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INTRODUCTION
What are acids and bases? Svante Arrhenius noticed that acids release hy-
drogen ions in solution. He classified acids and bases in this way: acids are
compounds that dissociate, or break up, in water to give a proton; bases are
compounds that dissociate in water to accept a proton.
Acids that break up completely in water; these are strong acids. Acids that
do not dissociate completely in water are call weak acids. Can you guess
what bases that dissociate completely are called? Do you think weak bases
dissociate completely?
Things that don’t dissociate completely are most interesting - and often
overlooked when it comes to safey! Take vinegar, for example. Did you
know that this common kitchen item is actually a weak acid called acetic
acid? While it’s not dangerous in the form you buy at the grocery store be-
cause it’s very dilute (about 4-5% acetic acid), more concentrated solutions
of acetic acid can be quite dangerous even though it’s a weak acid.
Take a look at the dissociation of acetic acid in water.
Consider an acetic acid burn on your skin. Because acetic acid is neutral,
it can be absorbed into the skin readily. Now you’ve got acetic acid in your
hypodermis - the lower layer of your skin! If it were to remain neutral and
not dissociate, this wouldn’t be a problem, but it does dissociate. The ions
then attack the layers of skin and burn it. All the while, more weak acid can
enter the hypodermis causing further damage, and a very severe burn.
CH
3
COOH H
+ CH
3
COO
How about the chemical workers -
are they unionized or ionized?
H
2
O
Alternating Current
Switch
Bulb
Filament
Electrodes
Sample
Sample
ELECTROLYTES
Solution
We can do the same type of experiment with
solutions. Solutions are homogeneous mixtures
of a solute dissolved in a solvent (module #).
When substances are dissolved in water, the
solutions can be tested by seeing if the solu-
tion conducts electricity, just as the metal of the
screwdriver does. An electrolyte is a substance
whose aqueous solutions conduct electricity.
Electrolytes are often identified by experiments
testing the conductivity of solutions.
A simple conductivity experiment is the following:
A plastic screwdriver handle is placed against the
electrodes - no light! When the metal part is used
- bulb lights brightly! The handle is an insulator,
whereas the metal is a conductor!
In the following pictures, you will be asked to
determine if the solutions are conductors or non-
conductors. These pictures will help you realize
the defi nitions of some key chemical terms, and
hopefully some chemical concepts!
►The diagram below represents the apparatus
with electrodes dipped into liquid water. Is liquid
water a conductor? Look at the top left-hand
corner of the conductivity experiment (the light
bulb) to answer the question.
Be aware of the potential danger of alter-
nating current electricity. Do not attempt
any of these experiments without appropri-
ate supervision.
0.7 M Acetic Acid Solution.
Solute:
H
CO
2
CH
3
CO
2
H
(aq)
NH
4
Cl
-
OH
-
_______
HCO
3
-
►Name the ions in solution. Which species
is the anion? Which species is the cation?
Is the solute an electroyte?
►Name the ions in solution.
Name the molecules in solution.
Is the solute an electrolyte?
►Write the formula for the anion.
Write the formula for the cation.
Write the formula for the solute.
Write the formula for the solvent.
►Is the solute completely dissociated?
Complete the list above of the chemical spe-
cies in solution.
ELECTROLYTES
NaOH Na
+ OH
3
COOH H
+ CH
3
COO
Let’s do some ChemLogs in order to further
understand strong electrolytes.
►Why is the equilibrium arrow used for this
weak electrolyte dissociation?
Let’s do some ChemLogs in order to further
understand weak electrolytes.
Na
NaOH
Na
OH
CH
333
COOH
HH
CHCHCH
33
COOCOOCOO
CH
333
COOH
►List the weak electrolytes. List the strong
electrolytes.
Notice the solutes which dissociate completely
cause the light to shine brightly. The dim light
is caused by partial dissociation. Solutes that
dissociate completely are called STRONG
ELECTROLYTES. Solutes that dissociate only par-
tially are called WEAK ELECTROLYTES.tially are called WEAK ELECTROLYTES.
0 1 2 3 4 5 6 7 89 10
10 9 8 7 6 5 4 3 21
0 1 2 3 4 5 6 7 89 10
10 9 8 7 6 5 4 3 21
H
NH
44
Cl
(aq)(aq)
An Ammonium Cloride Solution.
Solute:
CO
2 (aq)2 (aq)
A Carbon Dioxide Solution.
Solute:
(x
mol L
mol L
mol L )
(x
mol L
mol L
mol L )
NaOH
(aq)(aq)
0.7 M Sodium Hydroxide Solution.
Solute:
CH
3
CO
2
H
Dissociation is the separation of ions that occurs
when an ionic compound dissolves.
An ACID is a proton (H
ion) donor. A BASE is
a proton acceptor. This is called the Bronsted-
Lowry system.Lowry system.
SO
4
2-
H
2
SO
4 (aq)
KCl
(aq)
HNO
3 (aq)3 (aq)
KOH
(aq)(aq)(aq)(aq)
K
pH is a measure of the concentration of H
ELECTROLYTES
►Fill in the “empty” ChemLogs for each of the
following experiments:
H
OH
-
NO
3
-
H
K
0.8 M Nitric Acid Solution.
0.8 M Potassium Hydroxide Solution.
0.8 M Potassium Chloride Solution.
0.8 M Sulfuric Acid Solution.
H
2
SO
4
KCl
HNO
33
KOH
H
2
SO
4
KCl
HNO
3
KOH
SO
4
2-
K
Cl
2H
►Strong electrolytes that dissociate completely
to give H
ions are called strong acids. Name
the strong acids on this page.
►Strong electrolytes that completely dissociate
to give a H
acceptor are called strong bases.
Name the strong bases on this page.
►There are other solutes that are strong elec-
trolytes. These are called salts. Name the salts
on this page.
►For each of the solution on this page, state
whether the ph is less than 7, greater than 7, or
equal to 7.
► Complete the chemical reaction and the
chemLog where necessary.
(x
mol L
mol L
mol L )
Solute:
Solute:
Solute:
Solute:
Cl
10 9 8 7 6 5 4 3 21 0 1 2 3 4 5 6 7 89 10
10 9 8 7 6 5 4 3 210 1 2 3 4 5 6 7 89 10
0 1 2 3 4 5 6 7 89 10 10 9 8 7 6 5 4 3 21 0 1 2 3 4 5 6 7 8910 10 9 8 7 6 5 4 3 21
(x
mol L
mol L
(x10 mol L )
mol L
mol L
mol L )
(x
mol L
mol L
mol L )
When pH is between 1 and 7, the solution is
acidic. When the pH is between 7 and 14, the
solution is basic. The pH of pure water is 7 (neu-
tral pH).
pH = –log [H
]
CONJUGATE ACIDS AND BASES
HNO
2
H
NO
2
-
We’ve already seen the following reaction earlier,
but let’s study it more.
We know that HNO
2
is an acid because it gives a
H
ion;
HNO
2
H
NO
2
-
However, at equilibrium, NO
2
-
accepts a proton to
become HNO
2
. That makes NO
2
-
a base;
HNO
2
H
NO
2
-
So, HNO
2
is an acid, and NO
2
-
is its conjugate
base. On the other hand, NO
2
-
is a base, and
HNO
2
is its conjugate acid.
The species formed when an acid loses a proton is
called a CONJUGATE BASE. The species formed
when a base gains a proton is called a CONJU-
GATE ACID.
A simple way to remember how to identify acids
and their conjugate bases, and bases and their
conjugate acids, is the following table:
from the acid,
substract H
conjugate base + H
to the base,
add H
conjugate acid
►What do you notice about the arrangement of
the water molecules around the H
ion in picture
The oxygen atoms in the water molecules have a
slightly negative charge (negative dipole). The H
ion is actually bonded to several water molecules.
As you can see from the picture, a more complete
way of writing the reaction of the dissociation of
nitrous acid is
NO
2
-
H
Nitrous Acid Solution.
HNO
2
HNO
2
H
3
O
NO
2
-
+ H
2
O
We call the H
3
O
ion the hydronium ion.
Note: Water must be acting as a base in this reac-
tion (see next section).
HNO
2
H
11
O
5
NO
2
-
+ 5 H
2
O
Obviously, this could get complicated (and long!)
very fast. A more common way to
►Identify the acid and conjugate base in the fol-
lowing reaction:
We chemists usually just write the most simpli-
fi ed version using H
. You may notice that some
books do use the hydronium ion. As a student, it
is important to recognize that these two versions
indicate the same thing - a proton surrounded by
water molecules.
CH
3
COOH H
+ CH
3
COO
ACIDS
BASES
When NH
3 (g)
is dissolved in water, it acts as a
base. Bases accept protons.
NH
3
NH
4
OH
-
+ H
2
O
Notice that H
2
O is acting as an acid. We cannot
exclude H
2
O as a reactant because we need to
have a balanced chemical reaction.
The dissociation of water gives both an acid
(substance that provides a proton) and a base
(substance that accepts a proton). We call water
AMPHIPROTIC.
K
w
) = 1 x 10
H
2
O
H
+ OH
In the general equilibrium module, we learned to write
equilibrium constants. For water,
K
c
[H
][OH
]
[H
2
O]
Since the concentration of H
2
O in water is 55.6 M,
it can be regarded as a constant. A simplified
version of K
c
is called the ion-product constant
of water.
K
w
= [H
][OH
]
K
w
is a constant. At 25°, the concentrations of
both H
and OH
are 1 x 10
mol L
►Will the value of K
w
change at varying tem-
peratures? Explain.
Remember the definition of pH from earlier in this
module.
CONJUGATE ACIDS AND BASES
Other examples include:
pH = – log [H
]
pOH = – log [OH
]
pK
w
= - log K
w
►Determine the pH, pOH, and pK
w
for the
dissociation of water. Check your answers by
using the fact that pH + pOH = pK
w
►What is the pH of a solution in which the pOH = 4?
Which is greater for this solution - the value of
[H
] or the value of [OH
]?
OH
H
H
2
O + H
2
O
H
3
O
+ OH
which can be simplified to
For most purposes, we can simplify pictures of
solvent molecules. The neutral water molecules
in the following picture are simplified to a blue
background. The ions in the water are thus shown
more clearly.
Now that we’ve seen that acids in water give hy-
drogen ions, let’s look at pure water. We’ve stud-
ied the following picture earlier and determined
that water is a non-electrolyte because the light
bulb didn’t light. While all of this is true, there’s
more to this story.
H
2
O
Pure Distilled Water
►For a review, state the defi nition of a non-
electrolyte.
When a very sensitive light is used in the conduc-
tivity experiment, the light lights. How can this
be? Well, pure water does form ions in solution;
it forms protons and hydroxide ions. But, the con-
centration of these ions is so low that a sensitive
light is needed in the conductivity experiment to
see evidence of the ions. To show this in the form
of a reaction, we write
This calculation can be used for more than just the
concentration of H
. The pX of X (where X is any
number) is equal to the –log of X.
►Water is an example of a molecule which gives
2 conjugate acid/base pairs. What are they?
Note that there are no units associated with equi-
librium constant values, K. This is standard, and
you are not expected to show units.
º C K
w
pK
w
0 0.12 x 10
10 0.29 x 10
25 1.01 x 10
Ion Product of Water.
INDICATORS
Water molecules
HIn
In
-
HIn H
C
15
H
15
N
3
O
2
H
+ C
15
H
14
N
3
O
2
-
H
One way to determine the pH of a solution is to use an indicator. An indicator exists in different colored forms
depending on whether the compound is protonated or unprotonated. The indicator shown is purple cabbage extract.epending on whether the compound is protonated or unprotonated. The indicator shown is purple cabbage extract.
►The 1 x 12 array shown below shows the pH where the color changes.
Mark which pH values show [HIn] > [In
-
].
Mark which pH values show [HIn] < [In
-
].
Mark which pH values show [HIn] = [In
-
].
K
a
[In
-
]
[HIn]
=
[H
]
Therefore, when [HIn] = [In
-
],
K
a
[In
-
]
[HIn]
=
[H
]
= 1, so, K
a
= [H
].
►Which part of the solution causes the color, the solute or the solvent?
H
pH =pH = 4 5 6 7 8 9
►Is purple cabbage indicator an acid or a
base?
►Write the K
a
expression for HIn H
Rearranging the K
a
expression gives
To see “pure” color,
[In
-
]
[HIn]
=
for blue. For “pure” purple,
[In
-
]
[HIn]
=
K
a
=
[H
]
K
a
=
[H
]
K
a
=
[H
]
10 K
a
=
[H
]
Therefore, the pH range of an indicator is pK
a
Malachite green
Cresol red
Metacresol purple
Orange IV
Thymol blue
Methyl yellow
Bromophenol blue
Congo red
Methyl orange
Methyl red
Chlorophenol red
Litmus
para-Nitrophenol
Bromocresol purple
Bromocresol green
Neutral red
Phenol red
Cresol red
Curcumin α
meta-Cresol purple
Thymol blue
Phenolphthalein
Malachite green
Thymolphthalein
Alizarin yellow R
Curcumin ß
Clayton yellow
Various indicators
Bromothymol blue
Thompson-Markow
Universal Indicator
pH Ranges of Indicators
pH
[H
]
10
Here is a list of several acid/base indicators and
the pH range in which they are useful:
A universal indicator, shown at the bottom, is a
mixture of several indicators. Universal indica-
tors have a wider range of usefullness because
they mix indicators.
INDICATORS
►What is the pH range of the Thompson-Mar-
kow universal indicator?
What is [H
] at pH 0?
►The Thompson-Markow Universal Indica-
tor is a mixture of six indicators. Metacresol purple
is one of them; name the 5 others.
INDICATOR WELL TRAYS
pH
pH
pH
pOH
pOH
pOH
1 2 3 4 5 6 7 8 9 10 11 12 13 14
14 13 12 11 10 9 8 7 6 5 4 3 2 1
1 2 3 4 5 6 7 8 9 10 11 12 13 14
1 2 3 4 5 6 7 8 9 10 11 1213 14
14 13 12 11 10 9 8 7 6 5 4 3 2 1
14 13 12 11 10 9 8 7 6 5 4 3 2 1
pH
pOH
1 2 3 4 5 6 7 8 9 10 11 12 13 14
14 13 12 11 10 9 8 7 6 5 4 3 2 1
pH
pH
pH
pOH
pOH
pOH
1 2 3 4 5 6 7 8 9 10 11 12 13 14
14 13 12 11 10 9 8 7 6 5 4 3 2 1
1 2 3 4 5 6 7 8 9 10 11 12 13 14
1 2 3 4 5 6 7 8 9 10 11 1213 14
14 13 12 11 10 9 8 7 6 5 4 3 2 1
14 13 12 11 10 9 8 7 6 5 4 3 2 1
pH
pOH
1 2 3 4 5 6 7 8 9 10 11 12 13 14
14 13 12 11 10 9 8 7 6 5 4 3 2 1
pH
1 2 3 4 5 6 7 8 9 10 11 12 13 14
pH
pH
pOH
pOH
pOH
1 2 3 4 5 6 7 8 9 10 11 12 13 14
14 13 12 11 10 9 8 7 6 5 4 3 2 1
1 2 3 4 5 6 7 8 9 10 11 1213 14
14 13 12 11 10 9 8 7 6 5 4 3 2 1
14 13 12 11 10 9 8 7 6 5 4 3 2 1
pH
pOH
1 2 3 4 5 6 7 8 9 10 11 12 13 14
14 13 12 11 10 9 8 7 6 5 4 3 2 1
Below are several chemLogs showing the pH and pOH for each
well. Each well contains a solution of methyl orange at the differ-
ent pH values listed..
►What can you determine
about the relationship between
pH and pOH?
►Complete the following
equation:
pH + pOH = ______
►What is the pK
a
of methyl
orange?
►The approximate pH
range of any indicator is the
range between the pK
a
and the pK
a
pH range of methyl orange?
►Remember that K
a
x K
b
K
w
for any conjugate acid/
base pair. What is the pK
b
for
methyl orange? Using pK
b
determine the pOH range for
methyl orange.
pH
TITRATION
ACID BASE EQUILIBRIAACID BASE EQUILIBRIA
HCl
Five drops of 0.01 M HCl were added to each of
the wells in the 1 x 12 tray.
►Is the solution in each of these
wells acidic, neutral, or basic?
Are the solutions at equilibrium?
To each well above, drops of 0.01 M NaOH were
added. The numbers above the wells represents
the drops of NaOH added.the drops of NaOH added.
►Which solution is acidic?
Which solution is neutral?
Which solution is basic?
Where are the water molecules that are shown
coming from?
Write the net ionic equation for this titration
reaction.
Is [H
] higher in the well with 1 drop of base
added, or 5 drops added? What about pH?
Is [H
] higher in the well with 5 drop of base
added, or 10 drops added? What about pH?
►So, as [H
] increases, pH decreases, and
vice versa. What happens to pH as [OH
-
]
increases?
►What is the pH of the 0.01 M HCl that we
started with?
H
2
O
DEFINITION:
pH = - log [H
]
NaOH NaCl
HCl H
2
NaOH NaCl O
A solution of strong base, NaOH, can be titrated
with a solution of strong acid, HCl. The overall
chemical reaction for the titration is the following:chemical reaction for the titration is the following:
pH
Drops of 0.01 M NaOH Added
A B C D E B
�
�
�
�
�
�
�
�
�
��
��
���
►Experiment B is a strong acid titrated with
a strong base. Try to produce the titration
curve for this experiment. The chemLog is
shown at the right to help you with this.
TITRATION
ACID BASE EQUILIBRIAACID BASE EQUILIBRIA
Thompson-Markow
Universal Indicator
Titration of Strong Acid
with Strong Base
Titration of Weak Acid
with Strong Base
Titration of Strong Acid
with Weak Base
Titration of Weak Acid
with Weak Base
HCl
H
2
O
NaOH
NaCl
A solution of HCl is shown in each of the drawings.
TITRATION
ACID BASE EQUILIBRIAACID BASE EQUILIBRIA
���� ����
���� ����
►What is the limiting reagent for each of the
sketches that you drew?
►Which of the sketches that you drew is at
equilibrium?
►Are the drawings of HCl in solution (before
NaOH is added) at equilibrium?
►Draw a sketch for each sample after the
NaOH shown has been added. The water mol-
ecules are shown for you by the pink background.
A LIMITING REAGENT is a reactant that governs
the maximum yield of product that is possible.
For example, when 1 NaOH is added to a solu-
tion with 10 HCl, the NaOH is the limiting reagent
because it governs how many NaCl and water
molecules can be formed in the reaction.
►Is the pH of each of the solutions that you
have drawn acidic, neutral, or basic?
HCl H
2
NaOH NaCl O
BUFFERS
In experiments, we often need to control the [H
]
of the solution. This is done with buffers. A buff-
ered solution resists changes in [H
]. Buffers can
have a predetermined pH, with acidic buffers at
any chosen value between pH 0-7, and basic buf-
fers at any chosen value between pH 7-14.
Two dilute
solutions of
acetic acid
at the same
concentration.
Add acid.
Add base.
When acid is added, the buffer
accepts protons.
When base is added, the buffer
donates protons.
Add more
acid.
Add more
base.
The buffer capacity is an indication of the amount
of acid or base that can be added before the buffer
loses its ability to resist the pH change.
►At which point in each of the additions above
was the buffer capacity exceeded?
acetic acid
acetate ion
proton
sodium
Acidic buffers are solutions containing weak acids
and the salts of weak acids, eg. CH
3
COOH and
CH
3
COONa. Basic buffers are solutions contain-
ing weak bases and the salts of weak bases, eg.
NH
3
and NH
4
Cl.
CH
3
COOH H
+ CH
3
COO
CH
3
COONa Na
+ CH
3
COO
CH
3
COOH H
+ CH
3
COO
CH
3
COONa Na
+ CH
3
COO
HCl
NaOH
HCl
NaOH
HENDERSON-HASSELBALCH EQUATION
How can we design a buffer with a specifi ed pH?
Let’s examine the chemical reaction for solution
containing a general weak acid and its salt - a
buffer.
K
a
x
[A
[A
[A ]
[HA]
HA H
+ A
+ A
+ A
►Write the K
a
expression for this reaction.
After rearranging the K
a
expresstion to determine
what [H
] is equal to, we determine that
[H
] =
We can now determine the pH. Write the pH expres-
sion by taking the negative log of each side of the
[H
] expression.
pH = – log [H
] =
K
a
x
[A
[A
[A ]
[HA]
[A
[A
[A ]
[HA]
= – log K
a
[HA]
[A
[A
[A ]
= pK
a
Since acidic buffers are solutions of weak acids
and their salt, HA loses only a tiny fraction of its
protons. What are the dominant species (the most
concentrated) in solution? They HA, becasue HA
is a weak acid, and A
is a weak acid, and A
is a weak acid, and A , because it comes from the
complete dissociation of the salt NaA. Because
they are so dominant in solution, we can approxi-
mate [HA] and [A
mate [HA] and [A
mate [HA] and [A ] of a buffer by their initial con-
centrations. The Henderson-Hasselbalch equation
states that:
pH ≈
[HA]
[A
[A
[A ]
pK
a
initial
Let’s calculate the pH of a buffer solution in which
0.040 M CH
3
COONa
(aq)
and 0.080 M CH
3
COOH
(aq)
are mixed at 25 °C.
►Identify the acid and its conjugate base in this
mixture. Identify the ions from the salt in this
mixture.
►Write the equilibrium equation for acetic acid,
and rearrange the K
a
for this equation to give
[H
].
►Use the Henderson-Hasselbalch equation to
find the pH.
NaA Na
+ A
A
A
