CONTENTS
CONTENTS
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PAGE
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Abstract
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2
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Introduction
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3
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Literature review
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5
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Objectives
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7
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Methodology
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8
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Results
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9
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Discussion
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10
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Conclusion and recommendation
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15
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Reference
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16
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Appendix
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17
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This purpose of this experiment was to understand the nature of a
buffer, to prepare a buffer from acetic acid and sodium acetate and also to
test the ability of buffered and unbuffered solutions to resist pH changes when
strong acids and bases are added. This experiment can be divided into three
parts, where the first part was to prepare the buffer solutions needed, the
second part tested the buffering action towards acid while the third part
tested the buffering action towards base. In parts B and C, the acid and base were
added 1ml once, drop by drop using a dropper so that they dissolved well in the
solutions. The pH readings of the solution before and after adding each ml of
the acid or base were measured by a pH meter which was calibrated to ensure accurate
reading. During this experiment, several errors which can affect the results of
the experiment had occurred. For example, the beakers used to contain the
buffer solutions might be contaminated. This caused the pH reading taken to
differ from the theoretical values. Besides, the calibration of the pH meter
was not done well enough to acquire accurate readings for the measurements.
This explained why the pH readings of the solutions measured are all lower than
expected.
A buffer solution is one which resists changes in pH when
small quantities of an acid or an alkali are added to it.The buffer solution
also known as an aqueous
solution consisting of a mixture of a weak
acid and its conjugate
base or a weak
base and its conjugate
acid. It has the property that the pH of the solution changes very little when a small amount of strong
acid or base is added to it. Buffer solutions are used as a means of keeping pH
at a nearly constant value in a wide variety of chemical applications. Many
life forms thrive only in a relatively small pH range; an example of a buffer
solution is blood.
Buffer
solutions achieve their resistance to pH change because of the presence of an
equilibrium between the acid HA and its conjugate base A-.
HA 
H+ + A-
H+ + A-
When
some strong acid is added to an equilibrium
mixture of the weak acid and its conjugate base, the equilibrium is shifted to the left, in
accordance with Le Chatelier's principle. Because of this, the hydrogen ion
concentration increases by less than the amount expected for the quantity of
strong acid added. Similarly, if strong alkali is added to the mixture the
hydrogen ion concentration decreases by less than the amount expected for the
quantity of alkali added.
The effect
is illustrated by the simulated titration of a weak acid with pKa =
4.7. The relative concentration of undissociated acid is shown in blue and of
its conjugate base in red. The pH changes relatively slowly in the buffer
region, pH = pKa ± 1, centered at pH = 4.7 where [HA] = [A-],
but once the acid is more than 95% deprotonated the pH rises much more rapidly.
An acidic buffer solution
is simply one which has a pH less than 7. Acidic buffer solutions are commonly
made from a weak acid and one of its salts - often a sodium salt.
A common example would be a
mixture of ethanoic acid and sodium ethanoate in solution. In this case, if the
solution contained equal molar concentrations of both the acid and the salt, it
would have a pH of 4.76. It wouldn't matter what the concentrations were, as
long as they were the same.
An alkaline buffer solution
has a pH greater than 7. Alkaline buffer solutions are commonly made from a
weak base and one of its salts.
A frequently used example
is a mixture of ammonia solution and ammonium chloride solution. If these were
mixed in equal molar proportions, the solution would have a pH of 9.25. Again,
it doesn't matter what concentrations you choose as long as they are the same.
By adding an acid to buffer
solution,the buffer solution must remove most of the new hydrogen ions
otherwise the pH would drop markedly.
Hydrogen ions combine with
the ethanoate ions to make ethanoic acid. Although the reaction is reversible,
since the ethanoic acid is a weak acid, most of the new hydrogen ions are
removed in this way.
Since most of the new
hydrogen ions are removed, the pH won't change very much - but because of the
equilibria involved, it will fall a little bit.
By adding an alkali to
buffer solution,the alkaline solutions contain hydroxide ions and the buffer
solution removes most of these.
This time the situation is
a bit more complicated because there are two processes which
can remove hydroxide ions.
Removal by reacting with
ethanoic acid
The most likely acidic
substance which a hydroxide ion is going to collide with is an ethanoic acid
molecule. They will react to form ethanoate ions and water.
Because most of the new
hydroxide ions are removed, the pH doesn't increase very much.
Removal of the hydroxide
ions by reacting with hydrogen ions
Remember that there are
some hydrogen ions present from the ionisation of the ethanoic acid.
Hydroxide ions can combine
with these to make water. As soon as this happens, the equilibrium tips to
replace them. This keeps on happening until most of the hydroxide ions are
removed.
LITERATURE REVIEW
Buffer solutions achieve their
resistance to pH change because of the presence of an equilibrium between the
acid HA and its conjugate base A-.
HA H+ + A-
H+ + A-
When some strong acid is
added to an equilibrium mixture of the weak acid and its conjugate base,
the equilibrium is shifted to the left, in accordance with Le Chatelier's
principle. Because of this, the hydrogen
ion concentration increases by less than the amount expected for the quantity of
strong acid added. Similarly, if strong alkali is added to the mixture the
hydrogen ion concentration decreases by less than the amount expected for the
quantity of alkali added.
The effect is illustrated by the
simulated titration of a weak acid with pKa = 4.7. The relative
concentration of undissociated acid is shown in blue and of its conjugate base
in red. The pH changes relatively slowly in the buffer region, pH = pKa ±
1, centered at pH = 4.7 where [HA] = [A-], but once the acid is more
than 95% deprotonated the pH rises much more rapidly.
A. Acidic buffer solution
An
acidic buffer solution is simply one which has a pH less than 7. Acidic buffer
solutions are commonly made from a weak acid and one of its salts - often a
sodium salt.
A common example would be a mixture
of ethanoic acid and sodium ethanoate in solution. In this case, if the
solution contained equal molar concentrations of both the acid and the salt,
it would have a pH of 4.76. It wouldn't matter what the concentrations were,
as long as they were the same.
|
Example:
 |
You
can change the pH of the buffer solution by changing the ratio of acid to salt,
or by choosing a different acid and one of its salts.
B. Alkaline buffer solutions
An
alkaline buffer solution has a pH greater than 7. Alkaline buffer solutions are
commonly made from a weak base and one of its salts.
A buffer solution has to contain things
which will remove any hydrogen ions or hydroxide ions that you might add to
it - otherwise the pH will change. Acidic and alkaline buffer solutions
achieve this in different ways.
|
A
frequently used example is a mixture of ammonia solution and ammonium chloride
solution. If these were mixed in equal molar proportions, the solution would
have a pH of 9.25. Again, it doesn't matter what concentrations you choose as
long as they are the same.
Example
:
OBJECTIVES
1. To understand the nature or a buffer.
In this experiment, the initial and final pH
readings of the buffer solutions were measured, to study how the addition of
acid or base affect the pH of the buffer solutions.
2. To prepare a buffer from acetic acid and
sodium acetate.
This acidic buffer was prepared in this
experiment with different amount of sodium acetate dissolved by acetic acid.
3. To test the ability of buffer and unbuffered
solutions to resist pH changes when strong acids and bases are added.
Bufferd and unbuffered solutions prepared were
added with slowly increased amount of strong acids and bases, and the pH
readings were taken each time the acids and bases were added. With the recorded
readings, it can be studied the pattern of pH change in each different
solutions. Then, the stage at which pH changes were resisted can be determined.
METHODOLOGY
MATERIALS
Beakers, distilled water, sodium
choride, solid sodium acetate, acetic acid, stirring rod, hydrochloric acid, pipette,
pH meter, magnetic stirrer, spin bar, sodium hydroxide.
PROCEDURES
A.
The Preparation of the Buffer Solution
1.
7 100 ml beakers were labelled as 1 to 7.
2.
50 mL distilled water was added into beaker 1
and 6 and 50 mL 0.1 M sodium chloride into beaker 2 and 7.
3.
1 g solid sodium acetate (CH3COONa)
was weigh and transferred into beaker 3.
4.
5 g solid sodium acetate (CH3COONa)
was weigh and transferred into beaker 4.
5.
10 g solid sodium acetate (CH3COONa)
was weigh and transferred into beaker 5.
6.
50 mL 0.1 M acetic acid was added to beaker 3,
4, and 5. The solution was stir until all the solid dissolve.
B.
The Determination of Buffering Action toward
Acid
1.
The pH meter was calibrated.
2.
The pH value of distilled water in beaker 1
was determined and recorded.
3.
1 mL 6.0 M hydrochloric acid was added into
beaker 1. Mixed the solution and the new pH value of solution was determined.
4.
Step 3 was repeated until there is only a
slight change in pH value.
5.
Step 2 to step 4 were repeated for each beaker
2, 3, 4 and 5
C.
The Determination of Buffering Action Toward
Base
1.
All the steps in Section B were repeated by
replacing 6.0 M hydrochloric acid with 6.0 M sodium hydroxide as well as beaker
1-5 with beaker 6-7.
RESULTS
Beaker
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Content
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Initial pH
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pH when the stated amount of 6.0M hydrochloric acid is added
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Slight change in pH
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pHwhen the stated amount of6.0 M sodium hydroxide is added
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Slight change in pH
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||||
1 ml
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2ml
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3ml
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1ml
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2ml
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3ml
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|||||
1
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Distilled Water
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4.89
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1.1.
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0.55
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0.54
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0.01
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||||
2
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0.1 M Sodium Chloride
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2.91
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1.42
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0.73
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0.58
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0.15
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||||
3
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CH3COONa(1g) + 50ml 0.1 M
acetic acid
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3.00
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2.95
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-
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-
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0.05
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||||
4
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CH3COONa(5g) + 50ml 0.1 M
acetic acid
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4.09
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4.06
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-
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-
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0.03
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||||
5
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CH3COONa(10g) + 50ml 0.1
M acetic acid
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4.65
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4.64
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-
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-
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0.01
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||||
6
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Distilled Water
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4.67
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12.47
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12.72
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-
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0.25
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||||
7
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0.1 M Sodium Chloride
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4.34
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11.45
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13.74
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13.80
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0.06
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||||
DISCUSSION
The 0.1 M sodium chloride needed to
prepare the solutions in part A was prepared from solid sodium chloride. The
weight of solid sodium chloride needed was calculated as below.
0.1
M NaCl =
0.1 mol dm-3 Na Cl
=
0.1 (23 + 35.5)g dm-3 NaCl
= 5.85 g
dm-3 NaCl
1dm3 is equal to 1 L. So, 5.85 g of
solid NaCl was weighed and added into distilled water in a
1 L volumetric flask. Distilled water was then
added until the calibration mark. The volumetric flask was shaken vigorously to
produce evenly mixed 0.1 M NaCl.
The 0.1 M acetic acid was prepared
from raw liquid acetic acid using the same method. However, raw liquid acetic
acid was used instead of solid acetic acid because there would be a very
exothermic reaction when acid reacts with water.
M=0.1M
V=1000mL
Mw=60.05 g/mole
Purity= 1.05%
Pg= 2.1 g/cm3
mconcentrate=10.32 g
Vconcentrate= 10
mL
M1V1 = M2V2
M2 = M1V1/V2
= 2.1 x 1.05 x 1000/ 60.05
= 36.72
=36.7
V1 = 0.1 x 250/36.7
=0.68ml
Beaker 1
The initial pH reading of the
distilled water was 4.89, which showed a slightly acidic pH value. Distilled
water is supposed to be neutral, which means it has a pH value of 7. The lower
pH in the result might be caused by impurities in the distilled water, or
contamination of the beaker or other apparatus used.
After adding in 1ml of 6.0 M
hydrochloric acid, the pH of the solution dropped to 1.10, which is very
acidic. This is because the hydrochloric acid dissociated in the distilled
water to form H+ ions and Cl- ions. The H+
ions contributed to the acidity of the solution.
As the concentration of hydrogen ion
increased,
pH = -log [H+]
[H+] increases,
-log [H+] decreases.
As a result, pH decreases.
The pH of the solution decreased
greatly until the numbers of moles of other particles were insignificant. The
molarity of hydrogen ion remained almost constant. Therefore, the pH dropped
very slightly.
Distilled water does not have buffer
property because it cannot resist the pH change when the hydrochloric acid was
added.
Beaker 2
The pH reading of the solution was
lower than the theoretical value, due to the same reason as in Beaker 1.
The pH change in this solution was
similar to that in Beaker 1, meaning that the solution in Beaker 2 had very low
buffer capacity. It did not act as a buffer to resist the pH change when strong
acid (hydrochloric acid)was added into it.
Beaker 3, 4 and 5
In these beakers, the following
chemical equations took place.
CH3COOH (aq) CH3
CH3COONa (s) CH3COO-
(aq) + Na+ (aq)
The acetic acid was only partially
dissociated in water but the sodium acetate was fully dissociated. Thus, there
were low concentration of H+ ions and high concentrations of CH3COO-
ions and the undissociated CH3COOH molecules. When hydrochloric acid
was added into it, the H+ ions dissociated from the hydrochloric
acid combined with the CH3COO- ions to form molecules of
acetic acid. Since that most of the hydrogen ions, H+ was removed,
the concentration of hydrogen ions remained almost the same, and as a result,
the pH remained almost unchanged. As observed in these beakers, the pH changes
when hydrochloric acid was added were very slight.
Beaker 6 and 7
Beaker 6 had the same content as
Beaker 1, thus the initial pH readings of the two solutions were almost the
same. Beaker 7 had the same content as Beaker 2, which was 0.1 M sodium
chloride. However, the initial pH readings of the two solutions were
significantly different. This was caused by impurities present, affecting the
pH of the solution.
In these two beakers, when sodium
hydroxide was added into the two solutions, their pH changed greatly. This
showed that the solutions had very low buffer property. The two solutions
cannot act as effective buffer solutions.
In comparison, the change in pH
reading was smaller in Beaker 7 than in Beaker 6 when sodium hydroxide was
added. This showed that the solution in Beaker 7 had a greater buffer capacity
than the solution in Beaker 6.
1. From your results in this experiment, which
solution of those you tested had the greatest buffer capacity:
Toward strong acid?
Toward strong base?
Discuss.
Buffer capacity is the amount of acid or base which can be
absorbed by the buffer solution without a significant change in pH. The greater
the buffer capacity, the more acid or base can be added into it without
resulting in significant pH change. In this experiment, the smaller the pH
change when the acid or base was added, the greater the buffer capacity.
Hydrochloric acid is a strong acid because it dissociates
completely in water to produce hydrogen ions.
From the results in this experiment, the tested solution which had
the greatest buffer capacity towards strong acid is the solution in Beaker 5, 50 ml 0.1 M acetic acid with
10 g solid sodium acetate.
This is shown through its slightest change in pH when 6.0 M
hydrochloric acid was added into it.
Sodium hydroxide is a strong base because it dissociates
completely in water to produce hydroxide ions.
The tested solution which had the greatest buffer capacity towards
strong base is the solution in Beaker 7,
50 ml 0.1 M sodium chloride.
This is shown through its slighter change in pH when sodium
hydroxide was added into it, as compared to that in solution in Beaker 6.
In comparing the buffer solutions in Beaker 3 to 5, they are the
mixture of the same substances, but different concentrations. This makes it
clear that buffer capacity depends on the concentration of the weak acid and
its conjugate base, which is CH2COO-. In general, the
maximum buffer capacity exists when the concentrations of the weak acid and its
conjugate base are approximately the same, and in large amount.
2. Why was distilled water used to rinse off
the pH probe?
Distilled water
was used to rinse the pH probe to make sure that the pH probe was neutral
before measuring the pH value. Distilled water is neutral. This is to avoid the
influence of the measurement of pH due to the hydrogen or hydroxide ions
present on the pH probe.
3. Define buffer solution.
A buffer solution is a solution whose pH
value changes only very slightly when small amounts of an acid or alkali are
added.
4. Specify which of these systems can be
classified as a buffer system:
a. KCl/HCl b. NH3/NH4NO3 c. NaHPO4/NaH2PO4
KCL/HCL is not a buffer system because HCl
is not a weak acid. Buffer solutions can only be formed from weak acid or base
so that the acid or base dissociates partially, and a reversible reaction is
present.
NH3/NH4NO3 is a buffer
system because NH3 is a weak base and NH4NO3 is its conjugate acid. Therefore, the combination of these two
chemicals will form buffer solution.
In aqueous state, NH3 reacts with H+ in the
water to form NH4+.
NH4+ NH3 + H+
NH4NO3 NH4+ +
NO3-
The above reactions take place in the solution and maintain the
concentration of hydrogen ion or hydroxide ion when acid or base is added into
it.
NaHPO4/NaH2PO4is
a buffer solution with the following equations.
NaH2PO4 Na+ +
H+ +HPO42-
NaHPO4- Na+ + HPO42-
NaH2PO4is the weak
acid in this buffer system while NaHPO4-
dissociates completely to provideHPO42-, its conjugate
base.
5.
Calculate
the pH of the buffer system 0.15M NH3 / 0.35 M NH4C
Kb for NH3 = 1.8 x 10-5
mol dm-3
By using the Henderson-Hasselbalch equation,
pOH =
- lgKb + lg
= - lg 1.8 x 10-5 + lg
= 5.11
pH = 14.00 - 5.11= 8.89
CONCLUSION AND RECOMMENDATIONS
Based on the results of the experiment, the solutions in Beakers 1,2,6
and 7 are unbuffered solutions while the solutions in Beakers 3,4 and 5are
buffered solutions. The buffer solutions are able to resists pH changes when titrate
with strong acid or strong base. Beaker 4 has the greatest buffer capacity towards
strong acid while beaker 7 has the greatest buffer capacity towards strong
base.
During the experiment, the pH meter must be rinse using distilled
water and the solution inside the pH meter to make sure it is neutral before
measure the pH value. This is to make sure that accurate pH measurement can be
obtained.
To ensure accurate pH measurement, the apparatus and materials
must also be clean from impurities.
The calibration of the pH meter must be done properly until
accurate readings can be obtained.
REFERENCES
- Chang, Raymond (2007). Chemistry.3rd ed. New York: McGraw-Hill
- Umland and Bellama (1999). General Chemistry. 3rd ed. Pacific Grove, CA: Brooks/Cole Publishing Company
- Zumdahl, Steven S (2005). Chemical Principals. 5th ed. New York: Houghton Mifflin Company
- Tan Yin Toon, Physical Chemistry for STPM, (2004), Fajar Bakti.
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