Acids and Bases Titration Example Problem

Titration is an analytical chemistry technique used to find an unknown concentration of an analyte (the titrand) by reacting it with a known volume and concentration of a standard solution (called  the titrant). Titrations are typically used for acid-base reactions and redox reactions. Heres an example problem determining the concentration of an analyte in an acid-base reaction: Titration Problem A 25 ml solution of 0.5 M NaOH is titrated until neutralized into a 50 ml sample of HCl. What was the concentration of the HCl? Step-by-Step Solution Step 1 - Determine [OH-] Every mole of NaOH will have one mole of OH-. Therefore [OH-] 0.5 M. Step 2 - Determine the number of moles of OH- Molarity # of moles/volume # of moles Molarity x Volume # of moles OH- (0.5 M)(.025 L)# of moles OH- 0.0125 mol Step 3 - Determine the number of moles of H When the base neutralizes the acid, the number of moles of H the number of moles of OH-. Therefore the number of moles of H 0.0125 moles. Step 4 - Determine the concentration of HCl Every mole of HCl will produce one mole of H, therefore the number of moles of HCl number of moles of H. Molarity # of moles/volume Molarity of HCl (0.0125 mol)/(0.050 L)Molarity of HCl 0.25 M Answer The concentration of the HCl is 0.25 M. Another Solution Method The above steps can be reduced to one equation MacidVacid MbaseVbase where Macid concentration of the acidVacid volume of the acidMbase concentration of the baseVbase volume of the base This equation works for acid/base reactions where the mole ratio between acid and base is 1:1. If the ratio were different as in Ca(OH)2 and HCl, the ratio would be 1 mole acid to 2 moles base. The equation would now be MacidVacid 2MbaseVbase For the example problem, the ratio is 1:1 MacidVacid MbaseVbase Macid(50 ml) (0.5 M)(25 ml)Macid 12.5 MmL/50 mlMacid 0.25 M Error in Titration Calculations There are different methods used to determine the equivalence point of a titration. No matter which method is used, some error is introduced, so the concentration value is close to the true value, but not exact. For example, if a colored pH indicator is used, it may be difficult to detect the color change. Usually, the error here is to go past the equivalence point, giving a concentration value that is too high. Another potential source of error when an acid-base indicator is used is if water used to prepare the solutions contains ions that would change the pH of the solution. For example, if hard tap water is used, the starting solution would be more alkaline than if distilled deionized water had been the solvent. If a graph or titration curve is used to find the endpoint, the equivalence point is a curve rather than a sharp point. The endpoint is a sort of best guess based on the experimental data. The error can be minimized by using a calibrated pH meter to find the endpoint of an acid-base titration rather than a color change or extrapolation from a graph.