Experiment 1
CHEM 131
Experiment 1B – Analysis of a H2SO4 Solution
In Experiment 1A, you used double displacement reactions in the qualitative analysis of an
unknown solution. The same reaction types can be used in quantitative analysis. In this
experiment, two quantitative methods will be used to determine the molarity of a sulfuric acid
solution.
Titration
Acid-base reactions, including those with gas formation, can be used in titration
experiments. In a titration, a solution with known concentration (the titrant) is slowly added
from a buret to a measured sample of the solution with unknown concentration. An indicator
changes color at the endpoint of the titration, when a stoichiometric amount of the titrant has
been added. The moles of solute in the unknown solution can be calculated from the volume of
titrant added. The concentration can be determined by dividing by the volume of the unknown
sample.
For example, the concentration of a phosphoric acid solution could be determined by
titrating it with a standardized NaOH solution. The titration reaction would be:
H3PO4 (aq) + 3 NaOH (aq)  3 H2O (l) + Na3PO4 (aq)
If a 25.00-mL sample of the phosphoric acid solution were used, and 28.14 mL of 0.2055 M
NaOH were required to reach the endpoint, the molarity of the phosphoric acid solution would be
calculated as follows.
Volume of titrant
0.02814 L 

Moles of titrant
0.2055 mol NaOH
1L


1 mol H 3PO 4
3 mol NaOH
Moles of unknown
= 1.9276  10–3 mol H3PO4
Moles of unknown
 Molarity of unknown
Liters of unknown
1.9276  103 mol H3 PO4
0.02500 L
CHEM 131
=
0.07710 M H3PO4
Experiment 1B
1
If you are planning to analyze a solution by acid-base titration, there are several points of
experimental design to consider:

Of course, the solute must undergo an acid-base reaction with some readily available reagent,
the titrant.

You will need an indicator to show the endpoint of the titration. If you are adding a strong
base to an acidic solution, phenolphthalein is a good indicator; it changes from colorless to
pink at the endpoint. If you are adding a strong acid to a basic solution, methyl orange is a
good indicator; it changes from yellow to pink at the endpoint.

The precision of your result depends on three factors:
1) The titrant must be standardized; that is, you need to know its precise concentration.
2) You need to measure the volume of the unknown sample precisely, with a pipet.
3) You need to know the precise volume of titrant needed to reach the endpoint. Since
this is measured with a buret, and burets can be read to the nearest 0.01 mL, 4
significant figures can be obtained if more than 10 mL of titrant is used.
The least precise of these three quantities determines the number of significant figures you
may report in your result.
In this experiment, you will use a standard NaOH solution to titrate the sulfuric acid
solution, with phenolphthalein as the indicator. The titration reaction will be:
H2SO4 (aq) + 2 NaOH (aq)  2 H2O (l) + Na2SO4 (aq)
Gravimetric Analysis
Precipitation reactions can be used in gravimetric analysis experiments. When
gravimetric analysis is used to determine an unknown concentration, an excess of precipitating
reagent is added to a measured sample of the unknown solution. The precipitate is filtered out of
the mixture, dried, and weighed. The moles of solute in the unknown solution can be calculated
from the mass of the precipitate. The concentration can be determined by dividing by the
volume of the unknown sample.
For example, the concentration of a magnesium chloride solution could be determined by
precipitating its chloride ions as silver chloride. If silver nitrate were used as the source of silver
ions, the precipitation reaction would be:
MgCl2 (aq) + 2 AgNO3 (aq)  2 AgCl (s) + Mg(NO3)2 (aq)
If a 25.00-mL sample of the magnesium chloride solution were used, and 1.228 g of AgCl were
obtained in the precipitation, the molarity of the magnesium chloride solution would be
calculated as follows.
CHEM 131
Experiment 1B
2
Mass of precipitate
1.228 g AgCl


Moles of precipitate 
1 mol AgCl
143.32 g AgCl

1 mol MgCl2
2 mol AgCl
Moles of unknown
= 4.2841  10–3 mol MgCl2
Moles of unknown
 Molarity of unknown
Liters of unknown
4.2841  10 –3 mol MgCl 2
0.02500 L
=
0.1714 M MgCl2
If you are planning to do a gravimetric analysis of a solution (e.g., for Experiment 1C), there are
several points of experimental design to consider:

The solute must undergo a precipitation reaction with some readily available reagent.

The precision of your results depends on how precisely you measure the unknown
sample, and on how precisely you are able to weigh the precipitate. Since our most
precise balances have a measurement uncertainty of ±0.0002 g, 4 significant figures can
be obtained if at least 0.1 g of precipitate forms.

Thus, it is desirable for the precipitate to have a high molar mass. The higher the molar
mass of the precipitate, the more grams of it will form.
There are some other design considerations that are specific to the particular precipitate formed:

In this experiment, you will add excess BaCl2 solution to a sample of the sulfuric acid
solution. The precipitation will be:
H2SO4 (aq) + BaCl2 (aq)  2 HCl (aq) + BaSO4 (s)
BaSO4 is a very fine precipitate, which sometimes passes through filter paper. Two steps
in the procedure were added in order to prevent this: adding HCl to the mixture, and
heating it. Both steps help to form large particles of BaSO4 that will not pass through
filter paper.

In the handout for Experiment 1C, special considerations are provided for a variety of
different precipitates that might result from gravimetric analysis of your unknown.
CHEM 131
Experiment 1B
3
Experiment 1B – Analysis of a H2SO4 Solution
Procedure
Collect about 50 mL of H2SO4 solution in dry graduated cylinder. Store it in a dry Florence
flask, with a stopper to prevent contamination. Record the concentration of the H2SO4 solution
so that you may check your work.
A – Acid-Base Titration
Collect about 80 mL of standardized NaOH solution. Record the concentration of this
solution. Rinse a 50-mL buret with distilled water, and then with three small portions of the
standardized NaOH solution. Be sure to rinse the tip of the buret. Fill the buret to just below the
0.00 mL mark with NaOH solution. Record the initial buret reading to the nearest 0.01 mL.
Pipet exactly 10.00 mL of the H2SO4 solution into a 250-mL Erlenmeyer flask. Add two
or three drops of phenolphthalein. Place a piece of white paper under the flask and add NaOH
from the buret, with swirling, until a faint pink color persists for at least 20 seconds. As you near
the endpoint, rinse down the inside walls of the flask with distilled water from your wash bottle.
Record the final buret reading to the nearest 0.01 mL.
Find the volume of NaOH used, and calculate the molarity of the H2SO4 solution.
Repeat the titration until you have two results that agree within 1%. Average these two
results. Then calculate the percent error in the average H2SO4 concentration.
B – Gravimetric Analysis
Pipet exactly 10.00 mL of the H2SO4 solution into a 150-mL beaker. Add about 5 mL of
3 M HCl from a graduated cylinder. Add a glass bead to the beaker to prevent the solution from
bumping when heated. Heat the solution to near-boiling on a hot plate. Measure 20.0 mL of 0.1
M BaCl2 solution into a 100-mL beaker using a graduated cylinder. Warm this solution on the
hot plate too. Using beaker tongs to hold the hot glassware, slowly pour the hot BaCl2 solution
into the 150-mL beaker containing the H2SO4 solution, with stirring. Use distilled water from
your wash bottle to rinse any solid from your stirring rod back into the beaker. Cover the 150mL beaker with a watch glass, and continue to heat it on the hot plate for 30 minutes. Allow the
mixture to slowly cool to room temperature. Use forceps to remove the glass bead. Rinse any
adhering precipitate from the bead with distilled water from your wash bottle.
Use pencil to write your name on the edge of a piece of filter paper. Fold it in quarters,
open it in a cone, and tear the corner off of the outside flap. Weigh the torn filter paper to the
nearest 0.0001 g. Obtain a long stemmed funnel, if available, and support it with a split stopper
and clamp over another beaker. Hold the filter paper firmly inside the funnel, and moisten it
with distilled water from your wash bottle so that it seals to the walls of the funnel. For fastest
filtration, the entire stem should fill with water, with no air bubbles.
Pour the BaSO4 suspension into the funnel, being careful to keep all of the liquid inside
the filter paper. Use distilled water to rinse all of the precipitate into the funnel. Use a few
CHEM 131
Experiment 1B
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additional mL of distilled water to rinse any excess BaCl2 from the filtered precipitate. Carefully
spread the wet filter paper with the BaSO4 precipitate on a piece of paper towel, and let it dry in
your locker until the next laboratory period.
If time permits, do a second trial.
Dispose of any solutions containing barium ions in the metal waste container. Other
waste can be flushed down the sink with plenty of water.
At the beginning of the next laboratory period, weigh the filter paper and precipitate to
the nearest 0.0001 g. Find the mass of precipitate, and from this mass calculate the molarity of
the H2SO4 solution. Calculate the percent error in this result.
Experiment 1B – Analysis of a H2SO4 Solution
Pre-lab Questions
An aqueous solution of Ba(OH)2 was analyzed by two different methods.
1. A 10.00-mL sample of the Ba(OH)2 solution was pipetted into a flask. An appropriate
indicator was added, and it was titrated with standardized hydrochloric acid. 27.90 mL of
0.1062 M HCl were required to reach the endpoint of the titration.
(a) Write a balanced equation for the acid-base reaction.
(b) Calculate the molarity of the original Ba(OH)2 solution from the volume of HCl used.
2. A 10.00-mL sample of the same Ba(OH)2 solution was pipetted into a beaker. An excess of
aqueous sodium phosphate was added. The resulting precipitate was filtered and dried. The
mass of the precipitate was 0.2974 g.
(a) Write a balanced equation for the precipitation reaction.
(b) Calculate the molarity of the Ba(OH)2 solution from the mass of the precipitate.
CHEM 131
Experiment 1B
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