Study of the pH Change in the Titration of [M/10] HCl with [M/10] NaOH using a Universal Indicator

pH change is indicated with the help of a universal indicator. And a universal indicator is defined as a pH indicator made of a solution of several compounds that exhibits several smooth colour changes over a wide range of pH values to indicate the acidity or alkalinity of solutions. There are some natural indicators also which are found in nature and those indicators are-red cabbage, grape juice, turmeric, curry powder, cherries, beetroots, turnip skin, tomato, onion etc. Some examples of synthetic indicators are methyl orange, phenolphthalein etc.

Table of Content

• Aim

• Theory

• Apparatus Required

• Procedure

• Observation

• Result

• Precautions

• Lab Manual Questions

• Viva Questions

• Practical Based Questions

Aim

To study the pH change in the titration of \[\dfrac{M}{{10}}\] HCl with \[\dfrac{M}{{10}}\] NaOH using a universal indicator.

Materials Required

1. Burette

2. Pipette (20 ml)

3. Titration flask

4. Beakers

5. Funnel

6. Universal indicator solution

7. 0.1 M HCl

8. 0.1 M NaOH

Theory

1. Titration of HCl with NaOH 

During titration, the titrant (NaOH) is added slowly to the unknown solution. As it is added, the HCl slowly reacts away. The point at which exactly enough titrant (NaOH) has been added to react with all the analytes (HCl) is called the equivalence point.

2. Universal Indicators 

A universal indicator is a mixture of different types of indicators that exhibits different colouration at different levels. It can be in the form of a paper strip or a solution. Examples are Methyl red and Phenolphthalein.

Procedure

1. Take a clean (well-washed) burette and rinse (fill) it with 0.1 M HCl solution and then fill (pour) it with this solution.

2. Rinse (fill) the pipette with 0.1 M NaOH (basic) solution. Pipette out 20.0 mL of 0.1 M NaOH in the conical flask and then add about 10 drops of the universal indicator solution to it.

3. Swirl (stir) the solution until the colour of the solution (formed) becomes uniform. Compare the colour of the solution (formed) with the ‘pH Indicator Chart’ and then estimate (calculate) the pH of the solution.

4. Now add (mix) 0.1 M HCl from the burette (instrument) to the solution slowly. After the addition (mixing) of 1 mL solution, compare the colour of the solution (formed) with the ‘pH Indicator Chart’ and estimate (find) the pH of the solution.

5. Keep on adding (mixing) 0.1 M HCl and estimate (find) the pH of the solution after the addition (mixing) of each 1 mL solution. In this way, add about 30 mL of 0.1 M HCl solution and then record the data in the table.

Observations

The volume of 0.1 M NaOH (basic) solution taken = 20.0 mL.

Result

• The pH of the solution decreases with the addition of 0.1M HCl.

• The decrease in pH is slow in the beginning.

• The point where there is a sharp fall in pH corresponds to the equivalence point.

Precautions

1. Handle the solutions carefully.

2. Burettes and beakers must be clean.

3. Solutions should be freshly prepared.

4. Universal indicators must be added carefully as per the requirement.

Lab Manual Questions

1. Name some universal indicators.

Ans: Some universal indicators are Methyl red, Phenolphthalein, Thymol blue etc.

2. Give the equation when the NaOH solution was added to the HCl solution.

Ans: The overall equation for this reaction is

NaOH+HCl→H2O+NaCl

3. Describe the titration of HCl with NaOH.

Ans: During the titration, the titrant (NaOH) is added slowly to the unknown solution. As it is added, the HCl slowly reacts away. The point at which exactly enough titrant (NaOH) has been added to react with all the analytes (HCl) is called the equivalence point.

4. What amount of HCl has been used in this experiment?

Ans: 0.1 M amount of HCl has been used in this experiment.

Viva Questions

1. Define pH.

Ans: It is defined as the negative logarithm of hydronium ion concentration in moles per litre. pH = – log [H3O+].

2. What do you mean by pOH?

Ans: It is a negative logarithm of OH- ion concentration. pOH = – log [OH-] = 14—pH.

3. What happens to the pH of the solution if a little acid is added to water?

Ans: When a little acid has been added, the concentration of H3O+ ions in the solution increases. Thus, the pH of the solution decreases.

4. Out of lemon juice and apple juice, which one would have lower pH?

Ans: Lemon juice would have lower pH as it is more acidic.

5. What is the effect of dilution on the pH of (i) an acidic solution and (ii) a basic solution?

Ans: The effect of dilution on the pH: 

• The pH of an acidic solution increases on dilution. 

• The pH of a basic solution decreases on dilution.

6. What do you mean by universal indicator?

Ans: It is a mixture of several indicators having different pH ranges. It shows many colour changes over a wide range of pH. Each colour corresponds to a certain approximate pH.

7. Which of the following solutions has lower pH: 0.1 M HCl or 0.1 M CH3COOH?

Ans: 0.1 M HCl would have lower pH because HCl being a strong acid produces a higher concentration of hydronium ions.

8. The pH of the sodium carbonate solution would be less than 7 or more than 7.

Ans: More than 7 because sodium carbonate is a salt of a strong base and weak acid, giving an alkaline solution due to hydrolysis.

9. What is the relationship between the pH and pOH of an aqueous solution?

Ans: The relationship is pH + pOH =pkw= 14 (at 298 K).

Practical Questions

1. Which of the following titrations will have the equivalence point at a pH of more than 8?

1. NH3 and HCl

2. HCl and NaOH

3. CH3COOH and NaOH

4. CH3COOH and NH3

Ans: CH3COOH and NaOH have an equivalence point at a pH of more than 8.

2. Which of the following is used as an indicator in the titration of a strong acid and a weak base?

1. Fluorescein

2. Thymol blue

3. Methyl orange

4. Phenolphthalein

Ans: Methyl orange is used as an indicator in the titration of a strong acid and a weak base.

3. What is the molarity of the solution of barium hydroxide, if 35 ml of 0.1 M HCl is used in the titration of 25 ml of the barium hydroxide solution?

1. 0.07

2. 0.28

3. 0.35

4. 0.14

Ans: 0.07

4. Which of the following is used as an indicator in the titration of a weak acid and a strong base?

1. Methyl red (5 to 6)

2. Methyl orange (3 to 4)

3. Phenolphthalein (8 to 9.6)

4. Bromothymol blue (6 to 7.5)

Ans: Phenolphthalein (8 to 9.6) is used as an indicator in the titration of a weak acid and a strong base.

5. A difference between strong and weak acids is:

1. Negative and positive pH

2. Complete and partial ionisation

3. Proton donation and electron acceptance

4. none

Ans: A difference between strong and weak acids is complete and partial ionisation.

6. Which of the following is a buffer solution?

1. NaCl + NaOH

2. CH3COONa + CH3COOH 

3. H2SO4+CuSO4

4. None

Ans: CH3COONa + CH3COOH is a buffer solution.

7. The ideal indicator for the titration of strong acid and weak base should have a pH range between:

1. 7-8

2. 8-10

3. 5-8

4. 4-6

Ans: 4-6 is the required pH range.

8. Find the concentration of HCl, if 10 ml of 0.5 M Ca(OH)2 is required to titrate 50 ml of HCl.

1. $\left [ \dfrac{1}{{10}} \right ]$ M

2. 10 M

3. $\left [\dfrac{1}{5}\right ]$ M

4. 5 M

Ans: $\left [\dfrac{1}{5}\right ]$ M is the required concentration.

Conclusion

From the above experiment, we can conclude that the pH of the solution decreases with the addition of 0.1 M HCl. The decrease in pH is slow in the beginning. After the addition of about 19.0 mL solution, the further addition shows a sharp fall in pH. After the sharp fall, the decrease in pH again becomes slow. The point where there is a sharp fall in pH (from about 10 to 3) corresponds to the equivalence point.