Essential Engineering Chemistry Laboratory Experiments

Experiment 1: Strength of Hydrochloric Acid

Aim

To determine the strength of the given dilute hydrochloric acid (HCl) solution by titrating it against a standard sodium carbonate (Na₂CO₃) solution.

Chemicals Required

• Standard sodium carbonate (Na₂CO₃) solution
• Dilute hydrochloric acid (HCl) solution
• Methyl orange indicator

Apparatus Required

• Burette
• Burette stand
• Pipette
• Conical flask
• Funnel

Chemical Reaction

Na₂CO₃ + 2HCl → 2NaCl + H₂O + CO₂

Principle

Normality is defined as the number of gram equivalents of solute present in one litre of solution. According to the law of equivalence, the relationship between the reacting solutions is given by:

N₁V₁ = N₂V₂

Where:
N₁ - Normality of HCl (unknown)
V₁ - Volume of HCl used
N₂ - Normality of Na₂CO₃ (known)
V₂ - Volume of Na₂CO₃ (10 mL)

Procedure

1. Clean the burette thoroughly with distilled water and rinse it with the given HCl solution.
2. Fill the burette with HCl solution and adjust the level to the zero mark.
3. Using a pipette, transfer 10 mL of standard sodium carbonate solution into a clean conical flask.
4. Add 2–3 drops of methyl orange indicator; the solution will appear yellow.
5. Titrate the solution with HCl until the colour changes from yellow to light pink.
6. Record the final burette reading.
7. Repeat the titration to obtain concordant readings.

Result

The strength of the given hydrochloric acid solution is 3.285 g/L.

Precautions

Do's:

• Wear a laboratory coat during the experiment.
• Read chemical labels carefully before use.
• Wear gloves and wash hands thoroughly after completing the experiment.

Don'ts:

• Do not touch the face or eyes without washing hands.
• Never work alone in the laboratory.
• Never taste any chemical.

Experiment 2: Water Hardness Determination

Aim

To determine the temporary and permanent hardness of a given water sample using EDTA solution.

Chemicals Used

• Standard calcium carbonate (CaCO₃) solution
• Approximately 0.02 N EDTA solution
• Eriochrome Black-T (EBT) indicator
• Ammonium hydroxide–ammonium chloride buffer solution (pH 9–10)
• Given water sample (unknown hardness)

Apparatus Required

Conical flask, Burette, Pipette, Measuring cylinder, Funnel, Filter paper, Chopper, Beaker, Watch glass.

Principle

Hardness of water is defined as its soap-consuming capacity. Temporary hardness is caused by bicarbonates of calcium and magnesium, whereas permanent hardness is due to chlorides, sulphates, and carbonates of calcium and magnesium. Temporary hardness can be removed by boiling, while permanent hardness requires chemical treatment.

Calcium (Ca²⁺) and magnesium (Mg²⁺) ions present in water form stable complexes with EDTA. These ions are titrated with a standard EDTA solution using Eriochrome Black-T (EBT) as the indicator. The indicator forms a wine-red complex with metal ions, which changes to blue at the end point when EDTA forms a more stable metal-EDTA complex. The pH is maintained between 9 and 10 using an NH₄OH–NH₄Cl buffer solution.

Procedure

1. Prepare standard hard water by dissolving 1 g of calcium carbonate in one litre of distilled water.
2. Pipette out 10 mL of the standard hard water into a conical flask.
3. Add 5 mL of buffer solution followed by 2–3 drops of EBT indicator.
4. Titrate the solution with EDTA until the colour changes from wine red to blue, indicating the end point.
5. Repeat the titration to obtain concordant readings.
6. Take the given hard water sample and boiled water sample and repeat the above procedure for each.

Reactions

M²⁺ (Ca²⁺, Mg²⁺) + EBT → [M–EBT] (wine-red, less stable complex)
[M–EBT] + EDTA → [M–EDTA] + EBT (blue, more stable complex)

Result

• Total hardness of water = 0.2 g/L
• Permanent hardness = 0.12 g/L
• Temporary hardness = 0.2 g/L - 0.12 g/L = 0.08 g/L

Precautions

• Use distilled water for washing and rinsing glassware.
• Prepare the EDTA solution using double-distilled water.
• Add the same quantity of indicator in each titration.
• Maintain the pH between 9 and 10 throughout the titration by using buffer solution.

Experiment 3: Chloride Content in Water

Aim

To determine the chloride content of a given water sample by the Argentometric method (Mohr's Method).

Requirements

Apparatus

Conical flask, burette, beaker, measuring flask, pipette.

Reagents

Standard silver nitrate (AgNO₃) solution.

Indicator

Potassium chromate (K₂CrO₄) solution.

Theory

Chlorides are commonly present in water as sodium chloride, magnesium chloride, and calcium chloride. Concentrations above 250 ppm impart a salty taste, making water unsuitable for drinking.

In the Argentometric method, chloride ions are titrated with standard silver nitrate using potassium chromate as an indicator. The pH must be between 7 and 8. Silver nitrate reacts with chloride ions to form a white precipitate of silver chloride. Once all chloride ions are consumed, excess silver ions react with chromate ions to form a red precipitate of silver chromate, indicating the endpoint.

Reactions involved:

• Ag⁺ + Cl⁻ → AgCl (white precipitate)
• 2Ag⁺ + CrO₄²⁻ → Ag₂CrO₄ (red precipitate)

Procedure

(A) Titration with Blank Solution

1. Pipette 10 mL of distilled water into a conical flask.
2. Add 3–4 drops of potassium chromate indicator.
3. Titrate with standard silver nitrate solution with continuous shaking.
4. The appearance of a light yellow colour indicates the endpoint.
5. Repeat until a concordant reading V₁ is obtained for blank correction.

(B) Titration with Sample Water

1. Pipette 10 mL of the given water sample into a conical flask.
2. Add 3–4 drops of potassium chromate indicator.
3. Titrate with standard silver nitrate solution until the endpoint is reached.
4. Repeat to obtain a concordant reading V₂.

Result

The chloride content present in the given water sample is 569 ppm.

Precautions

• All glassware should be thoroughly cleaned.
• Freshly prepared standard silver nitrate solution and potassium chromate indicator should be used.
• The pH of the solution should be maintained between 7 and 8 during titration.

Experiment 4: Chlorine in Bleaching Powder

Aim

To determine the percentage of available chlorine present in a given sample of bleaching powder.

Requirements

Apparatus

Burette, pipette, conical flask, volumetric flask, beaker.

Chemicals

Bleaching powder, potassium iodide solution, dilute acetic acid, 0.1 N sodium thiosulphate solution, distilled water.

Indicator

Freshly prepared starch solution.

Principle

The quantity of chlorine liberated by the action of a dilute acid on bleaching powder is known as available chlorine. The active constituent is calcium hypochlorite, Ca(OCl)₂. When treated with dilute acetic acid, it liberates chlorine gas.

Reaction:
Ca(OCl)₂ + 2CH₃COOH → Ca(CH₃COO)₂ + H₂O + Cl₂

The available chlorine is estimated by an iodometric method. In an acidic medium, liberated chlorine reacts with potassium iodide to release iodine, which is then titrated against standard sodium thiosulphate.

Reactions involved:

• Cl₂ + 2KI → 2KCl + I₂
• I₂ + 2Na₂S₂O₃ → Na₂S₄O₆ + 2NaI
• Starch + I₂ → Blue coloured complex

Procedure

• Pipette out 20 mL of bleaching powder suspension into a 250 mL conical flask.
• Add 10 mL of 10% potassium iodide solution followed by 5 mL of dilute acetic acid.
• Cover the flask with a watch glass and allow the reaction to proceed.
• Titrate the liberated iodine with standard sodium thiosulphate until the dark brown colour changes to pale yellow.
• Add 1 mL of freshly prepared starch solution to produce a blue colour.
• Continue titration until the blue colour disappears.
• Repeat to obtain concordant readings.

Result

The amount of available chlorine present in the given sample is 0.1007 g.
Percentage of available chlorine = 6.035%

Precautions

• All apparatus should be thoroughly rinsed with distilled water before use.
• Measure all solutions accurately using a pipette and burette.

Experiment 5: Urea-Formaldehyde Resin

Objective

To prepare urea-formaldehyde resin.

Requirements

Apparatus

Beaker, conical flask, glass rod, measuring cylinder, fractional weight box, water bath.

Chemicals

Glacial acetic acid, 40% formaldehyde solution, urea, concentrated sulphuric acid, phenol, concentrated hydrochloric acid, distilled water.

Theory

Urea-formaldehyde resin is a condensation polymer formed by the reaction of urea with formaldehyde in the presence of an acid or alkaline catalyst.

Procedure

1. Take 5 mL of 40% formaldehyde solution in a clean beaker.
2. Add about 2.5 g of urea gradually with continuous stirring until a saturated solution is obtained.
3. Add a few drops of concentrated sulphuric acid slowly while stirring continuously.
4. A voluminous white solid mass forms suddenly, indicating resin formation.
5. Wash the solid residue thoroughly with water.
6. Dry the product and calculate the percentage yield.

Result

The percentage yield of the prepared urea-formaldehyde resin is 48%.

Precautions

• Handle formaldehyde and concentrated sulphuric acid carefully.
• Add sulphuric acid dropwise with constant stirring.
• Perform the experiment in a well-ventilated area.
• Wash the resin thoroughly and dry it completely before weighing.

Experiment 6: Iron Content in Iron Ore

Aim

To determine the iron content in the given iron ore sample by titrimetric analysis against potassium dichromate solution using potassium ferricyanide as an external indicator.

Requirements

Standard potassium dichromate solution (N/10), potassium ferricyanide indicator, dilute sulphuric acid, conical flask, burette, pipette, white glazed tile, iron ore solution (Mohr’s salt).

Principle

The iron ore sample contains ferrous ions (Fe²⁺), which are oxidized to ferric ions (Fe³⁺) by potassium dichromate in the presence of dilute sulphuric acid. Potassium dichromate acts as a strong oxidizing agent.

The end point is detected using potassium ferricyanide as an external indicator. Before the end point, ferrous ions react with the indicator to produce a greenish-blue (Turnbull’s blue) precipitate. After complete oxidation, no blue colour is formed.

Reactions

Oxidation reaction:
Cr₂O₇²⁻ + 6Fe²⁺ + 14H⁺ → 2Cr³⁺ + 6Fe³⁺ + 7H₂O

Indicator reaction (before end point):
3Fe²⁺ + 2[Fe(CN)₆]³⁻ → Fe₃[Fe(CN)₆]₂ (Turnbull’s blue precipitate)

Procedure

1. Pipette 10 mL of the sample solution into a clean conical flask and add 3 mL of distilled water.
2. Rinse and fill the burette with N/10 potassium dichromate solution.
3. Place a few drops of potassium ferricyanide indicator on a white glazed tile.
4. Titrate the sample solution with potassium dichromate.
5. After each addition, place one drop of the reaction mixture on an indicator drop on the tile.
6. Continue titration until no blue colour is produced with the indicator.
7. Repeat until three concordant readings are obtained.

Result

The ferrous iron content present in the supplied sample is 0.01428 g.

Precautions

• Use freshly prepared and standardized potassium dichromate solution.
• Add potassium dichromate slowly near the end point.
• Use potassium ferricyanide only as an external indicator.
• Avoid excess contact of the reaction mixture with air to prevent oxidation of Fe²⁺.

Experiment 7: Phenol-Formaldehyde Resin

Aim

To prepare phenol-formaldehyde resin (Bakelite).

Requirements

Apparatus

Beaker, conical flask, glass rod, measuring cylinder, water bath.

Chemicals

Phenol, 40% formaldehyde solution, glacial acetic acid, concentrated sulphuric acid, concentrated hydrochloric acid, distilled water.

Theory

Phenol-formaldehyde resins are condensation polymers. The preparation of Bakelite occurs in stages: formation of hydroxybenzyl alcohols, formation of Novolac resin, and finally conversion into Bakelite, a hard thermosetting plastic.

Procedure

1. Take 5 mL of glacial acetic acid and 2.5 mL of 40% formaldehyde solution in a 500 mL beaker.
2. Add 2 g of phenol followed by 1 mL of concentrated hydrochloric acid.
3. Heat the mixture gently on a water bath with continuous stirring for about 5 minutes.
4. A pink-coloured plastic mass is formed.
5. Wash the solid residue repeatedly with distilled water and dry it completely.
6. Weigh the dried resin and calculate the percentage yield.

Result

The yield of obtained Bakelite is 4.3 g.

Precautions

• Phenol is toxic and corrosive; handle it carefully using gloves.
• Add concentrated acids slowly and with caution.
• Heat the reaction mixture gently to avoid charring.
• Stir continuously to ensure uniform polymerization.

Experiment 8: Lead-Acid Storage Battery

Aim

To study the construction and working of a lead-acid storage battery and to demonstrate its charging and discharging characteristics.

Apparatus

Lead plates (Pb), lead dioxide plates (PbO₂), dilute sulphuric acid (electrolyte), battery container, porous separators, connecting wires, ammeter, voltmeter, DC power supply, hydrometer.

Theory

A lead-acid battery is a secondary electrochemical cell. It converts chemical energy into electrical energy during discharge and vice versa during charging.

Construction of Lead-Acid Battery

• Positive plate: Lead dioxide (PbO₂)
• Negative plate: Sponge lead (Pb)
• Electrolyte: Dilute sulphuric acid (H₂SO₄) with specific gravity 1.28–1.30
• Separator: Porous rubber or PVC sheets
• Cell Voltage: Each cell produces 2 V; a 12 V battery has six cells in series.

Working Principle

During Discharge

• Chemical energy is converted into electrical energy.
• Both electrodes convert into lead sulphate (PbSO₄).
• The concentration of sulphuric acid decreases.

During Charging

• Electrical energy is converted into chemical energy.
• Lead sulphate reconverts into Pb and PbO₂.
• The concentration of sulphuric acid increases.

Procedure

A. Demonstration Setup

1. Arrange a transparent container and place PbO₂ and Pb plates alternately with separators.
2. Fill the container with dilute sulphuric acid.
3. Connect the plates externally to form a 2 V cell.

B. Discharging Process

1. Connect a load (bulb) across the terminals.
2. Observe the decrease in voltage and electrolyte specific gravity.

C. Charging Process

1. Connect the battery to a DC power supply with correct polarity.
2. Observe the increase in voltage, gas bubble formation, and specific gravity.

Observations

Result

The construction, working, charging, and discharging of a lead-acid storage battery were successfully demonstrated as a reversible electrochemical system.

Precautions

• Handle sulphuric acid carefully to avoid burns.
• Do not short-circuit the battery terminals.
• Ensure proper ventilation during charging to manage gas emission.
• Wear safety goggles and rubber gloves.