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Lab 11A:  Titration of Hydrochloric Acid                                                                   SCH 3U










*Taken from unit11 guide



Both the hydrochloric acid and sodium hydroxide used in this experiment are corrosive substances. They could cause damage to skin and clothing. Wash away any spills with plenty of water.


Goggles must be worn during this lab.

Rubber gloves are to be worn when handling the solutions.

Waste solutions are to be flushed down the sink with copious amounts of water.




It is important that you review the lab outline thoroughly and fully understand both the purpose of the lab, and all of the procedures, before you attempt to complete the lab. Doing so will help to avoid time loss due to unfamiliarity with the lab procedures, and will ensure that your lab period is productive!




In this experiment, you will titrate a measured volume of hydrochloric acid with a solution of sodium hydroxide of known concentration. The acid and the base react with one another according to the equation:


HCl(aq)  +  NaOH(aq)   =>    NaCl(aq)  +  H2O(l)


The hydrochloric acid is placed in an Erlenmeyer flask, and phenolphthalein indicator is added. The sodium hydroxide solution is added from a buret into the flask containing the acid. During the first stages of the titration, the sodium hydroxide will be completely neutralized, and an excess of acid will remain. However, eventually there will be a point, the theoretical endpoint, at which the acid and the base have neutralized one another exactly, and no more base should be added to the flask.


The phenolphthalein indicator is used to determine experimentally the point, called the experimental endpoint, at which the base has neutralized the acid. Phenolphthalein is colourless in acid solution. It turns pink when the acid is

completely neutralized and a slight excess of base is present. In this titration, a successful endpoint is achieved if one drop of base turns the solution in the flask from colourless to a permanent light pink color.


In this experiment, you can use separate burets for the acid and the base. It is advantageous to use two burets. If you should overshoot the endpoint by adding too much base, you will be able to add an additional measured volume of acid from the acid buret. The additional acid will neutralize the excess base, and you can then add more base to reach a new endpoint.


Since you know both the concentration in moles per litre and the volume in millilitres (which you can convert to litres) of the sodium hydroxide, you can calculate the number of moles of hydrochloric acid used. At the endpoint, the number of moles of hydrochloric acid equals the number of moles of sodium hydroxide used. Thus, you know the number of moles of hydrochloric acid in a measured volume of acid and you can calculate the concentration of the hydrochloric acid in moles per litre.



For example, if 25.0 mL of 0.200 mol/L NaOH is able to neutralize 30.0 mL of HCl, we can calculate the concentration of the acid solution. First we determine the number of moles of NaOH:


25.0 ml NaOH   x   0.200 mol NaOH/1000 mL NaOH

 =  0.005 mol NaOH 


n = 0.005 mol NaOH


From the balanced equation the moles of HCl equals moles of base at the endpoint. Thus, 30.0 mL of HCl solution must contain 0.005 mol HCl, and we can calculate the concentration of the HCl solution:


25.0 ml NaOH


0.200 mol Na OH

1000 ml NaOH


1 mol HCl

1 mol NaOH



0.0300 L HCl

= 0.167 mol/L HCl


Remember! Always approach a chemistry problem by starting with a balanced chemical equation!


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