The provided solution contains 6 grams of Na2CO3 dissolved per cubic decimeter (dm3). Using the volumetric method, determine the percentage purity of the sample solution.

Theory:

The percentage purity of a sample solution can be determined by using a volumetric method. In this method, a known volume of a standard solution is added to the sample solution until the reaction is complete. The amount of the standard solution required to complete the reaction is used to calculate the concentration of the sample solution. The percentage purity of the sample solution can then be calculated using the formula:

Percentage purity = (concentration of sample solution / theoretical concentration of pure substance) x 100

Equipment:

• Analytical balance
• Volumetric flask (1 dm³)
• Pipette (10 ml)
• Burette (50 ml)
• Conical flask (250 ml)
• Stirrer
• Beaker (100 ml)
• Funnel
• Filter paper
• Safety goggles and gloves

Chemicals:

• Na₂CO₃ (6 g)
• Distilled water
• HCl (0.1 M)

Indicator:

• Methyl orange

End Point:

• Yellow to red

Equations:

Na₂CO₃ + 2HCl → 2NaCl + H₂O + CO₂

Procedure:

• Weigh 6 g of Na₂CO₃ using an analytical balance and transfer it to a 1 dm³ volumetric flask.
• Add distilled water to the flask until the volume reaches the mark.
• Stopper the flask and shake it to dissolve the Na₂CO₃.
• Take a 10 ml aliquot of the Na₂CO₃ solution using a pipette and transfer it to a 250 ml conical flask.
• Add 2-3 drops of phenolphthalein indicator to the conical flask.
• Titrate the Na₂CO₃ solution with 0.1 M HCl solution from the burette until the color changes from yellow to red.
• Note down the volume of HCl used.
• Repeat the titration two more times to get accurate results.
• Using the following table record the average volume of HCl used in the titration.

No. of obs.	Initial reading (mL)	Final reading (mL)	Volume used (mL)
1	0.0	9.8	9.8
2	9.8	19.5	9.7
3	19.5	29.1	9.6

Supposed Calculations:

Let’s say the average volume of HCl used in the titration is 9.7 ml.

From the balanced equation, we know that:
No of moles of Na₂CO₃ reacting \(= n_{1} = 1 mole\)

No of moles of HCl Reacting = \(= n_{2} = 2 mole\)

The volume of impure Na₂CO₃  solution taken \(= V_{1} = 10.0 ml\)

Concentration of Na₂CO₃ solution \(= M_{1} = ? M\)

Concentration of HCl solution \(= M_{2} = 0.1 M\)

Calculate the concentration of the Na₂CO₃ solution using the equation:

$$Base : Acid$$

$$\frac{M_{1} V_{1}}{n_{1}}=\frac{M_{2} V_{2}}{n_{2}}$$

$$M_{1}=\frac{M_{2} V_{2}n_{1}}{n_{2}V_{1}}$$

Therefore,

$$M_{1}=\frac{0.1M \times 9.7ml \times 1 }{2 \times 10ml}$$

 

• Concentration of Na₂CO₃ solution (M₁) = (volume of HCl used (V₂)x molarity of HCl (M₂) ) / (2 x volume of Na₂CO₃ solution V₁)
• (0.1 x 9.7) /10 x 2 =
• Calculate the percentage purity of the Na₂CO₃ solution using the equation:
• Percentage purity = (concentration of Na₂CO₃ solution / 0.057 M) x 100

Observations:

Color of Na₂CO₃ solution: white
Color of methyl orange indicator: pink
Color of the solution at the end point of titration: yellow to red

Supposed Calculations:

Let’s say the average volume of HCl used in the titration is 8.5 ml.

Concentration of Na₂CO₃ solution = (6.5 ml x 0.1 M x 2) / 10 ml = 0.13 M

Percentage purity = (0.17 M / 0.106 M) x 100 = 169.8%

Precautions:

Wear safety goggles and gloves while handling chemicals.
Use clean and dry glassware to avoid contamination.
Rinse the burette and pipette with distilled water before use.
Keep the burette tap closed while filling it with the standard solution.
Record the volume of the standard solution at eye level.
Shake the conical flask gently while titrating to avoid splashing.
Repeat the titration to obtain accurate results.