Pharmaceutical Analytical Chemistry (PHCM223-SS16)
Lecture 3
ACID- BASE EQUILIBRIUM-III
“pH of buffers”
Dr. Rasha Hanafi
PHCM223,SS16
Lecture 3, Dr. Rasha Hanafi.
1
LEARNING OUTCOMES
By the end of this session the student should be able to:
1.
2.
3.
4.
5.
Define main properties of buffer solutions.
Apply Henderson-Hasselbalch equation
Determine buffer capacity.
Calculate the pH of buffered solutions.
Calculate the pH changes in buffered solution upon adding
acids or bases.
6. Write a scientific description of a buffer
7. Comprehend preparation steps of a buffer
PHCM223,SS16
Lecture 3, Dr. Rasha Hanafi.
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IMPORTANCE OF pH CONTROL ON HUMAN BODY
Most people who suffer from
imbalanced pH are acidic.
This condition forces the body
to borrow minerals (Ca, Na, K,
Mg) from vital organs and bones
to buffer (neutralize) the acid
and safely remove it from the
body.
Because of this strain, the body
can suffer severe and prolonged
damage due to high acidity
(may go undetected for years).
PHCM223,SS16
Lecture 3, Dr. Rasha Hanafi.
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WHAT IS A BUFFER?
A buffer is an acid-base equilibrium system that resists the
change in pH of the soln. upon addition of small amounts of
acids (H+) or bases (OH-).
weak acid and its conjugate
base
CH3COOH
CH3COONa
PHCM223,SS16
CH3COO- + H+
 CH3COO- + Na+
weak base and its conjugate
acid
NH4+ +OH-
NH3 + H2O
Lecture 3, Dr. Rasha Hanafi.
NH4Cl

NH4+ + Cl4
WHAT IS A BUFFER?
If H+ is added to solution, the equilibrium shifts to the left (common ion H+),
the pH remains constant.
If H+ is removed (e.g. by adding OH-) then the equilibrium shifts to the right,
releasing H+ to keep the pH constant.
0.01 M HCl
Human blood
Buffered at 7.4
 pH 7.3 
H+
H+
H+
Neutral
soln.
H+
OHOHOH-
OH-
0.01 M HCl
Water pH 7
Non-buffered
 pH 2 !!
Response to a rise in pH
HCO3- + H+
H2CO3
Response to a drop in pH
The main source of HCO3 is the salt, because the acid only weekly dissociates
-
PHCM223,SS16
Lecture 3, Dr. Rasha Hanafi.
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HOW DOES A BUFFER MAINTAIN A CONSTANT pH?
H+
H+
H+
H+
If too much H+ is added, the
equilibrium is shifted all the
way to the left, and there is
no longer any more HCO3- to
“absorb” H+. At that point the
solution no longer resists
change in pH; and is useless
as a buffer. Same applies if
too much OH- is added.
H2CO3
OHOHOH-
OH
H+
-
OH
-
H+
OH
H+
OH
-
-
HCO3-
H+
OH-
+ H+
The main source of HCO3- is the salt, because the acid only weekly dissociates
PHCM223,SS16
Lecture 3, Dr. Rasha Hanafi.
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pH OF A BUFFER SOLUTION
IA

HA(aq)
[M – x]
pH of buffer solution:
Ka =
[H+][A–]
[HA]
I+
H+
[x]
+
+
AA–(aq)
[A- + x]
x (A + x)
=
(M – x)
As long as the concentrations of these species are reasonably high (> 0.1 M), we can
neglect the ionization of the acid and the hydrolysis of the salt (Remember that
dissociation of weak acids is minimal Ka< ~ 10–4 and hence
M – x ≈ M and A + x ≈ A
+] [A-]
x
(A)
[H
thus, Ka ≈
≈
(M)
[H+] ≈
[H+] ≈ Ka
PHCM223,SS16
Ka[HA]
[A–]
Molarity of HA
Molarity of A–
[HA]
(general buffer system)
- log
pH = pKa + log
Lecture 3, Dr. Rasha Hanafi.
[conj. base ]
[acid]
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HENDERSON- HASSELBACH EQUATION
IA
HA(aq)

pH = pKa + log
I+
H+
+
+
AA–(aq)
[conjugate base ]
[acid]
THE BUFFER CAPACITY
The buffering capacity of a solution is a measure of the effectiveness of the
buffer.
The higher the capacity the larger the amount of protons or hydroxide ions
that the buffer can absorb with no significant change in pH.
The capacity depends on the amount of acid and conjugate base from which
the buffer is made. The larger the amount, the greater the buffering capacity.
When [conj. Base] = [acid]  log 1 = 0  pH = pKa
PHCM223,SS16
Lecture 3, Dr. Rasha Hanafi.
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Test yourself
The following diagrams represent solutions containing a weak acid HA
and/or its sodium salt NaA.
Which solutions can act as a buffer? (a) and (c)
Which solution has the greatest buffer capacity? (c)
The Na+ ions and water molecules are omitted for clarity.
PHCM223,SS16
Lecture 3, Dr. Rasha Hanafi.
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PREPARING A BUFFER SOLUTION WITH A SPECIFIC pH
Suppose we want to prepare a buffer solution with a specific
pH. How do we go about it?
Henderson Hasselbach equation indicates that if the molar concentrations of
the acid and its conjugate base are approximately equal, that is,
if [acid]= [conjugate base] log [conjugate base] /[acid] = 0  pH = pKa
Thus, to prepare a buffer solution, we work backwards:
1- we choose a weak acid whose pKa is close to the desired pH.
2- we substitute the pH and pKa values in the equation to obtain the ratio
[conjugate base]/[acid].
3- This ratio can then be converted to molar quantities for the preparation
of the buffer solution.
PHCM223,SS16
Lecture 3, Dr. Rasha Hanafi.
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WHAT TO REPORT WHEN WRITING ABOUT A BUFFER?
• The identity of the buffer (name or chemicals)
• The molarity of the buffer
• The pH of the buffer
Examples:
 “We used a 0.5M phosphate buffer, pH 7.4”
 “The reaction was carried out in a 0.1M boric acid –
sodium hydroxide buffer adjusted to pH 9.2.”
PHCM223,SS16
Lecture 3, Dr. Rasha Hanafi.
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Test yourself
PHCM223,SS16
Lecture 3, Dr. Rasha Hanafi.
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Test yourself
PHCM223,SS16
Lecture 3, Dr. Rasha Hanafi.
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SUMMARY
pH
Proton
availability
Buffer
solution
Stable
Weak acid / conjugate base
Weak base/conjugate acid
pKa
determined by the magnitude of [HA] and [A-].
PHCM223,SS16
Lecture 3, Dr. Rasha Hanafi.
pH
[acid]/[base]
Buffer capacity
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REFERENCES
1.
2.
3.
Chemistry, 10th ed., Raymond Chang, ISBN 978-0-07-017264-7, McGraw
Hill. Chapter 16
http://www.nutritionalhealthenterprises.com/articles.ews?articles.ewdi
d=11
https://www.youtube.com/watch?v=NJyAme5GVF8
PHCM223,SS16
Lecture 3, Dr. Rasha Hanafi.
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