Back Titration Lab Report. In my experiment, I hoped to find the amount of calcium carbonate in some mineral limestone using the back titration method

I than mounted the burette on the retort stand, and filled it with the 0.10 mol dm-3 Sodium Hydroxide solution up to the 50 cm3 mark.

After doing all this, I titrated the solution of the limestone against the base from the burette, until the colour of the solution in the conical flask just turned pink. I repeated the procedure to obtain 3 sets of readings and will take the average value as my result.

Data processing and presentation:

With all the above processes being done, I wrote down the values which I obtained in the raw data table below.

Raw data

I then processed my data to find the average volume of acid used.

Vavg = V1+V2+V3

                 3

Vavg = 41.6 + 41.4 +

                        3

Vavg = 41.5 (± 0.05 cm3)  

Now I used the various formulas related to the mole concept to find the amount of calcium carbonate in the sample of limestone.

The equations of the reactions are as follows:

1) 2HCl + CaCO3 → CaCl2 + CO2 + H2O

                               

                                   2) HCl + NaOH → NaCl + H2O

Moles of acid in first reaction = Molarity x Average volume

                                                                   1000

                                                = 1 x (50 ± 0.02 cm3)

                                                           1000

                                                = 0.05 moles

                                             

Moles of base used = Molarity x Average volume

                                                  1000

                               = 0.1 x (41.5 ± 0.05 cm3)  

                                              1000

                               =4.15 x 10-3 moles of Sodium Hydroxide.

Moles of acid = 4.15 x 10-3 moles

Molarity of HCl = Moles x 1000

                                   Volume

                           = 4.15 x 10-3 x 1000

                                      25 ± 0.1

                           = 0.166 mol dm-3

Moles of HCl in 250 cm3 solution = Morality x Average volume

                                                                         1000

                                                             = 0.166 x (250 x ± 0.2 cm3)

                                                                               1000

                                                             = 0.0415 moles of HCl in 250 cm3 solution

Moles of acid used in reaction = 0.05 – 0.0415

                                                    = 8.15 x 10-3 moles

Ratio of moles of CaCO3 : HCl

                                   1      :   2

                                    x     : 8.15 x 10-3

                                    x    = 8.15 x 10-3 x 1  

                                                        2

                                          = 4.075 x 10-3 moles

Mass of CaCO3 in 1.5g of limestone  = Moles x RMM         (RMM of CaCO3 = 100)

                                                            = 4.075 x 10-3 x 100

                                                            = 0.41g

Percentage purity of calcium carbonate = 0.41 x 100

                                                                   1.5

                                                                = 27.3 ± 0.35 %

                                                                                                                                       

Conclusion:

A possible source of error in this experiment is the determination of the end-point, which is characterised by the solution just turning orange. This is because a slightly greater volume of acid may have been used than required to produce the pink colour. To try to reduce the effects of this error I would like to carry out a large number of titrations and their average used in the calculation.

Another possible way to reduce error is by using more accurate measuring instruments like a more precise burette so as to reduce the uncertainty of the measurements.