) HClO4
ClO4-
b) H3PO4
H2PO4-
c) CH3NH3+
CH3NH2
Acid Strength
Strong Acid:
-
Its equilibrium position lies far to the right.
(HNO3)
-
Yields a weak conjugate base. (NO3)
Acid Strength
(continued)
Weak Acid:
-
Its equilibrium lies far to the left.
(HC2H3O2)
-
Yields a much stronger (water is relatively
strong) conjugate base than water.
(C2H3O2-)
14_1577
H+
- H+
A
A H+
AH+
A(a)
AH+
H+
H+
A-
H+
HB
A-
HB
H+
HB
HB
H+
H+
H+
HB
HB
A-
AA-
A-
H+
HB
A-
HB
HB
B-
HB
HB
(b)
A strong acid is nearly 100 % ionized, while a weak acid
is only slightly ionized.
14_322
Before dissociation
HA
After dissociation,
at equilibrium
H+ A–
(a)
HA
HA
H+ A–
(b)
Diagram a represents a strong acid, while b represents a weak
acid which remains mostly in the molecular form.
14_323
Relative
acid strength
Very
strong
Relative
conjugate
base strength
Very
weak
Strong
Weak
Weak
Strong
Very
weak
Very
strong
The relationship of acid strength and conjugate base
strength for acid-base reactions.
Bases
Bases are often called alkalis because they
often contain alkali or alkaline earth metals.
“Strong” and “weak” are used in the same
sense for bases as for acids.
strong = complete dissociation (hydroxide ion
supplied to solution)
NaOH(s)  Na+(aq) + OH(aq)
Bases
(continued)
weak = very little dissociation (or reaction
with water)
NH3(aq) + HOH(l)  NH4+(aq) + OH(aq)
Water as an Acid and a Base
Water is amphoteric (it can behave either as an
acid or a base).
H2O + H2O <---> H3O+ + OH
acid 1 base 2
conj
acid 2
conj
base 1
Kw = 1  1014 M2 at 25°C
Kw = [H+][OH-]
Only about two molecules in a billion ionize!!
Ion product Constant, Kw
Kw is called the ion-product constant or
dissociation constant.
neutral solution [H+] = [OH-] = 1.0 x 10 -7 M
acidic solution [H+] > [OH-] [H+] > 1.0 x 10-7 M
basic solution [H+] < [OH-] [H+] < 1.0 x 10-7 M
No matter what the concentration of H+ or OH- in an
aqueous solution, the product, Kw, will remain
the same.
[H+] & [OH-] Calculations
Calculate the [H+] for a 1.0 x 10-5 M OH-.
Kw = [H+][OH-]
[H+] = Kw/[OH-]
[H+] = 1.0 x 10-14 M2/1.0 x 10-5 M
[H+] = 1.0 x 10-9 M
[H+] & [OH-] Calculations
Continued
Calculate the [OH-] for a 10.0 M H+.
Kw = [H+][OH-]
[OH-] = Kw/[H+]
[OH-] = 1.0 x 10-14 M2/10.0 M
[OH-] = 1.0 x 10-15 M
[H+] & [OH-] Calculations
Calculate the [H+] for a 2.0 x 10-2 M OH-.
Kw = [H+][OH-]
[H+] = Kw/[OH-]
[H+] = 1.0 x 10-14 M2/2.0 x 10-2 M
[H+] = 5.0 x 10-13 M
The pH Scale
pH = log[H+]
pH in water usually ranges from 0 to 14.
Kw = 1.0  1014 M2 = [H+] [OH]
pKw = 14.00 = pH + pOH
As pH rises, pOH falls (sum = 14.00).
Figure 15.5: Indicator
paper being used to
measure the pH of a
solution
Figure 15.4: A pH meter
pOH = 14
1x 10-14
pOH = 7
1 x 10-7
pOH = 0
1 x 100
OH -
OH-
+
O
H3
H3O+
OH
H3O+
1 x 100
pH = 0
1 x 10-7
pH = 7
1 x 10-14
pH = 14
Logarithms
-log 1.00 x 10-7 = 7.000
7.000
characteristic
mantissa
The number of significant digits in 1.00 x 10-7
is three, therefore, the log has three
decimal places. The mantissa represents
the log of 1.00 and the characteristic
represents the exponent 7.
14_324
[H+] pH
10–14 14
–13
10
Basic
13
10–12 12
–11
10
1 M NaOH
11
Ammonia
(Household
cleaner)
10–10 10
10–9
9
10–8
8
Neutral 10–7
7
10–6
6
10–5
5
10–4
4
10–3
3
10–2
2
10–1
1
1
0
Acidic
Blood
Pure water
Milk
Vinegar
Lemon juice
Stomach acid
1 M HCl
pH scale and pH values for common substances. A pH of
1 is 100 times more acidic than a pH of 3.
pH & Significant Figures
log
# Significant Figures -------> # decimal places
<------inv log
pH = - log [H+]
[H+] = inv log (-pH)
[H+] = 1.0 x 10-5 M
pH = 5.00
pH Calculations
Calculate the pH value for the following
solution at 25 oC.
[H+] = 1.0 x 10-9 M
pH = - log [H+]
pH = - log [1.0 x 10-9]
pH = 9.00
pH Calculations
Calculate the pH for the following solution at
25 oC.
[OH-] = 1.0 x 10-6M
pH + pOH = 14.00
pOH = - log [OH-]
pH = 14.00 - pOH
pOH = - log [1.0 x 10-6]
pH = 14.00 - 6.00
pOH = 6.00
pH = 8.00
pH Calculations
What is the pOH, [H+], & [OH-] for human
blood with a pH of 7.41?
pH + pOH = 14.00
pOH = 14.00 - pH
pOH = 14.00 - 7.41
pOH = 6.59
pH Calculations
Continued
What is the pOH, [H+], & [OH-]
for human blood with a pH of
7.41?
pH = - log [H+]
[H+] = antilog (-pH)
[H+] = antilog (-7.41)
[H+] = 3.9 x 10-8 M
Note: The number of
significant figures in the
antilog is equal to the number
of decimal places in the pH.
pH Calculations
Continued
What is the pOH, [H+], & [OH-]
for human blood with a pH of
7.41?
pOH = - log [OH-]
[OH-] = antilog (-pOH)
[OH-] = antilog (-6.59)
[OH-] = 2.6 x 10-7 M
Note: The number of
significant figures in the
antilog is equal to the number
of decimal places in the pOH.
pH of Strong Acid Solutions
Calculate the pH of a 0.10 M HNO3 solution.
Major species are: H+, NO3-, and H2O
Sources of H+ are from HNO3 and H2O -amount from water is insignificant.
[H+] = 0.10 M
Note: The number of
significant figures in
the [H+] is the same as
the decimal places in
the pH.
pH = - log [H+]
pH = - log [0.10]
pH = 1.00
A Buffered Solution
. . . resists change in its pH when either H+ or
OH are added.
1.0 L of 0.50 M HC2H3O2
+ 0.50 M Na C2H3O2
pH = 4.74
Adding 0.010 mol solid NaOH raises the pH
of the solution to 4.76, a very minor change.
Preparation of Buffered Solutions
Buffered solution can be made from:
1. a weak acid and its salt (e.g. HC2H3O2 &
NaC2H3O2).
2. a weak base and its salt (e.g. NH3 & NH4Cl).
Other examples of buffered pairs are:
H2CO3 & NaHCO3
H3PO4 & NaH2PO4
NaH2PO4 & Na2HPO4
Na2HPO4 & Na3PO4
Characteristics of a Buffer
1. The solution contains a weak acid HA and its
conjugate base A-.
2. The buffer resists changes in pH by reacting
with any added H+ or OH- so that these ions do
not accumulate.
3. Any added H+ reacts with the base A-.
4. Any added OH- reacts with the weak acid HA.
Buffered Solutions
Used when need to maintain a certain pH in the system.
 Blood
Buffers work by reacting with added H+1 or OH-1 ions
so they do not accumulate and change the pH.
Buffers will only work as long as there is sufficient
weak acid and conjugate base molecules present.
Buffering Mechanism
HC2H3O2(aq) <---> H+(aq) + C2H3O2-(aq)
The buffering materials dissolved in the
solution prevent added H+ or OH- from
building up in solution.
Buffering Capacity
. . . represents the amount of H+
or OH the buffer can absorb
without a significant change
in pH.