Classic Titration Lab
AP Chemistry
Titration is a method used to measure the concentration of a particular solute in a solution. It is type of volumetric
quantitative analysis (the use of volume measurements to analyze an unknown). If we know the molarity and
volume of one of the solutions, we can determine the molarity of the other solution if we know its volume.
Titrations exploit the reaction between two solutions. For example, when an acid and a base are mixed together
they react to form water and a salt. The pH will change as the acid and base are mixed together. Oxidationreduction reactions and precipitation reactions can also be used.
A different chemical species which changes color with pH can be used to indicate when an acid-base neutralization
reaction is complete, i.e. when stoichiometrically equivalent amounts of the two solutions have been mixed.
The equivalence point of a titration is the point where the stoichiometrically correct number of moles of each
reactant is present. The end point is where the indicator changes. They are not the same. We choose an indicator
which changes color close to the equivalence point the better.
Purpose: to determine the concentration of an unknown H2SO4 solution
Materials:
Ring stand
50 mL buret
buret clamp
2 - 150 mL beakers
2 - 250 mL beakers
125 mL Erlenmeyer flask
funnel
25.0 mL graduated cylinder
2 Beral-type pipets
Distilled water wash-bottle
3” x 5” card with slit
~ 100 mL 0.125 M HCl (standard)
~ 100 mL NaOH of unknown concentration
~ 100 mL H2SO4 of unknown concentration
Phenolphthalein, 0.5%, ~ 1 mL
Safety: Use apron and goggles. NaOH is a strong base, HCl and H2SO4 are strong acids. All solutions are
harmful to skin, especially eyes.
Pre-lab questions
1. What volume of 0.250 M NaOH is needed to neutralize 25.0 mL of 0.125 M HCl?
2. What is the purpose of the phenolphthalein?
3. What happens to the number of moles of HCl in the beaker if you add distilled water to increase the
volume?
4. What volume of 0.250 M NaOH is needed to neutralize 25.0 mL of 0.100 M H2SO4?
5. One student accidentally “overshot” the endpoint and added more NaOH solution. Is the calculated
concentration of the NaOH likely to be too high or too low as a result of this error?
6. Create a suitable data table in your lab notebook to record your results.
1
Procedures
1. Rinse the buret with several mL of distilled water. Allow to drain upside down.
2. Rinse the buret twice with about 5 mL of the NaOH solution.
3. Clamp the buret to the ringstand using the buret clamp and place the beaker
under the buret (this is the waste beaker). See Figure 1 for final setup.
4. Close the stopcock and fill the buret with the NaOH solution above the zero mark,
using a funnel.
5. Open the stopcock and allow the NaOH solution to flow into the waste beaker
until there are no air bubbles.
6. Fill the buret to the zero line, using a funnel. Open the stopcock and slowly
run the NaOH solution into the waste beaker until the bottom of the meniscus
is exactly at “0”.
Figure 1
7. Use the 3” x 5” card in such a way that you can easily read the markings on the buret.
8. Collect about 100 mL of 0.125 M HCl in a 150 mL beaker from the prep table.
Collect about 200 mL of the NaOH solution in a 250 mL beaker from the prep table.
9. Carefully measure 25.0 mL of the HCl into a graduated cylinder. Use a Beral-type pipet to make sure the
bottom of the meniscus is exactly on 25.0 mL.
10. Pour the HCl into the 125 mL Erlenmeyer flask and add a stir bar.
11. Add approximately the same volume of distilled water (easy to estimate with 50 mL line).
12. Add a couple of drops of phenolphthalein.
13. Begin your titration adding NaOH in 1 mL increments. Turn on the magnetic stirrer to make sure the
NaOH and HCl mix thoroughly.
14. When you are within a few mL of the volume of NaOH you predicted in pre-lab question 1, add smaller
and smaller volumes of the titrant. When you are very close, the phenolphthalein will change to fuchsia,
then disappear. When you reach the endpoint, which is one drop beyond the equivalence point, the solution
will turn completely fuchsia in color.
15. Record the volume of NaOH at the endpoint in your data table. The volume should be to two decimal
places, as buret markings are to the first decimal place (sig figs!).
16. Rinse your Erlenmeyer flask twice with distilled water.
17. Repeat your titration twice, and be sure to record each volume in your data table. Rinse your Erlenmeyer
flask with distilled water before trials. You may top up your buret with NaOH each time, or continue on,
recording the beginning and ending volumes. There is no need to drain and refill your buret.
18. Repeat the titration three times with the unknown concentration of H2SO4. Record your results in your data
table.
19. All solutions can be poured down the sink. Rinse the beakers and flasks with distilled water and leave
them upside down on paper towels at your lab station.
2
Analysis and Conclusions
1. Calculate the molarity of the NaOH solution. Show your work.
2. Calculate the molarity of the H2SO4 solution. Show your work.
3. Consider the following potential sources of error in the titration. Fill in the blank:
“H” if the error would have caused the calculated molarity of NaOH to come out too high.
“L” if the error would have caused the calculated molarity of NaOH to come out too low.
“N” if the error would have had no effect on the calculated molarity of NaOH.
a.
There was a little distilled water in the Erlenmeyer flask before the titration began. _______
b. There was a little HCl in the Erlenmeyer flask before the titration began. _______
c. There was a little distilled water in the buret before you began and you forgot to rinse it out with
NaOH. _______
d. You added 3 drops of phenolphthalein instead of 2 drops. ________
e. Some NaOH solution dripped into the Erlenmeyer flask before the initial NaOH volume was measured.
______
f.
While you were titrating, some NaOH dripped out onto the table instead of into the Erlenmeyer flask.
_____
LH, Adapted from Flinn ChemTopic Labs, Volume 13, 2002
3