Acids and Bases
Acids and Bases 2007-2008
1
Acids
• Svante Arrhenius, a Swedish chemist,
defines an acid as a substance that
yields hydrogen ions (H+) when dissolved
in water.
• Formulas for acids contain one or more
hydrogen atoms as well as an anion.
Acids and Bases 2007-2008
2
Naming Acids
In some cases two different names seem to be
assigned to the same chemical formula.
HCl(g)
HCl(l)
HCl(aq)
hydrogen chloride
hydrogen chloride
hydrochloric acid
The name assigned to the compound depends on
its physical state. In the gaseous or pure liquid
state, HCl is a molecular compound called
hydrogen chloride. When it is dissolved in
water, the molecules break apart into H+ and
Cl- ions; in this state, the substance is called
hydrochloric acid.
Acids and Bases 2007-2008
3
Binary acids (formed by hydrogen and
one other element) are named with a
“hydro-” prefix and an “-ic” ending on
the anion root.
Ex- HCl
HBr
Hydrochloric acid
Hydrobromic acid
Acids and Bases 2007-2008
4
The formulas for oxoacids, (acids that contain
hydrogen and an anion containing oxygen) are
usually written with the H first, followed by
the anion, as illustrated in the following
examples:
H2CO3
HNO3
HClO2
Carbonic acid
Nitric acid
Chlorous acid
If the anion ends in “-ate” then the acid ends in “-ic”,
if the anion ends in “-ite”, then the acid ends in “-ous”.
Remember, “ic” goes with the higher oxidation state, N
+5 in HNO3 (nitric acid) and
has an oxidation state of ___
+3 in HNO2 (nitrous acid)
___
Acids and Bases 2007-2008
5
Acids
• Acids have a sour taste; for example, vinegar
owes its sourness to acetic acid, and lemons
and citrus fruits contain citric acid.
• Acids cause color changes in plant dyes; for
example, they change the color of blue litmus
red.
• Acids react with certain metals to produce
hydrogen gas.
• Acids react with carbonates and bicarbonates
to produce carbon dioxide gas.
• Aqueous acid solutions conduct electricity.
Acids and Bases 2007-2008
6
Brønsted Acid
• Arrhenius’s definitions of acids are
limited in that they apply only to
aqueous solutions. Broader definitions
were proposed by the Danish chemist
Johannes Brønsted. A Brønsted acid is
a proton donor.
HCl(aq)
H+(aq) + ClRemember, the H+ ion is really just a proton (a
hydrogen atom is one proton and one electron,
you pull off the electron and all you are left
Acids and Bases 2007-2008
with is…
7
The size of a proton is about 10-15 m,
compared to the diameter of 10-10 m for
an average atom or ion. Such an
exceedingly small charged particle
cannot exist as a separate entity in
aqueous solution owing to its strong
attraction for the negative region of
the polar water molecule. Consequently,
the proton exists in a hydrated form as
H3O+, and is referred to as the
hydronium ion
H+ + H2O
H3O+
Acids and Bases 2007-2008
8
Since the acidic properties of the proton
are unaffected by hydration, we will
generally use H+(aq) to represent the
hydrated proton. This notation is for
convenience only, because H3O+ is closer
to reality. Keep in mind that both
notations represent the same species in
aqueous solution.
H+(aq) = H3O+(aq)
Acids and Bases 2007-2008
9
Monoprotic acids
HCl
each unit of acid yields one hydrogen
upon ionization
H2CO3
Diprotic acids
each unit of an acid gives up two H+ ions
Triprotic acids
H3PO4
yields three H+ ions upon ionization
Acids and Bases 2007-2008
10
Diprotic acids give up their two H+ ions in
separate steps:
H2SO4(aq)
H+(aq) + HSO4-(aq)
HSO4-(aq)
H+(aq) + SO4-2(aq)
Triprotic acids give up their H+ ions in
three separate steps.
Is HSO4- a strong or weak acid? Explain
Weak, only partially ionizes
Acids and Bases 2007-2008
11
Bases
• In another definition formulated by
Svante Arrhenius, a base can be
described as a substance that yields
hydroxide ions (OH-) when dissolved in
water. Some examples are
NaOH
Sodium hydroxide
KOH
Potassium hydroxide
Barium hydroxide
Ba(OH)2
Acids and Bases 2007-2008
12
Ammonia (NH3) is also classified as a
common base. At first glance this may
seem to be an exception to the
definition of a base. Note that as long
as a substance yields hydroxide ions
when dissolved in water, it need not
contain hydroxide ions in its structure
to be considered a base. In fact, when
ammonia dissolves in water, the
following reaction occurs:
Thus it is
properly
NH3 + H2O
NH4+ + OH- classified
Acids and Bases 2007-2008
as a base.13
Bases
• Bases have a bitter taste
• Bases feel slippery; for example, soaps,
which contain bases, exhibit this
property
• Bases cause color changes in plant dyes;
for example, they change the color of
red litmus blue
• Aqueous base solutions conduct
electricity
Acids and Bases 2007-2008
14
Brønsted Base
• A Brønsted base is defined by Johannes
Brønsted as being a substance capable
of accepting a proton.
Acids and Bases 2007-2008
15
Lewis Acids and Bases
So far we have discussed acid-base
properties in terms of the Brønsted
theory.
G.N. Lewis formulated a definition for
what is now called a Lewis base – a
substance that can donate a pair or
electrons. A Lewis acid is a substance
that can accept a pair of electrons.
Acids and Bases 2007-2008
16
The significance of the Lewis concept is that it
is much more general than other definitions.
For example, the reaction between boron
trifluoride and ammonia is a Lewis acid-base
reaction.
Acids and Bases 2007-2008
17
Strength of Acids and Bases
Strong acids are strong electrolytes,
which, for practical purposes, are
assumed to ionize completely in water.
That means that at equilibrium,
solutions of strong acids will not contain
any nonionized acid molecules.
Like strong acids, strong bases are all
strong electrolytes that ionize
completely in water.
Acids and Bases 2007-2008
18
The most common strong acids are HClO4,
HCl, HNO3 and H2SO4. Hydroxides of
alkali metals and alkaline Earth metals
are strong bases (like NaOH, KOH and
Ba(OH)2).
Other strong acids and strong bases are
listed on your Relative Strengths of
Acids and Bases Reference Sheet.
Acids and Bases 2007-2008
19
The strength of an acid is measured by
its tendency to ionize:
HX  H+ + XThe strength of the H-X bond influences
the extent to which an acid undergoes
ionization. The stronger the bond (the
higher the bond dissociation energy in
kJ/mol), the more difficult it is for the
HX molecule to break up and hence the
weaker the acid.
Acids and Bases 2007-2008
20
Bond Dissociation Energies for
Hydrogen Halides and Acid Strengths
Bond Bond Dissociation
Energy (kJ/mol)
H-F
568.2
Acid Strength
weak
H-Cl
431.9
strong
H-Br
366.1
strong
H-I
298.3
strong
Acids and Bases 2007-2008
21
The Strength of Oxoacids
Oxoacids contain hydrogen, oxygen,
and one other element Z, which
occupies a central position. To
compare oxoacid strength, it is
convenient to separate the oxoacids
into two groups.
Acids and Bases 2007-2008
22
Oxoacids having different central atoms that
are from the same group of the periodic
table and that have the same oxidation
number. Within this group, acid strength
increases with increasing electronegativity of
the central atom.
HClO3 > HBrO3
The Cl pulls more strongly on the electron
pair shared with the O, making the O-H
bond more polar, therefore making it
easier to ionize (and the acid stronger)
Acids and Bases 2007-2008
23
Oxoacids having the same central atom but
different numbers of attached groups.
Within this group, acid strength increases as
the oxidation number of the central atom
increases.
HClO4 > HClO3 > HClO2 > HClO
The greater the number of O atoms pulling on
the Cl, the more that the electrons are pulled
away from the O-H bond, making the O-H
bond more polar, therefore making it easier to
ionize (and the acid stronger)
Acids and Bases 2007-2008
24
You are going to remember these trends
because…
•HCl is a strong acid and HF (with the higher
bond dissociation energy) isn’t.
•HClO4 is a strong acid and HClO3 (where the
Cl has a lower oxidation number because it
has fewer oxygen atoms attached to it) isn’t.
Acids and Bases 2007-2008
25
Note: H3O+ is the strongest acid that can
exist in aqueous solutions. Acids
stronger than H3O+ react with water to
produce H3O+ and their conjugate bases.
Thus, HCl, which is a stronger acid than
H3O+, reacts with water completely to
form H3O+ and Cl-.
HCl(aq) + H2O(l)
H3O+(aq) + Cl-(aq)
Acids and Bases 2007-2008
26
The OH- ion is the strongest base that can
exist in aqueous solution. Bases stronger than
OH- react with water to produce OH- and
their conjugate acids.
For example, the oxide ion, (O-2) is a stronger
base than OH-, so it reacts with water
completely as follows:
O-2(aq) + H2O(l)  2OH-(aq)
For this reason, the oxide ion does not exist in
aqueous solutions.
Acids and Bases 2007-2008
27
Amphoteric Compounds
As you could see from the previous two examples,
water will act as either an acid or a base, depending on
the strength of the acid or base with which it is
reacting. Any species that can react as either an acid
or a base is described as amphoteric.
H2SO4(aq) + H2O(l)  H3O+(aq) + HSO4-(aq)
Proton acceptor
(base)
NH3(g) + H2O(l)  OH-(aq) + NH4+(aq)
Proton donor
(acid)
Acids and Bases 2007-2008
28
An extension of the Brønsted definition
of acids and bases is the concept of the
conjugate acid-base pair
CH3COOH(aq) + H2O(l)
Conjugate Conjugate
acid
base
Results from
the addition of
a proton to a
Bronsted base
Remains when
one proton has
been removed
from the acid
CH3COO-(aq) + H3O+(aq)
Acid
Proton (H+) donor
Base
Proton (H+) acceptor
A conjugate acid-base pair is defined as an acid and its conjugate base
(what’s left after the H+ was removed from the acid) or a base and its
conjugate acid (substance formed by the addition of the H+ to the base).
** Because the acid and base are always stronger than the conjugate acid and conjugate base, the
direction of the reaction proceeds from acid/base  conjugate acid/conjugate base.
29
Identify the acid, base, conjugate acid
and conjugate base in the following
reaction (**Reaction proceeds from stronger to weaker…)
NH3(aq) + H2O(l)
Conjugate
Base
What’s left
after H+ was
donated by
acid
(weaker base)
NH4+(aq) + OH-(aq)
Conjugate
Acid
Accepted
proton (H+)
(weaker acid)
Acid
Base
Proton
donor
Proton
(H+)
acceptor
(Stronger acid)
(stronger base)
Acids and Bases 2007-2008
30
The Acid-Base Properties of Water
Water is a very weak electrolyte and
therefore a poor conductor of
electricity, but it does undergo ionization
to a small extent:
H2O(l)
H+(aq) + OH-(aq)
This reaction is sometimes called the autoionization of water.
Acids and Bases 2007-2008
31
In the study of acid-base reactions in
aqueous solutions, the hydrogen ion
concentration is the key, because it
indicates the acidity or alkalinity of the
solution. Expressing the hydrogen ion as
H+, we can write the equilibrium constant
for the autoionization of water as
kw =
[H+][OH-]
[H2O]
Acids and Bases 2007-2008
Remember,
pure liquids and
solids are not
listed in the
ionization
equation,
therefore… 32
kw = [H+][OH-]
• kw is called the ion-product constant,
and is the product of the molar
concentrations of H+ and OH- ions at
a particular temperature.
Acids and Bases 2007-2008
33
In pure water at 25 oC, the
concentrations of H+ and OH- ions are
equal and found to be [H+] = 1.0 x 10-7 M
and [OH-] = 1.0 x 10-7 M. Thus,
kw = [H+][OH-]
kw = (1.0 x 10-7)(1.0 x 10-7)
kw = 1.0 x 10-14
Acids and Bases 2007-2008
34
Whether we have pure water or a solution
of dissolved species, the following
relation ALWAYS holds at 25 oC
kw = [H+][OH-] = 1.0 x 10-14
Acids and Bases 2007-2008
35
Because HCl is a strong acid…
HCl  H+ + ClCalculate the concentration of OH- ions in an HCl solution whose
hydrogen ion concentration is 1.3 M.
[HCl]
[H+]
[Cl-]
1.3
0.0
0.0
C (change)
-1.3
+1.3
+1.3
E (end)
0.0
1.3
1.3
I (initial)
kw = [H+][OH-]
1.0 x 10-14 = (1.3)[OH-]
[OH-] = 7.7 x 10-15 M
36
pH – A Measure of Acidity
Because the concentrations of H+ and
OH- ions in aqueous solutions are
frequently very small numbers and
therefore inconvenient to work with,
Soren Sorensen in 1909 proposed a more
practical measure called pH. The pH of a
solution is defined as the negative
logarithm of the hydrogen ion
concentration (in mol/L)
pH is a
pH = -log
[H+]
Acids and Bases 2007-2008
dimensionless
quantity (it will not
have a label)
37
Since pH is simply a way to express hydrogen ion
concentration, acidic and basic solutions at 25 oC can
be distinguished by their pH values, as follows:
Acidic solutions:
[H+] > 1.0 x 10-7 M, pH < 7.00
Basic solutions:
[H+] < 1.0 x 10-7 M, pH > 7.00
Neutral solutions
[H+] = 1.0 x 10-7 M, pH = 7.00
In an acidic solution there is an excess of H+ ions; [H+] > [OH-]
In a basic solution there is an excess of OH- ions; [OH-] > [H+]
Whenever [H+] = [OH-], the aqueous solution is said to be neutral.
**Note – when concentration has two significant figures, pH
will have two numbers TO THE RIGHT OF THE DECIMAL!
Acids and Bases 2007-2008
38
Calculate the pH of a 1.0 x 10-3 M HCl solution.
Since HCl is a strong acid, it is completely ionized
in solution:
Remember,
these are
HCl(aq)  H+(aq) + Cl-(aq)
molarities
I
C
E
HCl(aq)
H+(aq)
Cl-(aq)
1.0 x 10-3
0.0
0.0
-1.0 x 10-3 +1.0 x 10-3
0.0
1.0 x 10-3
+1.0 x 10-3
1.0 x 10-3
Thus, [H+] = 1.0 x 10-3 M
pH = -log(1.0 x 10-3)
pH = 3.00
Acids and Bases 2007-2008
39
The concentration of H+ ions in a bottle of table
wine was 3.2 x 10-4 M right after the cork was
removed. Only half of the wine was consumed.
The other half, after it had been standing open
to the air for a month, was found to have a
hydrogen ion concentration equal to 1.0 x 10-3 M.
Calculate the pH of the wine on these two
occasions.
When the wine was first opened
pH = -log [H+]
pH = -log (3.2 x 10-4) = 3.49
After the wine sat for a month
pH = -log [H+]
-3) = 3.00
pH
=
-log
(1.0
x
10
Why did the acidity increase?
Some of the ethanol converted to acetic acid, a reaction that takes place in
Acids and Bases 2007-2008
40
the presence of O2.
Given the pH of a solution, you can figure
out the [H+] concentration by using the
simple formula
[H+] = 10-pH
What is the hydrogen ion concentration of an acid
with a pH of 3.00?
[H+] = 10-pH
[H+] = 10-3.00
[H+] = 1.0 x 10-3
Acids and Bases 2007-2008
41
A pH meter is commonly used in the
laboratory to determine the pH of a
solution. Although many pH meters have
scales marked with values from 1 to 14, pH
values can, in fact, be less than 1 and
greater than 14.
Acids and Bases 2007-2008
42
A pOH scale analogous to the pH scale can
be devised using the negative logarithm of
the hydroxide ion concentration of a
solution. Thus we define pOH as
pOH = -log[OH-]
Acids and Bases 2007-2008
43
Now consider again the ion-product constant
for water:
kw = [H+][OH-] = 1.0 x 10-14
“Logs make
adders
multiply”
Taking the negative logarithm of both sides
we obtain
-log[H+] + -log[OH-] = -log(1.0 x 10-14)
-log[H+] + -log[OH-] = 14.00
pH
pOH
pH + pOH = 14.00
Acids and Bases 2007-2008
44
In a NaOH solution [OH-] is 2.9 x 10-4 M.
Calculate the pH of the solution.
First, figure out the pOH…
pOH = -log [OH-]
pOH = -log (2.9 x 10-4)
pOH = 3.54
Then use the pOH to figure out the pH…
pH + pOH = 14.00
pH = 14.00 – pOH
pH = 14.00 – 3.54 = 10.46
Acids and Bases 2007-2008
45
Calculate the pH of a 0.020 M Ba(OH)2 solution.
Ba(OH)2 is a strong base; each Ba(OH)2 unit
produces two OH-ions:
Ba(OH)2(aq)  Ba+2(aq) + 2OH-(aq)
I (M)
C (M)
E (M)
Ba(OH)2(aq) Ba+2(aq)
OH-(aq)
0.020
0.00
0.00
-0.020
+0.020
2(+0.020)
0.00
0.020
0.040
Thus, [OH-] = 0.040 M
pOH = -log 0.040 = 1.40
pH = 14.00 – pOH
pH = 14.00 – 1.40 = 12.60
Acids and Bases 2007-2008
46
To determine the hydroxide ion when given the
pOH, you need to use the formula
10-pOH = [OH-]
What is the molarity of a NaOH solution that has a
pH of 11.30?
10-11.30
10-pH = [H+]
= 5.0 x 10-11 = [H+]
[H+][OH-] = kw
(5.0 x 10-11)[OH-] = 1.0 x 10-14
[OH-] = 2.0 x 10-4
Because NaOH is a strong base, the [OH-] at the
end is equal to the initial concentration of NaOH.
Acidsthe
and Bases
2007-2008
The molarity of
NaOH
= 2.0 x 10-4 M
47
Weak Acids and Acid
Ionization Constants
Most acids are weak acids, which ionize
only to a limited extent in water. At
equilibrium, aqueous solutions of weak
acids contain a mixture of nonionized
acid molecules, H3O+ ions, and the
conjugate base.
The limited ionization of weak acids is
related to the equilibrium constant for
ionization, which is represented as Ka.
Acids and Bases 2007-2008
48
Consider a weak monoprotic acid, HA. Its
ionization in water is represented by
HA(aq) + H2O(l)
H3O+(aq) + A-(aq)
or simply
HA(aq)
H+(aq) + A-(aq)
Acids and Bases 2007-2008
49
Write the equilibrium expression for the
ionization of HA.
Ka=
[H+][A-]
[HA]
Ka, the acid ionization constant, is
the equilibrium constant for the
ionization of an acid.
Acids and Bases 2007-2008
50
At a given temperature, the strength of
the acid HA is measured quantitatively
by the magnitude of Ka. The larger Ka,
the stronger the acid – that is, the
greater the concentration of H+ ions at
equilibrium due to its ionization. Keep in
mind, however, that only weak acids
have Ka values associated with them.
Acids and Bases 2007-2008
51
You have a reference sheet that lists a
number of weak acids and their Ka
values at 25 oC. Although all of the
acids on that sheet are weak, within the
group there is great variation in their
strengths.
For example, Ka for HF (6.8 x 10-4) is
about 1.5 million times greater than that
for HCN (6.2 x 10-10).
Acids and Bases 2007-2008
52
Generally, we can calculate the hydrogen
ion concentration or pH of an acid
solution at equilibrium, given the initial
concentration of the acid and its Ka
value.
Alternatively, if we know the pH of a
weak acid solution and its initial
concentration, we can determine its Ka.
Acids and Bases 2007-2008
53
Suppose we are asked to calculate the pH
of a 0.50 M HF solution at 25 oC. The
ionization of HF is given by
HF(aq)
H+(aq) + F-(aq)
From your reference sheet we can
write
[H+][F-]
Ka =
= 6.8 x 10-4
[HF]
Acids and Bases 2007-2008
54
The first step is to identify all the
species present in solution that may
affect its pH. Because weak acids
ionize to a small extent, at equilibrium
the major species present are
nonionized HF and some H+ and F- ions.
Another major species is H2O, but its
very small Kw (1.0 x 10-14) means that
water is not a significant contributor to
the H+ ion concentration. Therefore,
unless otherwise stated, we will always
ignore the H+ or OH- ions produced by
the autoionization of water.
Acids and Bases 2007-2008
55
HF(aq)
H+(aq) + F-(aq)
We can summarize the changes in the concentrations
of HF, H+, and F- in the table below:
HF(aq) H+(aq)
I (M)
0.50
0.00
C (M)
-x
+x
E (M) 0.50 – x
x
Acids and Bases 2007-2008
F-(aq)
0.00
+x
x
56
The equilibrium concentrations of HF, H+
and F-, expressed in terms of the
unknown x, are substituted into the
ionization constant expression to give
Ka =
(x)(x)
0.50 - x
= 6.8 x 10-4
Rearranging this expression, we write
x2 + 6.8 x 10-4x – 3.4 x 10-4 = 0
Acids and Bases 2007-2008
57
This is a quadratic equation which can be
solved using the quadratic formula. Or,
we can try using a shortcut to solve for x.
Because HF is a weak acid and weak acids
ionize only to a slight extent, we reason
that x must be small compared to 0.50.
Therefore we can make the
approximation
0.50 – x ≈ 0.50
Acids and Bases 2007-2008
58
Now the ionization constant expression
becomes
x2
0.50 - x
≈
x2
0.50
= 6.8 x 10-4
Rearranging, we get
x2 = (0.50)(6.8 x 10-4) = 3.4 x 10-4
x = √3.4 x 10-4 = 0.018 M
Acids and Bases 2007-2008
59
Thus we have solved for x without having
to use the quadratic equation. At
equilibrium, we have:
[HF] = (0.50 – 0.018) M = 0.48 M
[H+] = 0.018 M
This is determined by going
[F-] = 0.018 M
back to the ICE chart
And the pH of the solution is
pH = -log(0.018) = 1.74
Acids and Bases 2007-2008
60
How good is this approximation? The
approximation is valid if the following
expression is equal to or less than 5%
Molarity of H+ at
equilibrium
0.018 M
Initial concentration
of weak acid
0.50 M
X 100 = 3.6%
If this is greater
than 5%, you must
use the quadratic
formula
61
The Quadratic Equation
-b± √b2 – 4ac
x=
2a
The values from the equation shown below,
(from slide 43), can now be substituted in
to the quadratic equation.
x2 + 6.8 x 10-4x – 3.4 x 10-4 = 0
a = 1; b = 6.8 x 10-4; c = -3.4 x 10-4
Acids and Bases 2007-2008
62
-b± √b2 – 4ac
x=
2a
-4 ± √(6.8 x 10-4)2 – 4(1)(-3.4 x 10-4)
-6.8
x
10
x=
2(1)
-4 ± .0014
-6.8
x
10
x=
2(1)
x = .018 M or -.018 M
Note – this
is the same
value as we
estimated
earlier!
The second solution is physically impossible because the concentration
of ions produced as a result of ionization cannot be negative.
pH = -log(0.018) = 1.74
63
Percent Ionization
We have seen that the magnitude of Ka
indicates the strength of an acid.
Another measure of the strength of an
acid is its percent ionization, which is
defined as
Percent ionization =
H+ concentration at equilibrium
Initial concentration of acid
Acids and Bases 2007-2008
X 100
64
The stronger the acid, the greater the
percent ionization.
The extent to which a weak acid ionizes
depends on the initial concentration of
the acid. The more dilute the solution,
the greater the percent ionization.
Acids and Bases 2007-2008
65
Diprotic and Polyprotic Acids
The treatment of diprotic and polyprotic
acids is more involved than that of
monoprotic acids because these
substances yield more than one
hydrogen atom per molecule.
These acids ionize in a stepwise manner,
that is, they lose one proton at a time.
An ionization constant expression should
be written for each ionization step.
Acids and Bases 2007-2008
66
Oxalic acid (H2C2O4) is a poisonous
substance used chiefly as a bleaching
and cleansing agent (for example, to
remove bathtub rings). Calculate the
concentrations of all the species
present at equilibrium in a 0.10 M
solution.
Acids and Bases 2007-2008
67
H2C2O4
H+ + HC2O4-
H2C2O4(aq)
H+(aq)
HC2O4-(aq)
I (M)
0.10
0.00
0.00
C (M)
-x
+x
+x
E (M)
0.10 – x
x
x
Acids and Bases 2007-2008
68
Ka =
Ka =
[H+][HC2O4-]
[H2C2O4]
(x)(x)
0.10 - x
= 5.6 x 10-2
= 5.6 x 10-2
Let me save you some work, you need to use the quadratic
formula for this one
x2 + 5.6 x 10-2x - 5.6 x 10-3 = 0
x = 0.052
Acids and Bases 2007-2008
69
When the equilibrium for the first stage
of ionization is reached, the
concentrations are:
[H+] =
0.052 M
[HC2O4-] = 0.052 M
[H2C2O4] = (0.10 - 0.052) M = 0.048 M
Acids and Bases 2007-2008
70
Next we consider the second stage of
ionization.
At this stage, the major species will be
HC2O4-, (this serves as the acid in the
second stage), H+, and C2O4-2 (the
conjugate base).
Acids and Bases 2007-2008
71
HC2O4-
H+ + C2O4-2
HC2O4-(aq)
H+(aq)
C2O4-2(aq)
I (M)
0.052
0.052
0.00
C (M)
-y
+y
+y
E (M)
0.052 – y
0.052 + y
y
Acids and Bases 2007-2008
72
Ka =
Ka =
[H+][C2O4-2]
[HC2O4-]
(0.052 + y)(y)
0.052 - y
= 5.4 x 10-5
= 5.4 x 10-5
Let me save you some work, you DON’T need to use the quadratic
formula for this one, 0.052 + y and 0.052 – y ≈ 0.052
Ka =
(0.052)(y)
0.052
= 5.4 x 10-5
y = 5.4 x 10-5
Acids and Bases 2007-2008
73
Testing the approximation - checking the
5% rule
5.4 x 10-5 M
0.052 M
X 100 =.10%
The approximation is valid.
Acids and Bases 2007-2008
74
At equilibrium
[H2C2O4] = 0.048 M
[HC2O4-] = (0.052 – 5.4 x 10-5) M = 0.052 M
[H+] = (0.052 + 5.4 x 10-5) M = 0.052 M
[C2O4-2] = 5.4 x 10-5 M
Acids and Bases 2007-2008
75
This example shows that for diprotic
acids, if Ka1 » Ka2, then we can assume
that the concentration of H+ ions is the
product of only the first stage of
ionization.
Acids and Bases 2007-2008
76
Weak Bases and Base Ionization
Constants
Weak bases, like weak acids, are weak
electrolytes.
Ammonia is a weak base that ionizes only
to a limited extent in water:
NH3(aq) + H2O(l)
NH4+ + OH-(aq)
Acids and Bases 2007-2008
77
The equilibrium constant is given by
Kb =
[NH4+][OH-]
[NH3]
Where Kb is called the base ionization
constant.
Acids and Bases 2007-2008
78
Follow the same procedures you used with
weak acids when solving problems
involving weak bases.
The main difference is that you will
calculate [OH-] first, rather than [H+].
Acids and Bases 2007-2008
79
The Relationship Between the Ionization
Constants of Acids and Their Conjugate
Bases
For any conjugate acid-base pair it is
always true that
KaKb = Kw
Acids and Bases 2007-2008
80
Calculate the Kb of the conjugate base of
acetic acid
CH3COOH
H+ + CH3COO-
Kw
Kb =
Ka
1.0 x 10-14
Kb =
1.8 x 10-5
Conjugate
base
Kb = 5.6 x 10-10
Acids and Bases 2007-2008
81
Acid – Base Properties of Salts
Salts (which are one of the products of
an acid-base neutralization reaction) are
strong electrolytes that completely
dissociate into ions in water. The term
salt hydrolysis describes the reaction
of an anion or a cation of a salt, or both,
with water. Salt hydrolysis usually
affects the pH of a solution.
Acids and Bases 2007-2008
82
Salts that Produce Neutral
Solutions
NaNO3
The cation
of this salt
came from
a strong
base
(NaOH)
H2O
Na+ + NO3-
The anion of this
salt came from a
strong acid
(HNO3)
Acids and Bases 2007-2008
83
Salts that Produce Neutral
Solutions
NaNO3
H2O
The Na+ ion and the
OH- from the water
would not stay
together (NaOH
would be a strong
base and would
dissociate)
Na+ + NO3The NO3- ion and the
H+ from the water
would not stay
together, (remember,
HNO3 is a strong
acid)
Acids and Bases 2007-2008
84
Salts that Produce Neutral
Solutions
NaNO3
H2O
Na+ + NO3-
Consequently, NaNO3 and other salts
formed from a strong acid and a
strong base do not affect the pH of a
solution.
Acids and Bases 2007-2008
85
Salts that Produce Basic
Solutions
NaCH3COO
The cation
of this salt
came from
a strong
base
(NaOH)
H2O
Na+ + CH3COO-
The anion of this
salt came from a
weak acid
(CH3COOH)
Acids and Bases 2007-2008
86
Salts that Produce Basic
Solutions
NaCH3COO
H2O
The Na+ ion and the
OH- from the water
would not stay
together (NaOH
would be a strong
base and would
dissociate)
Na+ + CH3COOThe CH3COO- ion and
the H+ from the
water WOULD stay
together, (remember,
CH3COOH is a weak
acid, meaning it only
ionizes slightly)
Acids and Bases 2007-2008
87
Salts that Produce Basic
Solutions
CH3
COO-
H2O
OH- + CH3COOH
Because the CH3COO- ions would bond
with the H+ ions, the OH- ions (which are
left behind when the H+ ions come off
of the water molecules) affect the pH
of the solution. In other words,
solutions produced by salts made from
strong bases and weak acids will be basic
in nature.
Acids and Bases 2007-2008
88
Salts that Produce Acidic
Solutions
NH4Cl
The cation
of this salt
came from
a weak
base
H2O
NH4+ + Cl-
The anion of this
salt came from a
strong acid
Acids and Bases 2007-2008
89
Salts that Produce Acidic
Solutions
NH4Cl
H2O
NH4+ + ClNH3 + H+
Because the Cl- ions wouldn’t bond with the
H+ ions, and the H+ ion that would separate
from the NH4+, it would leave excess H+ in
solution. In other words, solutions
produced by salts made from strong acids
and weak bases will be acidic in nature.
Acids and Bases 2007-2008
90
When the cation of a salt comes from a weak base and
the anion comes from a weak acid, you need to
compare the Ka and Kb values to determine if the
solution will be acidic or basic. For example,
NH4NO2
NH4+ + NO2-
The NH4+
dissociates
producing H+ ions
Ka = 5.7 x 10-10
NH4+
NH3 + H+
Kb = 1.4 x 10-11
NO2-
HNO2 + OH-
If Kb > Ka, the solution will
be basic,
if Ka > Kb, the solution will
be acidic..
therefore an aqueous
solution of NH4NO2
will be acidic.
Acids and Bases 2007-2008
NO2- has an affinity
for the H+ produced by
the autoionization of
water, leaving an
excess of OH- behind.
91
Oxides can be classified as acidic, basic, or
amphoteric.
• All alkali metal oxides and all alkaline earth
metal oxides except BeO are basic.
Na2O + H2O  2NaOH
BaO + H2O  2Ba(OH)2
• Beryllium oxide and several metallic oxides in
the boron family (Group 3A) and carbon
family (Group 4A) are amphoteric.
Acids and Bases 2007-2008
92
• Nonmetalic oxides in which the oxidation
number of the representative element is high
are acidic. The nonmetalic oxides that react
with water to form acids are sometimes
referred to as acidic anhydrides.
Representative elements in which the
oxidation number is low (for example, CO and
NO) show no measurable acidic properties.
CO2
SO3
N2O5
P4O10
Cl2O7
+
+
+
+
+
H2O 
H2O 
H2O 
H2O 
H2O 
H2CO3
H2SO4
2HNO3
4H3PO4
2HClO4
Acids and Bases 2007-2008
93
Buffer Solutions
A buffer solution is a solution of (1) a weak acid or a
weak base and (2) its salt; both components must be
present. The solution has the ability to resist
changes in pH upon the addition of small amounts of
either acid or base.
Buffers are very important to chemical and biological
systems. The pH in the human body varies greatly
from one fluid to another; for example, the pH of
blood is about 7.4, whereas the gastric juice in our
stomach has a pH of about 1.5. These pH values,
which are crucial for proper enzyme function and the
balance of osmotic pressure, are maintained by
buffers in most cases.
Acids and Bases 2007-2008
94
CH3COOH
NaCH3COO
CH3COO- + H+
H2O
CH3COO- + Na+
If you add a base to this
solution, the OH- will be
neutralized by the acetic
acid in the buffer,
therefore you will not
notice a significant
difference in the pH of
the solution.
Acids and Bases 2007-2008
95
CH3COOH
NaCH3COO
CH3COO- + H+
H2O
CH3COO- + Na+
If you add an acid to this solution, the
acetate ion will bond to the H+, so no
appreciable change in pH will be
observed because it is such a small
increase in H+
Acids and Bases 2007-2008
96
The buffering capacity, that is, the
effectiveness of the buffer solution,
depends on the amount of acid and
conjugate base from which the buffer is
made. The larger the amount, the
greater the buffering capacity.
Acids and Bases 2007-2008
97
In general, a buffer system can be
represented as salt-acid or conjugate
base-acid. Thus the sodium acetateacetic acid buffer system we discussed
can be written as
CH3COONa/CH3COOH
or simply
CH3COO-/CH3COOH.
Acids and Bases 2007-2008
98
Which of the following are buffer
systems?
NaClO4/HClO4
KF/HF
Acids and Bases 2007-2008
99
(a) Calculate the pH of a buffer system
containing 1.0 M CH3COOH and 1.0 M
CH3COONa.
(b) What is the pH of the buffer system
after the addition of 0.10 mole of
gaseous HCl to 1 L of the solution.
Assume that the volume of the
solution does not change when the HCl
is added.
Acids and Bases 2007-2008
100
CH3COOH
CH3COO-
H+
I
1.0
1.0
0.0
C
-x
+x
x
E
1.0 - x
1.0 + x
0.0 + x
Assume
ionization
is
negligible
Acids and Bases 2007-2008
101
CH3COOH
Ka =
CH3COO- + H+
[H+][CH3COO-]
[CH3COOH]
= 1.8 x 10-5
Rewriting the above equation…
[H+] = ka
[CH3COOH]
[CH3COO-]
[H+] = 1.8 x 10-5
Acids and Bases 2007-2008
(1.0)
(1.0)
102
[H+] = 1.8 x 10-5 M
-log[H+] = -log(1.8 x 10-5)
pH =
4.74
When the concentration of the acid and the
conjugate base are the same, the pH of the
buffer is equal to the pKa, which is determined
by taking the –log of the ka, of the acid
pH = pka = -log ka
Acids and Bases 2007-2008
103
(b) After the addition of HCl, complete
ionization of the 1.0 M HCl acid occurs;
HCl  H+ + Cl0.10 mol
0.10 mol
This 0.10 mol of H+ ions will bond with
0.10 mol of the CH3COO- ions,
DECREASING the amount of CH3COOby .10 mol and INCREASING the
amount of CH3COOH by .10 mol
Acids and Bases 2007-2008
104
So now instead of 1.0 mol of CH3COO- ions,
there will only be 0.90 mol (1.0 - .10)
And since there were originally 1.0 mol of
CH3COOH, you now have 1.10 mol of
CH3COOH with the added .10 mol
[CH3COO-] = .90 M
[CH3COOH] = 1.10 M
Acids and Bases 2007-2008
105
Ka =
[H+][CH3COO-]
[CH3COOH]
= 1.8 x 10-5
Rewriting the above equation…
[H+] =
Ka[CH3COOH]
[CH3COO-]
(1.8 x 10-5)(1.10)
[H+] =
(0.90)
Acids and Bases 2007-2008
106
[H+] = 2.2 x 10-5 M
pH =
-log(2.2 x 10-5)
pH =
4.66
2.2 x 10-5
1.8 x 10-5
= 1.2
This is a very slight change in pH, before
the HCl was added, the [H+] = 1.8 x 10-5,
after the addition of HCl, the [H+] = 2.2
x 10-5, this is an increase by a factor of
1.2 with a pH change from 4.74 to 4.66
Acids and Bases 2007-2008
107
Before the addition of HCl: [H+] = 1.0 x 10-7
After the addition of HCl: [H+] = 0.10 M
0.10 M
1.0 x 10-7 M
= 1.0 x 106
This would be a millionfold increase as
the pH changed from 7.00 to 1.00!
Acids and Bases 2007-2008
109
A "very convenient" equation for dealing
with buffer solutions is the HendersonHasselback equation.
[A-]
pH = pKa + log
[HA]
Acids and Bases 2007-2008
110
Previous question - Calculate the pH of a buffer
system containing 1.0 M CH3COOH and 1.0 M
CH3COONa.
[A-]
pH = pka + log
[HA]
[1.0]
pH = pka + log
[1.0]
log 1 = 0
pH = pka
pH = -log(ka)
pH = -log(1.8 x 10-5)
pH = 4.74
Acids and Bases 2007-2008
111
What is the pH of the previous buffer system after the
addition of 0.10 mole of gaseous HCl to 1 L of the
solution. Assume that the volume of the solution does
not change when the HCl is added.
[A-]
pH = pka + log
[HA]
[.90]
pH = pka + log
[1.10]
pH = -log(1.8 x 10-5) + -.041
pH = 4.66
Acids and Bases 2007-2008
Remember, the
.10 M H+ from
the HCl would
shift the
equilibrium to
the left taking
.1M CB,
decreasing the
amount of CB
to .90, and
increasing the
amount of A to
1.10
112
Acid – Base Titrations
Quantitative studies of acid-base
neutralization reactions are most
conveniently carried out using a
technique known as titration. In
titration, a solution of accurately known
concentration, called a standard solution,
is added gradually to another solution of
unknown concentration, until the chemical
reaction between the two solutions is
complete.
Acids and Bases 2007-2008
113
If we know the volumes of the standard
solution, we can calculate the
concentration of the unknown solution.
Sodium hydroxide is one of the bases
commonly used in the laboratory.
However, it is difficult to obtain solid
sodium hydroxide in a pure form because
it is hygroscopic, (it has a tendency to
absorb water from air), and its solution
reacts with carbon dioxide. For these
reasons, a solution of sodium hydroxide
must be standardized before it can be
used in accurate analytical work.
Acids and Bases 2007-2008
114
We can standardize the sodium hydroxide
solution by titrating it against an acid
solution of accurately known
concentration. The acid often chosen for
this task is a monoprotic acid called
potassium hydrogen phthalate
(abbreviated as KHP), for which the
molecular formula is KHC8H4O4.
Acids and Bases 2007-2008
115
The procedure for the titration of KHP and
NaOH is as follows:
1. Add a known amount of KHP to an
Erlenmeyer flask. Add some distilled
water to make up a solution.
2. Next, carefully add NaOH solution from
a buret until the equivalence point is
reached. The equivalence point is the
point at which the acid has completely
reacted with or been neutralized by the
base.
Acids and Bases 2007-2008
116
The equivalence point is usually signaled
by a sharp change in the color of an
indicator. In an acid-base titration,
indicators are substances that have
distinctly different colors in acidic and
basic solutions. One common indicator
is phenolphthalein.
Phenolphthalein indicates the presence
base
of a/n _________.
Phenolphthalein is
________
colorless in acidic solutions, _______
colorless
pink
in neutral solutions and ________
in
basic solutions.Acids and Bases 2007-2008
117
At the equivalence point, all the KHP
present has been neutralized by the
added NaOH and the solution is still
colorless. However, if we add just one
more drop of NaOH solution from the
buret, the solution will immediately turn
pink because the solution is now basic.
Acids and Bases 2007-2008
118
Apparatus for acid-base titration.
A NaOH solution is
added from the buret to
a KHP solution in an
Erlenmeyer flask.
A faint pink color
appears when the
equivalence point is
reached. If your
solution turns fuchsia,
you have gone past the
equivalence point
119
The neutralization reaction between NaOH and
KHP is one of the simplest types of acid-base
neutralization known. A neutralization
reaction is a reaction between an acid and a
base.
Aqueous strong acid-strong base reactions
produce water and a salt, (an ionic compound
made up of a cation other than H+ and an anion
other than OH- or O-2)
Acids and Bases 2007-2008
120
The reaction between KHP and sodium
hydroxide is:
KHC8H4O4(aq) + NaOH(aq)  KNaC8H4O4(aq) + H2O(l)
acid
base
Acids and Bases 2007-2008
salt
water
121
You can use the following format to solve
acid-base neutralization problems
Acid
M
mol
Base
L
M
Acids and Bases 2007-2008
mol
L
122
What is the molarity of the acid if 16.1 mL
of 0.610 M NaOH was required to
neutralize 20.0 mL of H2SO4?
Acid
M
mol
Base
L
M
.0200
.610
Acids and Bases 2007-2008
mol
L
.0161
123
Acid
M
mol
Base
L
M
.0200
mol
.610 .00982
L
.0161
.610 M = .610 mol =
1L
x
.0161 L
X = .00982
Acids and Bases 2007-2008
124
Acid
M
mol
Base
L
.00491 .0200
M
mol
.610 .00982
L
.0161
H2SO4 + 2NaOH  Na2SO4 + 2H2O
.00982 mol
x
NaOH
1 mol H2SO4
2 mol NaOH
= .00491 mol H2SO4
Acids and Bases 2007-2008
125
Acid
M
mol
Base
L
.0246 .00491 .0200
M = .00491 mol =
.0200 L
M
mol
.610 .00982
L
.0161
.0246 M
Acids and Bases 2007-2008
126
The reaction between HCl, a strong acid,
and NaOH, a strong base, can be
represented by
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
The pH profile of the titration of this
neutralization reaction is known as a
titration curve.
Acids and Bases 2007-2008
127
Consider the addition of 0.10 M NaOH solution (from a buret)
to an Erlenmeyer flask containing 25 mL of 0.10 M HCl.
Beyond the equivalence point,
the pH again increases slowly
with the addition of NaOH
Near the equivalence point,
the pH begins to rise steeply,
and at the equivalence point
the curve rises almost
vertically
Before the addition of NaOH,
the pH of the acid is given by
–log(0.10), or 1.00
When the NaOH is added,
the pH of the solution
increases slowly at first
Acids and Bases 2007-2008
128
It is possible to calculate the pH of a
solution at every stage of titration.
What is the pH of the solution after the
addition of 10.0 mL of 0.10 M NaOH to
25.0 mL of 0.10 M HCl?
Acids and Bases 2007-2008
129
0.10 mol NaOH =
x
1L
.0100 L
x = 0.00100 mol NaOH
0.10 mol HCl =
1L
x
.0250 L
x = 0.00250 mol HCl
0.00250 mol HCl
- 0.00100 mol NaOH
Excess
0.00150 mol HCl
Acids and Bases 2007-2008
acid
130
To determine the pH, you need to
calculate [H+]
Total
volume of
the
original
HCl and
NaOH
.00150 mol HCl
= .043 M HCl
.0350 L
pH = -log (.043)
pH = 1.37
Acids and Bases 2007-2008
131
At the equivalence point of a titration between
a weak acid and a strong base, the pH will
be greater than 7.
Acids and Bases 2007-2008
132
Because…
at neutralization
CH3COOH + NaOH  CH3COONa + H2O
CH3COONa
CH3COO- + Na+
This acetate ion
has an affinity for
the H+ ion in the
water, (thus leaving
the OH- behind,
making the solution
basic.)
133
The pKa of a weak
acid can be
determined
experimentally
The flat portion of the titration curve before
the equivalence point is called the buffer region.
In this part of the pH scale, the acid and
conjugate base are both present in significant
concentrations and the solution resists changes
in pH. As base is added to a solution in this
buffer region, acetic acid reacts with it to form
acetate ion, without a large change in pH.
Acids and Bases 2007-2008
134
1/2 way
At the half-equivalence point,
[CH3COOH] = [CH3COO-], so
Ka = [H3O+]
In the middle of the buffer region lies the halfequivalence point. Here the volume of base added is
half that required to reach the equivalence point and
half the acetic acid has been converted to the
conjugate base, acetate ion. This means that the
concentrations of acetic acid and acetate ion are equal.
If we examine the equilibrium expression at the halfequivalence point, we find something interesting…
Acids and Bases 2007-2008
135
At the half-equivalence point,
[CH3COOH] = [CH3COO-], so
Ka = [H3O+]
Taking the negative log of both sides yields
pKa = pH
Ka = 10-pH
This gives us an experimental way to determine the Ka of a
weak acid, and using a Ka table, the identity of an
unknown weak acid.
Acids and Bases 2007-2008
136
A slightly different curve results when you titrate a strong acid vs
a weak base. At the equivalence point of a titration between a
strong acid and a weak base, the pH will be less than 7.
Acids and Bases 2007-2008
137
Because…
at neutralization
HCl + NH3  NH4Cl
NH4Cl
NH3 + H+
The pH of less
than 7 is due to
the presence of
H+ ions formed
by the hydrolysis
of NH4+
Acids and Bases 2007-2008
138
Acid-Base Indicators
An indicator is usually a weak organic acid or
base that has distinctly different colors in its
nonionized (molecular) form and ionized form.
The end point of a titration occurs when the
indicator changes color.
However, not all indicators change color at the
same pH, so the choice of indicator for a
particular titration depends on the nature of
the acid and base used in the titration (that
is, whether they are strong or weak).
By choosing the proper indicator for a titration,
we can use the end point to determine the
139
equivalence point.
Let us consider a weak monoprotic acid that
we will call HIn. To be an effective
indicator, HIn and its conjugate base, ___,
Inmust have distinctly different colors.
HIn
One
color
H+ + InA
different
color
140
HIn
H+ + In-
If the indicator is in a sufficiently acidic
environment, the equilibrium, according
to Le Chatelier’s principle, shifts to the
left
__________
and the predominant color
HIn
will be that of ______,
Acids and Bases 2007-2008
141
HIn
H+ + In-
In a basic environment, the equilibrium
shifts to the right because…
The H+ and the OH- will form water,
thus removing the H+ from the system
and therefore shifting the equilibrium
to the right. The predominant color
will then be that of In-.
Acids and Bases 2007-2008
142
The end point of an indicator does not
occur at a specific pH; rather, there is
a range of pH within which the end point
will occur. In practice, we choose an
indicator whose end point lies on the
steep part of the titration curve.
Because the equivalence point also lies
on the steep part of the curve, this
choice ensures that the pH at the
equivalence point will fall within the
range over which the indicator changes
color.
Acids and Bases 2007-2008
143
Some Common Acid-Base Indicators
Indicator
In acid
In base
pH range
Thymol blue
Red
Yellow
1.2 - 2.8
Bromophenol blue
Yellow
Bluish purple
3.0 – 4.6
Methyl orange
Orange
Yellow
3.1 – 4.4
Methyl red
Red
Yellow
4.2 – 6.3
Chlorophenol blue
Yellow
Red
4.8 – 6.4
Bromothymol blue
Yellow
Blue
6.0 – 7.6
Cresol red
Yellow
Red
7.2 – 8.8
Phenolphthalein
Colorless
Reddish pink
8.3 – 10.0
The pH range is defined as the range over which the
indicator changes from the acid color to the base color.
Acids and Bases 2007-2008
144
This is THE END
of your CHEM II
NOTES!!!
Acids and Bases 2007-2008
145