AP Chemistry
Name: ___________________________
Group Members:
___________________________
___________________________
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Standardization of Acid and
Base Solutions
Purpose and Objectives
In this lab you will standardize an acid solution by use of a primary standard. The standardized acid will then
be uses to determine the concentration of a base. The neutralization reaction involved in the standardization
of the acid solution is:
Hydrochloric acid and sodium hydroxide are the most common strong acids and bases used in the laboratory.
It is impossible to purchase completely pure HCl or NaOH to be used in making solutions, so solutions of
approximate concentrations are made, then standardized to learn their exact concentrations. In this
experiment, you will prepare Stock solutions of hydrochloric acid and standardize it by titration. This
standardized solution will be used to analyze unknown solution of sodium hydroxide. HCl is standardized by
reacting with a known mass of sodium carbonate:
2HCl (aq) + Na2CO3(aq)  2NaCl (aq) + H2CO3 (aq)
Sodium carbonate (Na2CO3) is a primary standard, meaning it can readily be purchased in very pure form and
weighed out in such a way that you know to the greatest extent possible exactly how much material you have.
However, there are two considerations you must keep in mind when using sodium carbonate. First, the solid
will rapidly absorb water from the air. It is possible to remove the water by heating the solid in an oven for
several hours (this will be done for you before you arrive in lab), but care must be taken to store the material
in a desiccator after it has been dried. Second, the reaction above generates CO2, which dissolves into the
solution to generate an acid. The presence of dissolved CO2 thus interferes with the pH and the detection of
the end point of the titration. However, the CO2 can be driven off by boiling the solution, enabling an
accurate titration
H2CO3 (aq)  H2O (l) + CO2 (g)
The neutralization reaction involved in standardization of the base solution is:
NaOH (aq) + HCl (aq)  NaCl(aq) + H2O (l)
AP Chemistry
Safety and Disposal
Normal laboratory precautions, including wearing goggles and aprons at all times, should be taken to protect
you at all times. Use of heated liquids can cause burns. Use hot hands to transfer hot liquids and take care
when transferring liquid from one container to another. Research the chemicals used in this lab and how
to dispose of them.
Equipment
Read through the procedures and make a list of the equipment you will need.
Procedure
Part I – Preparation of Burets
1. Wash the buret with soap and water.
2. Rinse thoroughly with tap water.
3. Install the stopcock.
a. Apply grease and rotate to distribute the grease.
b. Secure the stopcock with a clip
c. Make certain that the stopcock is closed and that there is a beaker below buret to catch any
leaks.
4. Rinse 3 times with 5 mL portions of distilled water.
5. Rinse 2 times with 5 mL portions of titrant (HCl)
6. Fill buret with titrant, making certain that there are no air bubbles and that the tip is filled with
titrant.
Part II – Titration of HCl Solution
1. Weigh out about 0.25 g of anhydrous Na2CO3 to the nearest mg. (rapidly absorbs atmospheric water.
If not Record the exact mass of your sample.
2. Dissolve the Na2CO3 in about 25 mL of distilled water. Add 2-3 drops of Methyl red indicator.
3. Record the initial volume of acid in your buret. Titrated the solution with the HCl until the solution
color just begins to change from yellow to pink.
4. Bring the solution to a boil over a burner or hot plate. The original yellow solution color should
return. (if it does not, too much acid has been added and the titration should be repeated with a new
sample of Na2CO3)
5. Continue to add acid one drop at a time until the color changes to pink with the addition of just one
drop. Record the final volume of acid in the buret.
6. Repeat steps 1-5 with a second sample of Na2CO3.
Part III – Titration of NaOH Solution
1. Obtain at least 30 mL of NaOH solution.
AP Chemistry
2. Using a volumetric pipet, transfer precisely 10.00 mL of the NaOH solution to an Erlenmeyer flask
and add 2-3 drops of phenolphthalein indicator. Record this volume accurately.
3. Titrate the base solution with the standardized HCl from part I until a color change from pink to
colorless is observed. Record your initial and final volumes accurately.
4. Repeat steps 2-3 with a second aliquot of NaOH solution.
Data
You will need to construct data tables to display measured and calculated values.
Data Analysis
1. Show a sample calculation for the molarity of HCl from titration with the primary standard Na2CO3
2. Sketch a titration graph (pH vs. volume of titrant added) of what you believe the titration would look
like if graphed continuously using sensor.
3. Show a sample calculation for the molarity of NaOH from the titration with the standard HCl
solution.
4. Sketch a titration graph (pH vs. volume of titrant added) of what you believe the titration would look
like if graphed continuously using sensor.
5. Using the mass of sodium hydroxide added to the 50-ml volumetric flask, what is the expected
molarity of the sodium hydroxide solution prepared in this experiment?
6. How do you account for the discrepancy between your experiment value from the titration and the
value calculated from the mass of sodium hydroxide used to prepare the sodium hydroxide solution?
Conclusion
You will need to write a proper conclusion relating what you saw in the lab and how it relates to the topics
covered in class. Be specific in your relation of observations to the results. Be sure to include any possible
sources of error with a discussion of how that error could have affected the results.