Medical Biochemistry and Molecular Biology Department
MEDICAL BIOCHEMISTRY
AND
MOLECULAR BIOLOGY DEPARTMENT
PRACTICAL GUIDE NOTES
ON
PH AND
BUFFERS
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Medical Biochemistry and Molecular Biology Department
PH AND
BUFFERS
ILO of the current topic:
By the end of this topic, the student will be able to: Identify different pH
solutions.
Each student should be checked for the ability to:
•
Define Acids/Bases, and Neutral solutions.
•
Observe general properties of acids and bases.

Differentiate between strong and weak acids and bases
•
Define pH, pKa, Kw
•
Show how to measure the pH of solutions of acids or bases
•
Describe how the indicators work.
•
Explain the action of a buffer
Background: pH Level - Similar to the body temperature, which has to be
maintained; the pH level in the various body fluids are kept at a narrow
range (fig) for the metabolic reactions to proceed properly. In chemistry,
pH is a measure of the acidity or alkalinity of an aqueous solution. Pure
water is said to be neutral, with a pH close to 7.0 at 25 °C (77 °F).
Solutions with a pH less than 7 are said to be acidic and solutions with a
pH greater than 7 are basic or alkaline. pH measurements are important in
medicine; because any minimum change of plasma pH more or less than
7.4 will be fatal.
Dissociation of Water:
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Water has limited capacity to dissociate (ionize) into H+ & OHIn fact, the H+ is always hydrated & present as hydronium ion (H3O +). While,
the ability of acids to react with bases depends on the tendency of hydrogen
ions to combine with hydroxide ions to form water
Both reactions take place simultaneously, but (V1) is so much faster
than (V2) that only a minute fraction of H2O molecules are dissociated.
Ion Product of Water
1 Liter of water has a mass of 1000 g, MW of water= 18g.
The number of moles in 1000 g of H2O = Weight in g/ Molecular weight.
So, (1000 g) / (18 g mol–1) = 55.5 mol.
For your knowledge the probability that hydrogen exist in water as ion = 1.8 x
10-9
The molar concentration of [ H+] = The probability that hydrogen exist in
water as ion X molar conc. of water
= 1.8 x 10-9 x 55.5 == 99.9x 10-9 ~ 102x 10-9= 10-7 mole /L
So an average of 10-7 mole of H+ and Of OH- will be dissociated at any time
([H+] = [OH-] are equal in water because it is neutral).
N.B: Square brackets [ ] refer to the concentrations (in moles /L) of the
substances they enclose.
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Medical Biochemistry and Molecular Biology Department
[H+] [OH–] = 1.00 × 10–14
This expression is known as the ion product of water, and it applies
to all aqueous solutions, not just pure water (look at the following table).
Acidic Solution
[H+] > [OH–]
Alkaline (Basic) Solution
[H+] < [OH–]
Neutral Solution
[H+] = [OH–] = 1.00×10–7 mol L–1
PH
pH is the negative logarithm of the concentration of hydrogen ions in a
solution.
As this ions concentration is absolutely small, it is suitable to express it by
means of logarithms of base ten.
pH = - log [H+]
And
pOH = – log [OH–]
BASES:
A base is a substance that can Accept proton (yields an
excess of hydroxide ions when dissolved in water).
NaOH(s) → Na+(aq) + OH–(aq)
A strong base completely dissociated in dilute aqueous solution, e.g. NaOH ,
KOH.
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Medical Biochemistry and Molecular Biology Department
ACID
An acid is a substance that can donate
proton (yields an excess of hydrogen ions
when dissolved in water).

When the acid, HA, loses a proton it forms a base, A-. When the base,
A-, accepts a proton back again, it obviously reforms the acid, HA.
These two are a conjugate pair [ A- is the conjugate base., and HA is
its conjugate acid].
Hydrochloric Acid
HCl +H2O→ H3O + + Cl–
Sulfuric Acid
H2SO4+H2O→ H3O + + HSO4–
Hydrogen Sulfite Ion
H2SO3+2H2O→ 2H3O + + SO32–
Acetic Acid

C H3COOH+H2O→ H3O + + H3COO-
The strong acids such as HCl and HNO3 are effectively
100% dissociated in solution. Most organic acids, such as
acetic acid, are weak; only a small fraction of the acid is
dissociated in most solutions.

To recognize the pH of week acids that only dissociate to a
small degree in solution (Partial dissociation). We quantify
their pH by using, what is called dissociation constant (Ka).
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Medical Biochemistry and Molecular Biology Department

Ka is a measure of the acid strength. When Ka is large there
is considerable dissociation and the acid is strong. When Ka
is small there is a small degree of dissociation, and the acid
is weak.
Value of Ka like [H], involve negative power of ten, and they are
expressed as pKa values where:
pKa=- log Ka.
The stronger the acid the smaller is its pKa .
Strong Acids
Effectively
dissociated (100%)
Weak Acids
Partialy
dissociated
Ka is large
Ka is small
Degree of
dissociation
Acid dissociation
constant Ka
pKA
smaller pKA
larger pKA
Examples
HCl and H2SO4
Organic acids e.g.
H2CO3
The Henderson-Hasselbalch Equation:
It is an important equation to determine the pH of the weak acids
as they are incompletely dissociated: It is used to prepare buffers
While for strong acids (completely
dissociated) it is simply
pH=-log[H+]
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Medical Biochemistry and Molecular Biology Department
Problem Example (1)
Please given the following pH values: 9, 4.5, 2, 11, 3.8.
a)
Arrange
the
pH
values
in
terms
of
increasing
acid
strength. (Remember the lowest acid strength is the highest numbers). So
"11" would be the least acid strength. You can turn the pH values to
concentration easily. The concentrations are 10-9, 10-4.5, 10-2, etc. Note
that 10-2 is 1/100 [hundredth], which is a larger number than 10 -9, which is
1/1000000000 [a billionth])
b) Arrange the pH values in terms of increasing base strength. (Base
strength is reverse of acid strength, so the order in the previous problem is
reversed)
Problem Example (2)
Please determine the pH of a 0.010 M solution of Ba(OH)2:
………………………………………………………………………………
Comparison between an acid and an base
Definition
ACID
BASES
A substance that
liberates hydrogen ions
into solution Proton
Donor
A substance that yields
an excess of hydroxide
ions when dissolved in
water. Proton Acceptor
Taste
A characteristic sour
taste (think of lemon
juice!)
Reaction with litmus paper
Red
Examples
HCl and HNO3 ,acetic
acid
A bitter taste
Blue
Na OH, ammonia
Ability to react with each other to form salts.
React with
certain metals to
produce gaseous H2
.
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Ability to react with
acids to form salts.
Medical Biochemistry and Molecular Biology Department
MEASURING THE pH

There are substances which have the property of changing their color
when they come in contact with an acidic or basic environment. These
substances are called pH indicators. Usually, these substances are
weak acids which have a color in the non dissociated form and
another in the dissociated form.

They might be in the form of solutions e.g.(Phenol red) or litmus
paper

Litmus paper: special papers which have been soaked with indicators
are used.

These papers change color when they are
Red
blue
immersed in acidic or basic liquids. This is the case of the well-known
litmus paper.

More recently, it has become possible to measure the pH with electrical
instruments like the pH meter.
 The pH meter is an electronic instrument supplied with a special bulb
which is sensitive to the hydrogen ions which are present in the
solution being tested. These instruments are much more precise
and convenient to use than the indicating papers
Bulb of pH
meter
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Medical Biochemistry and Molecular Biology Department
How simple indicators work
Indicator(I) is a weak acid. It has a seriously complicated
molecule which we will simplify to HI The "H" is the proton which
can be given away to something else. The "I" is the rest of the
weak acid molecule. There will be an equilibrium established
when this acid dissolves in water. Taking the simplified version
of this equilibrium:
Adding hydroxide ions:
Blue
Adding hydrogen ions:
red
If the concentrations of HI and I are equal: At some point
during the movement of the position of equilibrium, the
concentrations of the two colors will become equal. The color
you see will be a mixture of the two.
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Medical Biochemistry and Molecular Biology Department
Table - Acid-Base Indicators
INDICATOR
COLOR CHANGE
INTERVAL (PH)
ACID
BASE
Thymol Blue
1.2 - 2.8
red
yellow
Methyl Orange
3.1 - 4.4
red
yellow
Methyl Red
4.4 - 6.2
red
yellow
Chlorophenol Red
5.4 - 6.8
yellow
red
Bromothymol Blue
6.2 -7.6
yellow
blue
Phenol Red
6.4 - 8.0
yellow
red
Thymol Blue
8.0 - 9.6
yellow
blue
Phenolphthalein
8.0 - 10.0
colorless
red
More indicators: http://chemistry.about.com/library/weekly/aa112201a.htm
Buffer
Buffers are solutions that contain an acid and its conjugate base that are
designed to resist pH changes.
This is important in biological systems to maintain proper blood chemistry and
in the environment to help minimize effects of acid rain. A buffer works by
using the equilibrium between the acid and its conjugate base to minimize the
effect that the addition of a small amount of additional acid or base has on the
pH.
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Medical Biochemistry and Molecular Biology Department
Principles of Buffering:
Buffers work chemically by shifting the equilibrium between the weak acid and
the conjugate base in the following manner. If a strong acid is added, it uses
up some of the conjugate base and produces more of the weak acid.
The pH may shift slightly downward due to the small amount of dissociation of
the weak acid, but it is much less of a shift than would be seen by adding the
strong acid directly to water, where each mole of acid added adds directly to
the [H3O+].
Since no H3O+ is produced directly by the addition of the strong acid, there is
minimal shift in pH. Buffers offer similar protection against changes in pH due
to the addition of a strong base. In this case, the weak acid absorbs the base
Thus, in a similar manner to the addition of acid, no explicit change in OH- or
H3O+ is seen, so there is minimal change in the pH. The pH is governed only
by the equilibrium constant, Ka, not directly by the addition of base or acid as
would be seen in a water solution.

The pKa of the buffer depends on
o
Concentration: It is the total (formal) concentration of the buffer
species ( [acid] + [conjugate base] )
o
Temperature
o
Ionic strength
 The ideal buffer would be one with a pKa very near the pH of our solution.
When pH =pKa, [HA] = [A-]. This allows the buffer to consume both
acid and base (and, as noted briefly,minimizes the pH change for a
particular addition).
 A rule of thumb is that buffers are useful within about +/- 1 pH unit of
the pKa.
Problem Example 5 (6)
The pKa of acetic acid is 4.76, its buffer mixture is most effective at pH:
a) 2.3
b) 4.5
c) 3.7
Answer: b (Apply The Henderson-Hasselbalch Equation )
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Medical Biochemistry and Molecular Biology Department
Applications of buffers
Buffer solutions are necessary to keep the correct pH for enzymes in many
organisms to work. Many enzymes work only under very precise conditions; if
the pH moves outside of a narrow range, the enzymes slow or stop working
and can denature, thus permanently disabling their catalytic activity which will
be fatal in many conditions (e.g. renal failure most of patients die from acid
bas imbalance ).
Physiological buffers:
a) A buffer of carbonic acid (H2CO3) and bicarbonate (HCO3−) is present in
blood plasma, to maintain a pH between 7.35 and 7.45. (pKa= 6.1)
b) The majority of biological samples that are used in research are made in
buffers, especially phosphate buffered saline (PBS) at pH 7.4. Phosphate
buffer is formed of H2PO4 (the acid) and HPO4 (its conjugate base), it is
found mostly in the intracellular fluids.(pKa= 6.8).So Theoretically,
Bicarbonate buffer is < efficient than Phosphate buffer Because…
However, bicarbonate system is more efficient in our blood because:
a) It is of higher concentration than phosphate buffer (0.002m) “so
phosphate buffer is more important as intracellular buffer only”.
b) Bicarbonate system produces less harmful end products i.e.
NaHCo2 + HCl → NaCl + H2CO3
H2CO3 ↔ Co2 + H2O (by Carbonic anhydrase enzyme)
Lab activities include the following order:
1. To identify a solution is acidic, basic or neutral in the laboratory using litmus
paper.
2. To investigate how indicators can be used to test for the presence of
acids or bases and the effect of pH on Indicator Dyes and to understand the
concept of neutralization reaction as well.
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Medical Biochemistry and Molecular Biology Department
3. Demonstrate the reaction between an acid and a base (home activity)
4. Demonstrate how buffer resist changes in pH.
Experiment 1:
Purpose: To identify a solution is acidic, basic or neutral in the laboratory
using litmus paper.
Materials needed: Test tubes, test tube rack, dropper pipettes, litmus paper,
vinegar (acetic acid) or lemon juice, water or milk, ammonia or backing
powder.
Procedure:
1. Put 2 ml of each of the available materials in different test tubes; (Acetic
acid, tab water, and ammonia).
2. Put a piece of litmus paper in each of
the test tubes of first set.
3. Observe and record the color of litmus
paper in each test tube.
OBSERVATIONS
Conclusions:
………………………………………………………………………..
QUESTION: What are the names of other acid-base indicators and how is the
affect of these indicators to acidic and basic solutions?
Experiment 2:
Purpose: To investigate how indicators can be used to test for the presence
of acids or bases and the effect of pH on them and to understand the concept
of neutralization reaction as well.
Materials needed:
Phenol red indicator, Acetic
acid, ammonia, water, test
13 DW
+PR
+ Alk
Medical Biochemistry and Molecular Biology Department
tubes, test tube rack, dropper pipettes
Procedure:
1-In a container put about 5ml of water and add about 5
drops of phenol red. The phenol red is an indicator which is
orange in color with pH7 just like water
2- With a dropper, add some drops of ammonia to the
previous solution. Ammonia has irritating vapours, to avoid
these vapours, make this experiment in a well ventilated
+acid
+Alk
area, keeping at a distance away. After some drops, you
will see the solution suddenly become reddish purple.
3-Now, to the red-purple fluid you obtained, add some drops of vinegar. What
happens? The liquid becomes yellow. Watch it!
 Explain your results
Conclusions:
………………………………………….………
Experiment 3 (Home activity)
Purpose: To understand the concept of chemical reaction between acid and
basic substances react with each other, producing a salt and often other
substances like water and carbon dioxide.
Materials needed:
test tubes, test tube rack, dropper pipettes, distilled water, vinegar (5% acetic
acid), baking soda (Na bicarbonate- Na HCO3)
Procedure:
In half a glass of water, put a few teaspoons of baking soda and mix in order
to obtain a quite concentrated solution. In the same
glass, pour a spoon of vinegar. As you can see, there
will be an abundant production of foam. What
happened?
OBSERVATIONS
A chemical reaction occurred between the baking soda
(a basic substance) and vinegar (an acid substance).
These two substances reacted with each other,
producing a salt, water, and carbon dioxide. That is the gas which produced
the little bubbles you observed and the salt produced is sodium acetate:
CH3COOH + NaHCO3 = CH3COONa + H2O + CO2
Acetic acid Baking soda Sodium acetate Water Carbon dioxide
Conclusions:
…………………………………………………………………………………………
………………………………………………………………
Experiment 4:
Purpose: To demonstrate how buffer can resist changes in pH.
Materials needed:
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Medical Biochemistry and Molecular Biology Department
Small beakers, dropper pipettes, distilled water, prepared phosphate buffer of
PH7and phenol red indicator (range 6.4-8), Ammonia or Acetic acid
Preparation:






Add 20 ml of DW to the first beaker and buffer to the second beaker.
Add 5 drops of phenol red to each beaker.
Mix well and record your observation.(the intermediate color)
Then add one drop of alkali (e.g ammonia)or acid( e.g. acetic acid) to
the two beakers, and watch the difference, then record your
observation.(the intermediate color)
Continue to add more drops to both beakers and watch what happens.
then record your observation
Explain your results.
Conclusions:
…………………………………………………………………………………………
………………………………………………………………
Problems to solve:
1)What is the pH of the solution with a hydronium concentration
[H3O+] 1.47 x 10-4?
What is the pOH of this solution?
2) What is the pOH of the solution with a hydroxyl concentration
[OH-] 2.98 x 10-2? What is the hydronium concentration [H3O+] of this
solution?
3) What is the hydronium concentration [H3O+] of a solution with a pH
of 7.84? What is the hydroxyl concentration [OH-] of this solution?
4) What is the hydroxyl concentration [OH-] of a solution with a pH of
3.76?
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