If you are titrating 25.00 mL of a 1.00mg/ml CaCO3 solution, how many mL of 0.010M EDTA solution will be required to reach the equivalence point
To find out how many milliliters of 0.010M EDTA solution will be required to reach the equivalence point, we need to use the concept of molarity and stoichiometry.
The balanced chemical equation for the reaction between CaCO3 and EDTA is:
CaCO3 + EDTA → Ca(EDTA) + CO2
From the balanced equation, we can see that the stoichiometric ratio between CaCO3 and EDTA is 1:1. This means that 1 mole of CaCO3 reacts with 1 mole of EDTA.
Given that the molarity of the CaCO3 solution is 1.00 mg/mL, we can calculate the number of moles of CaCO3 in the 25.00 mL of the solution as follows:
Mass of CaCO3 = Volume × Concentration
= 25.00 mL × 1.00 mg/mL
= 25.00 mg
Moles of CaCO3 = Mass / Molar mass of CaCO3
= 25.00 mg / (40.08 g/mol + 12.01 g/mol + 3 * 16.00 g/mol)
= 25.00 mg / 100.09 g/mol
= 0.2496 mmol
Since the stoichiometric ratio between CaCO3 and EDTA is 1:1, the number of moles of EDTA required to reach the equivalence point is also 0.2496 mmol.
Next, we need to calculate the volume (in liters) of 0.010M EDTA solution containing 0.2496 mmol of EDTA:
Volume (in L) = Moles / Molarity
= 0.2496 mmol / (0.010 mol/L)
= 0.02496 L
= 24.96 mL
Therefore, approximately 24.96 mL of 0.010M EDTA solution will be required to reach the equivalence point in the titration.