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Lab 11A:
Titration of Hydrochloric Acid
SCH 3U |
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Introduction
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*Taken from unit11 guide Safety: 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. Note: 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! Introduction: 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:
=
0.167 mol/L HCl Remember! Always approach a chemistry problem by starting with a
balanced chemical equation! |
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