CBSE Chemistry Experiment - 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
Burette
Pipette (20 ml)
Titration flask
Beakers
Funnel
Universal indicator solution
0.1 M HCl
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
Take a clean (well-washed) burette and rinse (fill) it with 0.1 M HCl solution and then fill (pour) it with this solution.
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.
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.
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.
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
Handle the solutions carefully.
Burettes and beakers must be clean.
Solutions should be freshly prepared.
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
Which of the following titrations will have the equivalence point at a pH of more than 8?
NH3 and HCl
HCl and NaOH
CH3COOH and NaOH
CH3COOH and NH3
Ans: CH3COOH and NaOH have an equivalence point at a pH of more than 8.
Which of the following is used as an indicator in the titration of a strong acid and a weak base?
Fluorescein
Thymol blue
Methyl orange
Phenolphthalein
Ans: Methyl orange is used as an indicator in the titration of a strong acid and a weak base.
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?
0.07
0.28
0.35
0.14
Ans: 0.07
Which of the following is used as an indicator in the titration of a weak acid and a strong base?
Methyl red (5 to 6)
Methyl orange (3 to 4)
Phenolphthalein (8 to 9.6)
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.
A difference between strong and weak acids is:
Negative and positive pH
Complete and partial ionisation
Proton donation and electron acceptance
none
Ans: A difference between strong and weak acids is complete and partial ionisation.
Which of the following is a buffer solution?
NaCl + NaOH
CH3COONa + CH3COOH
H2SO4+CuSO4
None
Ans: CH3COONa + CH3COOH is a buffer solution.
The ideal indicator for the titration of strong acid and weak base should have a pH range between:
7-8
8-10
5-8
4-6
Ans: 4-6 is the required pH range.
Find the concentration of HCl, if 10 ml of 0.5 M Ca(OH)2 is required to titrate 50 ml of HCl.
$\left [ \dfrac{1}{{10}} \right ]$ M
10 M
$\left [\dfrac{1}{5}\right ]$ M
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.
FAQs on Study of the pH Change in the Titration of [M/10] HCl with [M/10] NaOH using a Universal Indicator
1. What is the balanced chemical equation for the neutralisation reaction between hydrochloric acid (HCl) and sodium hydroxide (NaOH), and what is the expected pH at the equivalence point for the 2025-26 CBSE curriculum?
The balanced chemical equation for this neutralisation reaction is: HCl + NaOH → NaCl + H₂O. Since hydrochloric acid is a strong acid and sodium hydroxide is a strong base, they completely neutralise each other. The resulting salt, sodium chloride (NaCl), is neutral and does not hydrolyse water. Therefore, the pH at the equivalence point is exactly 7 (neutral) at 298 K.
2. What is the specific role of a universal indicator in the titration of HCl with NaOH, and how does its colour change signify the progress of the reaction?
A universal indicator is a pH indicator composed of a mixture of several compounds that exhibits a spectrum of smooth colour changes over a wide pH range. In this experiment, its role is to visually track the pH as the acid is added to the base:
- Initially: In the conical flask containing NaOH (a strong base), the indicator shows a blue or violet colour, indicating a high pH (approx. 13).
- During Titration: As HCl is added, the pH gradually decreases, and the indicator's colour shifts through green towards yellow.
- At Equivalence Point: A single drop of HCl causes a very large and sudden drop in pH. The indicator's colour changes rapidly from greenish-yellow to orange or red, signalling that the neutralisation point has been reached.
3. From an exam perspective, what are two important precautions a student must take when performing this titration to ensure an accurate result?
For the Class 11 practical exam, two crucial precautions are:
- Rinsing the Apparatus Correctly: The burette must be rinsed with the titrant (HCl solution), and the pipette must be rinsed with the analyte (NaOH solution). This prevents any residual water from diluting the standard solutions, which would cause errors in the final calculation.
- Ensuring No Air Bubbles: Before starting the titration, it's essential to remove any air bubbles from the jet or tip of the burette. An air bubble displaced during titration would be counted as part of the volume of the titrant, leading to an inaccurate and higher burette reading.
4. Why does the pH change very slowly at the beginning and end of the titration but drop sharply near the equivalence point?
This phenomenon is due to the logarithmic nature of the pH scale and the changing concentrations of H⁺ and OH⁻ ions.
- At the start: The solution is strongly basic with a high concentration of OH⁻ ions. The initial drops of HCl are neutralised, but the overall concentration of OH⁻ remains high, so the pH changes only slightly.
- Near the equivalence point: Almost all the OH⁻ ions have been neutralised. At this critical stage, even a tiny amount of added HCl causes a massive change in the relative concentration of H⁺ ions, resulting in a steep, sudden drop in the pH value.
- After the equivalence point: The solution is now acidic and contains excess H⁺ ions. Adding more acid at this point does not significantly alter the already high H⁺ concentration, so the pH curve flattens out again at a low pH value.
5. How would the titration curve and equivalence point for this experiment (strong acid-strong base) be different if a weak acid like acetic acid (CH₃COOH) were used instead of HCl?
There would be several key differences in the titration curve, which are often asked as higher-order thinking questions:
- Starting pH: The titration with HCl starts at a very high pH (~13). The titration with weak acid (CH₃COOH) would start at a lower pH because the initial base is the same, but the acid is weaker. Wait, the prompt titrates acid *into* the base. So the initial pH in the flask is high for NaOH. The final pH after excess acid addition would be lower for HCl than for CH₃COOH.
- Equivalence Point: For the HCl-NaOH titration, the equivalence point is at a neutral pH of 7. For a CH₃COOH-NaOH titration, the salt formed (CH₃COONa) undergoes hydrolysis to produce OH⁻ ions, making the solution basic. Therefore, the equivalence point would be above pH 7.
- Buffer Region: The titration curve for a weak acid features a distinct buffer region before the equivalence point where the pH changes very little upon acid addition. This region is absent in a strong acid-strong base titration.
6. Explain how understanding the pH curve from this experiment helps in selecting the correct indicator for other types of acid-base titrations.
Understanding the pH curve is fundamental to indicator selection. An effective indicator is one whose pH range of colour change falls entirely within the steep, vertical portion of the titration curve.
- This ensures that the colour change (the endpoint) occurs precisely when the reaction reaches its equivalence point.
- For a strong acid-strong base titration, the vertical drop is large (e.g., from pH 10 to 3), so indicators like phenolphthalein (pH 8.2–10.0) and methyl orange (pH 3.1–4.4) are both suitable.
- However, for a weak acid-strong base titration where the equivalence point is >7, only phenolphthalein would be appropriate, as methyl orange would change colour too early.
7. A student titrates 25 mL of 0.1 M NaOH solution with 0.1 M HCl. What is the expected volume of HCl required to reach the equivalence point? This is a frequently asked calculation in exams.
To solve this, we use the stoichiometry of the reaction. Since HCl and NaOH react in a 1:1 molar ratio, we can use the formula M₁V₁ = M₂V₂.
- Let M₁ and V₁ be the molarity and volume of HCl.
- Let M₂ and V₂ be the molarity and volume of NaOH.
- M₁ = 0.1 M, V₁ = ?
- M₂ = 0.1 M, V₂ = 25 mL
(0.1 M) × V₁ = (0.1 M) × (25 mL)
V₁ = 25 mL
Therefore, 25 mL of 0.1 M HCl is required to reach the equivalence point.
