To determine the percentage purity of a sample of sodium carbonate.
Na2CO3 + 2HCl –> 2NaCl + H20 + CO2
Sodium Hydrochloric Sodium Water Carbon
Carbonate Acid Chloride Dioxide
* 4.00 g dm-3 of impure Sodium Carbonate in solution
* 1.00 mol dm-3 Hydrochloric acid
* 50cm3 burette burette clamp and stand
* 25cm3 bulb pipette pipette pump
* Teat pipette
* Conical flask white tile
* 250cm3 graduated flask
* 1 beaker for waste funnel
* Methyl Orange indicator
* Distilled water
To determine the percentage purity of a sample of impure sodium carbonate we must first find out how much, in volume, hydrochloric acid it takes to neutralise the solution (this is an acid and base titration and will need an indicator). To do this I will need to adhere to the following method;
1. Clamp the burette carefully. Fill the burette with the hydrochloric acid using the funnel. Remove the air space below the tap and use the wastage beaker to catch any of the acid. Remove the funnel
2. Use the pipette to take 25cm3 of the sodium carbonate solution and place it in the conical flask. Add 1-2 drops of methyl orange indicator. Place the conical flask directly under the burette.
3. Take note of the initial reading of the acid in the burette and titrate until the indicator changes from a yellow colour to a pink colour. Take note of the final reading of the acid in the burette.
4. Repeat a minimum of three times, and for each, calculate the volume of acid used by subtracting the initial reading from the final reading.
Before I start any titration, I must first dilute the hydrochloric acid to a suitable level
To calculate the volume of hydrochloric acid I will add to the 250cm3 graduated flask, I must first calculate the molarity of the sodium carbonate.
For sodium carbonate (Na2CO3):
Moles = Mass
= 0.0377 mol dm-3 (3s.f.)
= 0.4 (1s.f.)
From the formula equation I can see that 1 mole of sodium carbonate reacts with 2 moles of hydrochloric acid, therefore the molarity of the acid will be double that of the sodium carbonate.
Concentration of hydrochloric acid = 0.04 x 2 = 0.08 mol dm-3
To make up a solution of 250cm3 of 0.08 mol dm-3 hydrochloric acid using 1.00 mol dm-3 hydrochloric acid I must find out how much of the acid, and therefore how much distilled water, I should add to the graduated flask for a 250cm3 solution.
Volume needed using = volume x concentration = V x C
1.00 mol dm-3 moles 1000 1000
= 250 x 0.08 1000
According to the above equation, I will need to add 20cm3 of hydrochloric acid to a graduated flask that should also contain 230cm3 of distilled water to make up the rest of the 250cm3 solution of the acid. I should use a burette to add an accurate volume of hydrochloric acid to the graduated flask, and when filling up the flask a teat pipette should be used when making up the final few cm3 of the solution.
There is little risk involved in this experiment as all the solutions used are very dilute and are unlikely to harm anyone if spilled. However, for safety measures, a lab-coat and goggles should be worn and extra care should be taken when using the glassware.
The following table shows the error that can occur when reading off the measuring device. I must take these into consideration for my experiment as it could affect my results.
Volume / cm3
Error / cm3
Error per cent / %
The error taken for the burette is as such because I am taking a reading from it twice (initial and final reading), but I am reading to the nearest 0.05cm3 to give a greater degree of accuracy. I will also need to take into account the degree of accuracy the titration was, but first I will have to see the volume of hydrochloric acid that will be used in the titration.
To ensure accuracy in this experiment I must:
* Wash all equipment with distilled water and then whatever may be put in it. i.e. wash the burette out with a bit of acid.
* Make sure there are no air bubbles in jet in the burette
* Make sure the jet is below the conical flask
* Make sure that there is no funnel in the burette when titrating
* Read from the bottom of the meniscus
* Record all readings to the nearest 0.05cm3
* Make sure there are no air bubbles in the pipette
* Release solution and touch pipette onto surface
* Swirl the conical flask constantly
* Use single drops toward the end of the titration
* Record at least two concordant results after one rough titration
In the table below, I have tabulated my results for the titrations that I completed. The “Initial reading” and “Final reading” are indicating the level of hydrochloric acid used in the burette.
Initial reading / cm3
Final reading / cm3
Hydrochloric Acid used / cm3
The average of this titration shall be taken from the most similar results. In this case they are: 16.70, 16.65 and 16.60.
Average titration = 16.70 + 16.65 + 16.60
From the average titration I can work out the error percentage of the titration. This works out at 0.60%.
From the average titration volume I will be able to calculate the percentage purity of the sodium carbonate solution.
Moles HCl = Average HCl used x concentration HCl
= 16.65 x 0.08
Moles Na2CO3 = moles HCl
Moles Na2CO3 = Volume x Concentration
Concentration Na2CO3= Moles Na2CO3 x1000
= 0.00066 x 1000
Percentage purity of = Concentration Na2CO3 x RFM
= 0.0264x 106
Moles HCl = 16.65x 0.08
Moles Na2CO3 = 16.65 x 0.08
in 25cm3 1000 x 2
Moles in 250 cm3 = 16.65 x 0.08 x 10
1000 x 2
Mass Na2CO3 = 16.65 x 0.08 x 10 x 106
in 250 cm3 1000 x 2
Percentage purity = 70.596%
I now need to take into account the errors that could have been made.
Overall percentage = 0.50 + 0.24 +0.08 + 0.16 + 0.60 = 1.58%
So there is a possible error in my percentage purity of 1.58%,
Possible percentage = 70.596 ? 1.58 x 100
purity of Na2CO3 70.596
= 73.2% to 68.4%
Therefore the percentage purity of the sodium carbonate solution could range from the upper bound of 73.2%, to the lower bound of 68.4%.