Experiment 1 Preparation of Buffer Solutions

1EXPERIMENT 1: PREPARATION OF BUFFER SOLUTIONS

TABLE OF CONTENTS

Bil. Title Page1 Summary 22 Pictorial Methodology 33 Results 54 Discussions 65 Tutorial 86 Calculations 117 Conclusions 158 References 16


1.0 SUMMARY

In this experiment, the objective is to prepare buffer solutions and test the effects of

added acid or base upon the pH. It is also to review acid-base chemistry and calculation. The

experiment is started by preparing the buffer solutions citric acid and sodium citrate. Weight out

each substances (3.560g of citric acid and 9.255g of sodium citrate) and add distilled water to

make a buffer solution. Then, determine the pH of the solution using the pH electrode. The pH

value calculated is 5.00 compared with the experiment, the pH value obtained by the pH

electrode is 4.96. This slight different occur may come from several error that happened while

handling the solution or while conducted the pH electrode. As conclusions, buffer solutions are

solutions that resist change in hydronium ion and the hydroxide ion concentration upon addition

of small amounts of either acid or base, or upon dilution.


2.0 PICTORIAL METHODOLOGY

2.1 Preparation of buffer solution

2.1.1 Solid citric acid and sodium acetate have been mixed together to form a solution.

Then the solution has been stirred using mechanical stirrer.

2.1.2 500 mL of a 0.1 buffer solution with a pH of 5.0 using citric acid, C6H8O7 and sodium

acetate, NaC2H3O2 were prepared. The mixed solution of citric acid and sodium

acetate then dilute with distilled water in the volumetric flask.


2.2 Testing of Buffer Solution

2.2.1 The buffer solution was added into 100ml beaker and the pH was measured using the

pH electrode.

2.2.2 50ml deionized water was added into two different beakers and the pH was measured.

The first beaker was added with 1.0 ml of 1M HCL and the second beaker was added

with 1.0 ml of 1M NaOH. The pH of each solution was measured.

2.2.3 The above steps was repeated but by using buffer solution instead of deionized water.

The pH of each solution was measured.


3.0 RESULT

Solution pH

Initial After adding HCl After adding NaOH

Deionised water 5.76 2.32 12.09

Buffer 4.96 4.84 5.06


4.0 DISCUSSION

The pH value of the deionised water was easily influenced by the present of atmosphere

CO2. Theoretically, the pH value of deionised water is 5.5. The pH value might change if the

deionised water was left exposed to the atmosphere and this will cause the change in pH value

which will rise up to 7.00. Based on the experiment, the initial pH value for the deionised water

obtained was 5.76. This shows that the deionised water has been well preserved.

When the strong acid, HCl was added into the deionised water, the pH value obtained is

2.32 which mean the deionised water has turn into acid. The same phenomenon occurs while

strong base, NaOH was added into the deionised water. The pH value obtained is 12.09. The

result can be obviously seen, as the deionised water was turned into acid and base as the strong

acid and strong base was added respectively.

The Buffer solution was prepared by mixing the buffer agent’s form (HA) which is citrate

acid with its conjugate base (A-) which is sodium citrate. The pH value calculated is 5.00

compared with the experiment, the pH value obtained by the pH electrode is 4.96. This slight

different occur may come from several error that happened while handling the solution or while

conducted the pH electrode.

As HCl was added into the Buffer solution, the pH value changed from 4.96 to 4.84. It

shows that the difference occur is quite small which is only 0.12. At the same time while NaOH

was added, the pH value changed from 4.96 to 5.06. The buffer solution is still resist the pH

value from change, as the difference is only 0.1. Although the pH value of Buffer solution

obtained is different from the calculated, the activity of the Buffer solution is still can obviously

be seen. The change of pH value is only in the range of 0.1-0.12.

The error occurred gives a slight different as the calculated. The prepared Buffer solution

may have a little contamination causing the pH value to become 4.96 (which is supposed to be

5.00). Second possible error may occur while handling the Buffer solution with the pH electrode.

The data was taken before it is stable or the pH electrode itself was not well cleaned. Another

error might happened was during the chemical was taken. The chemical taken might not be

precise as desired (parallax error).


As the recommendation, before using any apparatus, it is necessary to clean it up first.

This is to ensure that the apparatus is free contamination. The cleaning processes have to take

charge while handling the pH electrode as well. This is a must in order to obtain a desirable pH

value. As for the deionised water, try not to leave it exposed to the air in a long period of time.


5.0 TUTORIAL

1. Describe how you would prepare a (CH3CO2H / CH3CO2-) buffer with a pH of 5.2 given that the pKa (CH3CO2H / CH3CO2-) is 4.76.

-Find the mole ratio of CH3CO2H / CH3CO2- using Henderson-Hasselbalch equation. The calculation as follows:

Henderson-Hasselbalch equation:

pH= pKa + log[CH3 CO2- ][ CH3 CO2 H ]

Rearrange the equation:

log[CH 3 CO2- ][ CH 3 CO2 H]

=pH -pKa (1)

Given the pKa and pH for citric acid:

pKa= 4.76 (2)

pH = 5.2 (3)

Substitute (2) and (3) into (1)

log[CH 3 CO2- ][ CH 3 CO2 H]

= 5.2 - 4.76

log[CH 3CO2- ][ CH 3 CO2H]

= 0.44 (4)

Antilog equation (4)

[CH 3CO 2- ][ CH3 CO2 H ]

= 100.44

[CH 3CO 2- ][ CH3CO2 H ]

= 2.754

So the mole ratio of citric acid to citrate ion is

[CH 3CO 2- ][ CH3CO2 H ]

= 2.754 (5)


Given:

n= MV V= 500 ml = 0.5 L (6)

nV

= M (7), then Expand equation (7) into

nCH 3CO 2-+n CH 3CO 2 H

V= M (8)

Rearrange equation (5)

nCH 3 CO2-=2.754n CH 3 CO2H (9)Substititute (9) into (8)

nCH 3CO 2-+n CH 3CO 2 H

0.5L= 0.1

molL

(10)

Solve equation (10)

3.754n CH 3 CO2H=0.5mol

n CH 3 CO2 H=0.5mol3.754

n CH 3 CO2 H=0.1332 mol

Using equation (9) to solve no of mole sodium citrate

nCH 3CO2-=2.754n CH 3 CO2 H

nCH 3CO2-=2.754 (0.1332 mol )

nCH 3CO2- = 0.3668 mol

Molar mass for CH3 CO2H = 60.05 g/mol

Molar mass for CH3 CO2- = 82.03 g/mol

Mass of citric acid = 0.1332 mol x 60.05 g/mol = 7.999g

Mass of sodium citrate = 0.3668 mol x 82.03 g/mol = 30.089 g

- Weight both of the citric acid and sodium acetate and put it inside the 500 ml volumetric flask.

-Then add distilled water into the volumetric flask till neck of flask. Dissolve the acid and salt.


- Add the distilled water till calibration mark.

2. What is buffer capacity? What determines the capacity of a buffer?

Buffer capacity is the ability of buffer solution to resist pH changes. Buffer capacity increases as the molar concentration of acid or salt increases. The buffer capacity also increases as the pH is closer to pKa. The buffer capacity is expressed as the molar concentration of sodium hydroxide required to increase the pH by 1.


6.0 CALCULATIONS

1. Calculation for 500 ml of 0.1M of buffer solution with pH 5

Calculation of buffer solution can be done by using Henderson-Hasselbalch equation to deter mine the mole ratio of the citric ion and citrate ion. The calculation as follows:

Henderson-Hasselbalch equation:

pH=pK a + log[ A - ]HA

rearrange the equation:

log [A - ]HA

=pH -pKa (1)

Given the pKa and pH for citric acid:

pKa= 4.77 (2)

pH = 5 (3)

Substitute (2) and (3) into (1)

log [A - ]HA

= 5-4.77

log [A - ]HA

= 0.23 (4)

Antilog equation (4)

[ A- ]HA

= 100.23

[ A- ]HA

= 1.698

So the mole ratio of citric acid to citrate ion is

[ A- ]HA

= 1.698 (5)

Given:

n= MV (6)


V= 500 ml = 0.5 L

nV

= M (7)

Expand equation (7) into

n A−¿+nHA

V= M ¿ (8)

Rearrange equation (5)

nA−¿=1.698 nHA ¿ (9)Substititute (9) into (8)1.698nHA+nHA0.5L

= 0.1 molL

(10)

Solve equation (10)

2.698 nHA=0.5mol

nHA=0.5mol2.698

nHA=0.0 1853 mol

Using equation (9) to solve no of mole sodium citrate

nA−¿=1.698 (0.01853)¿

nA−¿=0.03147 mol ¿ Molar mass for citric acid = 192.13 g/mol

Molar mass for sodium citrate = 294.1 g/mol

Mass of citric acid = 0.01853 mol x 192.13 g/mol

*Mass of citric acid = 3.560 g

Mass of sodium citrate = 0 .03147 mol x 294.1 g/mol

*Mass of sodium citrate = 9.255 g


2. For pH of buffer after adding acid

The equation of ionization of citric acid:

C6 H8O7 ↔C6 H5 O7- +3 H+¿ ¿

¿¿

“ICE” chart: Assume all acid consume.

Mol C6H8O7 H+ C6H5O7-

Initial 0.01853 mol 0.001 L(1 M) = 0.001 mol3 x 0.001 mol=0.003

0.03147 mol

Change +0.003mol -0.003 mol -0.003 molEquilibrium 0.02153 mol x 0.02847 mol

Find Ka

Ka = 10-4.77 =1.698 x 10−5

¿¿

¿¿

Convert to pH

pH= - log10¿¿

pH= - log10(1.284 × 10-5 )

pH= 4.90

(Refer to http://www.chem.purdue.edu/gchelp/howtosolveit/Equilibrium/Buffers.htm#Bufferaddacid)


3. For pH of buffer after adding base

Ionization equation of sodium citrate:

C6 H8O7 → H20+C6 H5O7-

Assume all base is consumed.

Mol C6H8O7 OH- C6H5O7-

Initial 0.01853 mol 0.001 L(1 M) = 0.001 mol

0.03147 mol

Change +0.001mol -0.001 mol +0.001 molEquilibrium 0.01953 mol 0 0.03247 mol

C6 H8O7 ↔C6 H5 O7- +3 H+¿ ¿

C6H8O7 H+ C6H5O7-

Initial 0.01853 0 0.03147 Change -x +x -xEquilibrium 0.01953-x x 0.03247-x

x=(1.698 x 10-5 )0.01953-x 0.03247-x

x=1.021x 10 -5

pH= - log10(1.021x 10-5 )

pH= 4.99


7.0 CONCLUSIONS

As for the conclusions, buffer solutions are solutions that resist change in hydronium ion

and the hydroxide ion concentration (and consequently pH) upon addition of small amounts

of either acid or base, or upon dilution. A buffer is a mixture of a weak acid and its

conjugate base. Because the solute can either absorb protons or release protons, the buffer

solution can maintain the pH at a constant value.

From the results obtained, it can be concluded that the initial pH value for the deionised

water obtained was 5.76. This shows that the deionised water has been well preserved.When the

strong acid, HCl was added into the deionised water, the pH value obtained is 2.32 which mean

the deionised water has turn into acid. The same phenomenon occurs while strong base, NaOH

was added into the deionised water. The pH value obtained is 12.09. The result can be obviously

seen, as the deionised water was turned into acid and base as the strong acid and strong base was

added respectively.

As HCl was added into the Buffer solution, the pH value changed from 4.96 to 4.84. It

shows that the difference occur is quite small which is only 0.12. At the same time while NaOH

was added, the pH value changed from 4.96 to 5.06. The buffer solution is still resist the pH

value from change, as the difference is only 0.1. Although the pH value of Buffer solution

obtained is different from the calculated, the activity of the Buffer solution is still can obviously

be seen. The change of pH value is only in the range of 0.1-0.12.

Due to some errors that occur, the experimental data values obtained are opposed to the

theoretical value. Therefore, some recommendations have been made so that the errors can be

overcome. Lastly, the two main objectives of this experiment are fully achieved.


8.0 REFERENCES

1. (Deionized Water With Acid and Base [online] Available from

http://www.finishing.com/2400-2599/2576.shtml [31st August 2012])

2. (Buffer Solutions [online] Available from

http://chemed.chem.wisc.edu/chempaths/GenChem-Textbook/Buffer-solutios-620.html

[31st August 2012])

3. (Buffer Solutions [online] Available

fromhttp://en.wikipedia.org/wiki/Buffer_solution[31st August 2012])

4. (Buffer Solutions [online] Available from

www.chemguide.co.uk/physical/acidbaseeqia/buffers.html[31st August 2012])

5. (Preparation of pH Buffer Solutions [online] Available from

http://delloyd.50megs.com/moreinfo/buffers2.html [31st August 2012])

Experiment 1 Preparation of Buffer Solutions

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