Chemistry 12
UNIT 4
ACIDS AND BASES
PACKAGE #7
BUFFERS
Buffers are solutions which resist LARGE pH changes when SMALL additions of acid or
base are made to the solution.
Buffers typically consist of equilibrium systems of weak acids and their conjugate base
or weak bases and their conjugate acids.
You need substantial amounts of the acid and its conjugate (or base and its conjugate) in
order to qualify as a buffer,
IN THE LAB:
In order for a pH meter to be accurate, it must be standardized for the range of pH's that it
is going to read.
Buffer solutions can be prepared to provide a wide range of reliable pH standards.
(These kinds of buffer solutions were used as standards for the pH lab that we did where
we diluted HCl and CH3COOH to produce a certain pH value).
By mixing equal molar quantities of a weak acid such as CH3COOH and its salt
NaCH3COO ( i.e. the conjugate base of CH3COOH), you can standardize a pH meter
within a certain pH range.
IN NATURE:
BLOOD:
Buffers exist in human blood to maintain the pH at or near 7.4. (Slightly alkaline).
Hemoglobin is the oxygen carrier in the blood and is involved in the equilibrium.
HHb +
O2
(hemoglobin)
+
H2O
⇄
H3O+ +
HbO2(oxyhemoglobin)
If pH is too low (called acidosis of the blood when pH is less than 7.20), which means
that the [H3O+] is too high, the equilibrium shifts sufficiently left resulting in the
[HbO2-] being too low. The result is that O2 will not bind with HHb properly.
If pH is too high (called alkalosis of the blood when pH is > 7.50) and thus the [H3O+] is
too low, then the equilibrium shifts sufficiently right resulting in the [HbO2-] being
preferentially formed. The result is that the O2 is prevented from being released from
the blood.
To prevent alkalosis or acidosis, two buffer systems are present in our bodies to regulate
the [H3O+] . Otherwise eating tomatoes or drinking lemon juice would affect the blood
pH and it would be fatal!
1)
The CO2 / HCO3- System
A product of metabolism and respiration is CO2(g). When CO2(g) diffuses into the
blood, it reacts with H2O to form carbonic acid:
CO2
+
H2O
⇄
H2CO3
The carbonic acid could ionize:
H2CO3 + H2O ⇄ H3O+ + HCO3But if this happened, the [H3O+] of the blood would increase and the pH would go
down.
BUT we have an excess of HCO3- in our blood, so therefore the H2CO3 does not
ionize.
And breathing out CO2(g) (i.e. CO2(g) ⇄ CO2(aq) ) upsets the [H3O+] in the
following reaction:
CO2(aq)
+
H2O ⇄ H3O+ + HCO3-
Breathe out the CO2 and the reaction shifts left (also excess HCO3- causes shift to left)
2)
The H2PO4- / HPO4-2 System
Both of the above species are present in the blood and our cells (for bone, tooth, and
DNA maintenance). The buffer
H2PO4- + H2O ⇄ H3O+ + HPO4-2
stabilizes the pH of the cells to a large degree.
As you can see, buffer solutions resist changes in pH.
BUFFERING ACHIEVES TWO RESULTS:
1) stable pH
2) produces more neutral initial solution (conjugate common ion reduces dissociation).
EXAMPLES:
•ACIDIC BUFFER:
MIX 1.0 mole of CH3COOH and 1.0 mole of NaCH3COO and dilute to 1.0 L of
solution.
When making up a buffer solution, you must ensure that the salts used are soluble.
This solution is referred to as an acidic buffer because it buffers the pH in the acidic
region.
See the equilibrium that is established in the following sample calculations.
•BASIC BUFFER:
MIX 1.0 mole of NH3 and 1.0 mole of NH4NO3 and dilute to 1.0 L of solution.
When making up a buffer solution, you must ensure that the salts used are soluble.
This solution is referred to as an basic buffer because it buffers the pH in the basic region.
The equilbrium that is established is as follows:
NH4+ +
H2O
⇄
NH3
+
H3O+
Note that the NH4+ came from: NH4NO3  NH4+
The NO3- is a spectator, but the NH4+ hydrolyzes
+
NO3-
Even if you showed the basic hydrolysis of ammonia, you would still be able to calculate
the pH using a Kb = Kw/Ka instead of the Ka that you would have used in the above
equilibrium.
NH3 +
H2O
⇄
NH4+ +
OH-
LIMITATIONS OF BUFFERS:
There is a limit to the amount of H3O+ or OH- which can be neutralized by a buffer. If
there is 1 mol of the conjugate base present, then a maximum of 1 mole of H3O+ that
can be neutralized. (and similarily 1 mol of conjugate acid can neutralize 1 mol of OH-).
SAMPLE BUFFER CALCULATIONS:
1. Calculate the pH of an unbuffered 0.010M acetic acid solution.
2. Calculate the pH of a buffered 0.010M acetic acid solution.
3. Calculate the pH of an unbuffered 0.010M acetic acid solution in which 0.004 M of
H+ is being added from an acidic source.
4. Calculate the pH of a buffered 0.010M acetic acid solution in which 0.004 M of H+ is
being added from an acidic source.
5. Calculate the pH of an unbuffered 0.010M acetic acid solution in which 0.004 M of
OH- is being added from an basic source.
6. Calculate the pH of a buffered 0.010M acetic acid solution in which 0.004 M of OHis being added from an basic source.
SAMPLE BUFFER CALCULATIONS - Answers
1. pH = 3.38 2. pH = 4.74 3. pH = 2.40 4. pH = 4.38 5. pH = 4.57 6. pH = 5.11
EXTRA BUFFER QUESTIONS:
7. Calculate the pH of a solution prepared by mixing 40.0 ml of 0.60 M HCOOH with
60.0 ml of 0.15 M HCOOK.
8. Calculate the pH of a buffer solution prepared by mixing 25.0 ml of 0.20 M acetic
acid with 75.0 ml of 0.16 M sodium acetate.
9. How many grams of NaF must be added to 0.100 L of 0.20 M HF to produce a
solution with a pH = 3.00
10. Find pH of a solution made by dissolving 2.0 g NH4NO3 in 50.0ml 0.400 M NH3
11. Calculate the pH of a buffered 0.010M acetic acid solution in which 0.004 M of OHis being added from an basic source.
EXTRA BUFFER QUESTIONS - Answers
1.
2.
3.
4.
5.
3.32
5.12
0.57 g
9.13
pH = 5.11
Unit 4 review list – the second ½ of Unit 4
4.13 Hydrolysis
4.14 Ka Calcs
4.15 Kb calcs
4.16 Acid-Base Titrations - p. 158!!!
4.17 Indicators – p. 162-163 !!!!
4.18 Primary Standards
Titrations Curves – p. 176-177!!!!
4.19 Buffers
4.20 Buffers in Biology
4.21 Anhydrides
Acid Rain – your own study notes
Stay Tuned for:
One large, or several small review packages
for Unit 4