Chemistry 12
Notes on Unit 4
Buffer Solutions
A buffer solution is a solution which resists changes in pH when a small amount of acid
or base is added. Or we could say it minimizes the change in pH when acid or base is
added.
Weak Acid Equilibria and the Common Ion Effect
Say you had some 1.0 M acetic acid (CH3COOH) solution.
This equilibrium becomes established:
CH3COOH + H2O
H3O+ +
1.0
0
0
-x
+x
+x
[I]
[C]
[E]
CH3COO-
Low conc. at equilibrium
A little less than 1.0 M at
equilibrium
Low conc. at equilibrium
Since CH3COOH is a WEAK acid, the [H3O+] and [CH3COO-] are quite low at equilibrium.
Now, lets add some sodium acetate (NaCH3COO) to the equilibrium mixture so that [
CH3COO-] is 1.0 M.
When we do this the [ CH3COO-] obviously goes up. However, by LeChatelier’s Principle,
the equilibrium will shift to the LEFT, causing [H3O+] to decrease and [CH3COOH] to
increase
CH3COOH
Concentration is a now
a little > 1.0 M
+
H2 O
H3O+
Concentration is a
now very low!
+
CH3COO-
Concentration is a now a
little < 1.0 M
So what we have produced is a solution that has a fairly high ( ≈ 1M ) of a WEAK ACID
(CH3COOH) and a WEAK BASE (CH3COO- ) in the same solution. This is how a buffer
solution is prepared. Since the acid and the base are both WEAK, they don’t neutralize
each other like a mixture of a SA and SB would. They co-exist in this equilibrium unless
disturbed!
Chemistry 12
Notes on Unit 4
There are two kinds of Buffer Solutions:


A Weak Acid and the Salt of It’s Conjugate Base (WASCB) eg. 1.0 M CH3COOH
& 1.0 M NaCH3COO
A Weak Base and the Salt of It’s Conjugate Acid (WBSCA) eg. 1.0 M NH3 & 1.0 M
NH4Cl
-
The first type (WASCB) or Acidic Buffers are useful as buffers in the acidic range
(solutions in which pH is 7 or lower)
-
The second type (WBSCA) or Basic Buffers are useful as buffers in the basic range
(solutions in which pH is 7 or higher )
NOTE: Buffers CANNOT be prepared using any STRONG ACIDS or STRONG
BASES!!! Strong acids and bases are too reactive, and will not remain in an equilibrium
mixture. They will react! So mixtures like 1.0 M HCl and 1.0 M NaCl or 1.0 M CH3COOH
and 1.0 M NaOH CANNOT be Buffers!
How Buffers Work to Minimize the Change in pH When Acids or Bases are
Added
Consider the buffer solution made up of 1M CH3COOH and 1M NaCH3COO.
The equation representing the equilibrium present in this buffer solution is:
CH3COOH
1M
+
H2O
H3O+
low
+
CH3COO1M
Now, we add a small amount of HCl to this solution.
The HCl produces H3O+, so the [H3O+] will immediately increase. (and the pH will go
down).
However, since this is in equilibrium and there is plenty of CH3COO- available, the
equilibrium will SHIFT to the LEFT and [H3O+] will go back down again (but not quite to
it’s original value)
This can be shown on a graph of [H3O+] vs. Time:
As equilibrium shifts left, [H3O+] decreases to
partially compensate for the sudden increase
[H3O+]
A very small net increase in
the [H3O+]
Time
HCl added
Chemistry 12
Notes on Unit 4
Using the same buffer solution:
CH3COOH
1M
+
H3O+
low
H2 O
+
CH3COO1M
This time we add a small amount of base NaOH:
The NaOH produces OH- which neutralizes H3O+ so the [H3O+] will immediately
decrease. (and the pH will go up).
CH3COOH + H2O
H3O+ +
CH3COOHowever, since this is in an equilibrium and there is plenty of CH3COOH available, the
equilibrium will SHIFT to the RIGHT and [H3O+] will go back up again (but not quite to
it’s original value)
Equilibrium shifts to the RIGHT
CH3COOH
+
H2O
H3O+
+
CH3COO-
This can be shown on a graph of [H3O+] vs. Time:
[H3O+]
A very small net decrease
in the [H3O+]
As equilibrium shifts right, [H3O+] increases to
partially compensate for the sudden decrease
Time
So, in summary, this buffer maintains a relatively constant pH when a small amount of
acid or base is added to it! ------- This is the function of a buffer solution! -------Basic Buffers work using the same principles. Work through the example by
filling in the blanks…..
Eg.) A buffer solution is prepared using 1M NH3 and 1M NH4Cl (WBSCA)
a) Write the equilibrium equation describing this buffer.
_______________________________________________________
b) When a small amount of HCl (SA) is added, the [OH-] quickly _____creases (the pH
goes ____ )
c) As a result, the equilibrium shifts to the _____________, and the [OH-] gradually
_____creases. (the pH goes back _________ )
d) As a result of adding HCl, there was a small net ____crease in the [OH-] (a small net
___ crease in pH )
Chemistry 12
Notes on Unit 4
e) Draw a graph of [OH-] vs. Time to illustrate what happened in b  d. Label each part.
[OH-]
Time
HCl added
f) Draw a graph of pH vs. Time to illustrate what happened in b  d.
pH
Time
HCl added
Eg.) A buffer solution is prepared using 1M NH3 and 1M NH4Cl (WBSCA)
a) Write the equilibrium equation describing this buffer.
_________________________________________________________
b) When a small amount of NaOH (SB) is added, the [OH-] quickly ___creases (the pH
goes ____ )
c) As a result, the equilibrium shifts to the _____________, and the [OH-] gradually
_____creases. (the pH goes back down )
d) So, as a result of adding NaOH, there was a small net _____crease in the [OH-] (a
small net ___ crease in pH )
e) Draw a graph of [OH-] vs. Time to illustrate what happened in b  d. Label each part.
[OH-]
Time
NaOH added
Chemistry 12
Notes on Unit 4
f) Draw a graph of pH vs. Time to illustrate what happened in b  d.
pH
Time
NaOH added
So, in summary, this buffer _____________________ when a small amount of acid or
base is added to it!
Limitations of Buffers
Say we have a buffer solution prepared using 1M CH3COOH and 1M NaCH3COO. The
equilibrium describing this buffer solution is:
CH3COOH
1M
+
H3O+
low
H2 O
CH3COO1M
+
Let’s say we add 1.5 moles of HCl to 1 litre of this solution. The [H3O+] will immediately go
up to 1.5 M. This is more than the 1 M CH3COO- can react with. There will still be an
excess of H3O+ large enough to bring the pH down significantly.
In the above case we have overcome the limitations of our buffer and it cannot hold the
pH relatively constant any more. Buffers only maintain a relatively constant pH when
SMALL amounts of acid or base are added to them!
Looking at the “buffer region” of the titration curve for a WA-SB titration illustrates how
the buffer “loses control” of the pH when the [base] overcomes the buffer:
Tit rat ion of weak ac id (25 c m ³) wi th s t rong bas e.
14
12
10
Point at which the buffer is
overcome by the base and is
unable to minimize the change
in pH
8
pH
Buffer
Region
6
4
2
0
10
20
30
40
volum e (c m ³)
Volume
of Base (mL)
50
Chemistry 12
Notes on Unit 4
Uses of Buffers
 Calibration of pH meters
 Control of pH in industrial
reactions
 Used in maintaining water quality
 Pools and hot tubs
Biological Buffer Systems
 Wine making
 pH balanced shampoos and
deodorants
 Soil pH
 Minimizing effects of acid rain
Pg . 182-183
For Hemoglobin to work properly, the pH of the blood needs to stay very close to 7.35
Equilibrium:
HHb
+
hemoglobin
O2

H3O+
+ HbO2oxyhemoglobin
When you inhale the [O2] in lungs is high. This diffuses through the thin alveoli walls into
the blood. So the equilibrium above shifts to the RIGHT, producing more oxyhemoglobin.
This “oxygenated” blood then takes the oxyhemoglobin to the cells of the body where [O2]
is low. Because [O2] is low, the equilibrium shifts to the LEFT, releasing O2 to the cells
where it can be used for cellular respiration.
Buffer Systems in the Blood
CO2 is produced during cellular respiration. It dissolves in the blood and can be thought of
as a solution of “carbonic acid” (H2CO3 ) Also present in our blood stream is
“bicarbonate” (HCO3-) which is the conjugate base of H2CO3. So we have an acidic
(WASCB) buffer system in our blood stream:
H2CO3
+
H2 O
H3O+
+
HCO3-
Actually exists as CO2 & H2O
When the [H3O+] tends to fluctuate in our blood, this buffer maintains the pH as close as
possible to 7.35
When a person “hyperventilates”, too much CO2 is lost and this equilibrium shifts to the
left, decreasing the [H3O+] and therefore increasing the pH. This can cause the person to
“black out”.
CO2 can be brought back up by “bag breathing” in a paper bag.
Read though the rather chemically interesting “chicken farmer” exercise 143
on page 183
HW pg 181 # 131 - 138
Chemistry 12
Notes on Unit 4
Answers
Consider the buffer solution made up of 1M CH3COOH and 1M NaCH3COO.
The equation representing the equilibrium present in this buffer solution is (just the same as the
ionization of the weak acid, CH3COOH):
CH3COOH
1M
+
H3O+
low
H2 O
+
CH3COO1M
Draw a graph of pH vs. Time when a small amount of HCl is added to the buffer above (Explain
each part)
pH rapidly decreases as
HCl (H3O+) is added
There is a small net DECREASE in
pH in the overall procedure.
pH
The equilibrium above shifts to
the LEFT, decreasing [H3O+], so
pH gradually increases.
Time
HCl added
CH3COOH
+
H3O+
H2O
+
CH3COO-
Draw a graph of pH vs. Time when a small amount of NaOH is added to the buffer above (Explain
each part)
The equilibrium above shifts to the right,
bringing [H3O+] up and pH down
pH
There is a small net INCREASE in pH in the overall procedure.
pH rapidly increases when NaOH is
added and [H3O+] decreases
Time
NaOH added
Chemistry 12
Notes on Unit 4
Basic Buffers work using the same principles. Work through the example by filling in the blanks…..
Eg.) A buffer solution is prepared using 1M NH3 and 1M NH4Cl (WBSCA)
a) Write the equilibrium equation describing this buffer.
NH3
+ H2O ⇆
NH4+
+ OH-
b) When a small amount of HCl (SA) is added, the [OH-] quickly decreases (the pH goes down )
c) As a result, the equilibrium shifts to the right, and the [OH-] gradually increases. (the pH
goes back up )
d) So, as a result of adding HCl, there was a small net decrease in the [OH-] (a small net decrease
in pH )
e) Draw a graph of [OH-] vs. Time to illustrate what happened in b  d. Label each part.
[OH-] rapidly decreases
as HCl (H3O+) is added
There is a small net DECREASE in
[OH-] in the overall procedure.
[OH-]
The equilibrium above shifts to
the RIGHT, increasing [OH-]
Time
HCl added
f) Draw a graph of pH vs. Time to illustrate what happened in b  d.
pH rapidly decreases as
HCl (H3O+) is added
and [OH-] decreases
There is a small net DECREASE in
pH in the overall procedure.
pH
The equilibrium above shifts to
the RIGHT, increasing [OH-], so
pH gradually increases.
Time
HCl added
Chemistry 12
Notes on Unit 4
Eg.) A buffer solution is prepared using 1M NH3 and 1M NH4Cl (WBSCA)
a) Write the equilibrium equation describing this buffer.
NH3
+ H2O ⇆
NH4+
+ OH-
b) When a small amount of NaOH (SB) is added, the [OH-] quickly increases (the pH goes up )
c) As a result, the equilibrium shifts to the left, and the [OH-] gradually dereases. (the pH
goes back down )
d) So, as a result of adding NaOH, there was a small net increase in the [OH-] (a small net increase
in pH )
e) Draw a graph of [OH-] vs. Time to illustrate what happened in b  d. Label each part.
The equilibrium above shifts to the LEFT,
bringing [OH-] down.
There is a small net INCREASE in [OH-] in the overall procedure.
[OH-]
[OH-] rapidly increases when NaOH
is added.
Time
NaOH added
f) Draw a graph of pH vs. Time to illustrate what happened in b  d.
The equilibrium above shifts to the LEFT,
bringing [OH-] down, thus pH also decreases.
There is a small net INCREASE in pH in the overall procedure.
pH
[OH-] rapidly increases when NaOH
is added. This causes pH to increase.
Time
NaOH added
So, in summary, this buffer minimizes (or resists) changes in pH when a small amount of acid or
base is added to it!