Titration Experiment

Titration Experiment



PURPOSE: The purpose of this laboratory experiment is to familiarize the student with the principles of the technique of titration.

BACKGROUND: The technique of titration involves measuring the exact volume of a solution of known concentration that is required to react completely with a measured volume of a solution of unknown concentration, or a weighed sample of unknown solid. The solution with the known concentration is called a standard solution. The point when just enough of the standard solution has been added to completely react with the solution of unknown concentration is called the equivalence point.

The goal in titration is to reach the equivalence point without adding any excess standard solution. To do this, a means is needed to identify when the equivalence point of the reaction is reached. A common method of identifying the equivalence point of a reaction is through the use of indicators. Typically, an indicator changes color at or near the equivalence point. The exact point where the indicator changes color is called the end point. Knowledge of the type of reaction being studied allows us to pick an appropriate indicator so that the end point it signals is very close to the actual equivalence point. For the reaction being studied in this experiment, you can assume that the end point occurs at the equivalence point. Thus the change in color signals that the reaction is complete and the titration is complete.

From the volume of standard solution used in the titration and its known concentration, one may calculate the number of moles of known solute. From the balanced equation for the reaction involved and the number of moles of known solute, one then may calculate the number of moles of unknown solute in the solution of unknown concentration. This then allows the determination of the concentration of the unknown solution using its number of moles and its original measured volume.


NaOH (aq) + HCl (aq) → H2O + NaCl(aq)

M = moles/Liter

At the equivalence point,

moles of HCl = moles of NaOH


M(NaOH) x V (Liters of NaOH) = M(HCl) x V(Liters of HCl)


M(NaOH) x V(mL of NaOH) = M(HCl) x V(mL of HCl)







25-mL pipet

100-mL beaker

250-mL beaker

250-mL Erlenmeyer flask

Phenolphthalein indicator

NaOH solution of known concentration (approximately 0.900 M)*

HC1 solution of known concentration (approximately 0.500 M)*

HC1 solution of known concentration (approximately 2.00 M)*

HC1 solution of unknown concentration

*The concentrations of the known solutions of HCl and NaOH are approximate. Record and use the exact concentrations on the bottles.

PROCEDURE: In this experiment, you will first titrate two solutions of HCl (with a known concentration) using a standard solution of NaOH. The purpose of these titrations is to develop your titration skills. You will then titrate the HCl solution (unknown concentration) using the standard NaOH solution and determine the concentration of the unknown HCl solution. The procedure for each titration is the same.

Obtain approximately 250 mL of the standard NaOH solution in a 250 mL beaker. This should be enough NaOH for all of your titrations.

Wash all glassware you will use. A final rinse of glassware with the solution you will use in it will minimize the requirement to dry the glassware.

  1. Titration of HC1 solution (approximately 0.500 M):
  1. Obtain approximately 60 mL of the HCl solution in a 100-mL beaker and record its exact concentration.
  1. Using a pipet, transfer 25.00 mL of the HCl solution to a 250-mL Erlenmeyer flask.
  1. Add approximately 50 mL of distilled water to the acid in the flask. This provides a larger amount of solution to see a color change. It has no effect on the amount of acid or base used in this experiment.
  1. Add three drops of Phenolphthalein indicator to the acid solution.
  1. Fill the buret with the standard NaOH solution. Be sure the tip of the buret is filled and record the level of the NaOH in the buret.

  1. Titrate the acid solution by adding small amounts of the NaOH solution from the buret while swirling the Erlenmeyer flask. As you near the end point, the solution will show traces of pink color. Just at the end point, the solution will change to a light pink color. Stop the titration and record the amount of NaOH remaining.
  1. Repeat this experiment a second time with a second sample of the approximately 0.500 M HC1 solution. The two experiments should use the same amount of NaOH to within 0.50 mL. If the volume of NaOH used in the two experiments is not within 0.50 mL of each other, do a third run. Use the average amount of NaOH used for all calculations.
  1. Complete the report form showing all calculations in your Lab Notebook.
  1. Titration of HCl solution (approximately 2.00 M): Use same procedures as above substituting 2.00 M HCl for the 0.500 M HCl.
  1. Titration of HCl solution (unknown concentration): Use same procedures as above substituting HCl solution of unknown concentration for the 0.500 M HCl. HINT: The concentration of the unknown HCl solution is between the two known concentrations you used earlier.


  1. Show all calculations in your Lab Notebook.
  1. Complete the Lab Report Form for this experiment and turn it in with the tissue pages from your Laboratory Notebook.
  1. Complete an error analysis in your Lab Notebook.