Chemistry 113.1
Introduction to Chemical Techniques
Experiments 9/10. Standard Solutions
(June 2012)
I. INTRODUCTION
A standard solution is a solution of known concentration that is prepared using exacting
techniques to ensure that the molarity is known to high accuracy. A stock solution is a large
volume of a common reagent prepared to perhaps 1 or 2 significant figures of accuracy. The
approximately 0.10 M NaOH and 0.020 M potassium permanganate solutions are such stock
solutions. You will determine their more exact concentrations by performing a standardization
for each. The reason this is necessary is that many common reagents decay over time or are
hygroscopic and collect additional water. Thus, in many common reagents, the concentration
changes as a function of time. For this reason, most standard solutions are prepared when
needed and only kept a short period of time (i.e., one to two weeks) before being re-standardized.
Once a solution is standardized, then this solution can be used to quantitatively measure the
amount of chemicals in other solutions, as long as these chemicals will react with the standard.
The first method for preparing standard solutions is to use solids that are very stable in air.
Because of the high stability, these salts can be weighed on a balance without fear that the salt
will gain or lose mass (because of oxidation by the oxygen in air or because of the absorbtion or
loss of water from/to the air). The second method for preparing standard solutions is to prepare a
solution and then react this solution with a separate standard solution.
In this experiment, you will use the first method to prepare standard solutions of potassium
hydrogen phthalate (Week 1) and ferrous ammonium sulfate hexahydrate (Week 2). You will
then use the acid/base reaction between potassium hydrogen phthalate and sodium hydroxide to
standardize the approximately 0.10 M sodium hydroxide stock solution (Week 1), and the
oxidation/reduction reaction between ferrous ammonium sulfate and potassium permanganate to
standardize the approximately 0.020 M potassium permanganate stock solution (Week 2). The
reason why solid sodium hydroxide cannot be used to prepare a standard solution via the first
method is that sodium hydroxide is extremely hygroscopic and, therefore, readily absorbs water
from the air. It also reacts with carbon dioxide in the air and cannot be stored for significant
periods of time. Solid potassium permanganate cannot be used to prepare a standard solution
because it is a strong oxidizer that reacts with any trace organic compounds in the deionized
water. Thus, since the concentration of the permanganate ion will change during these side
reactions, the final amount of permanganate in solution will not be known. Thus, stock solutions
of sodium hydroxide solutions and potassium permanganate solutions must be standardized
using the second standardization method.
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Chemistry 113.1
Introduction to Chemical Techniques
Experiments 9/10. Standard Solutions
(June 2012)
EXPERIMENT 9A. PREPARATION OF A STANDARD POTASSIUM HYDROGEN
PHTHALATE (KHP) SOLUTION.
Potassium hydrogen phthalate (KHP, KHC8H4O4) is a monoprotic acid that is inexpensive, is
available in high purity, and is easily dried. The high molar mass of this compound allows for
high precision when determining the number of moles in a solution of low volume (i.e., < 100
mL). Thus, KHP is commonly used to create a standard acidic solution using mass and
volumetric flasks.
1. Obtain a 100 mL volumetric flask from the stockroom. A volumetric flask has been
standardized and gives a volume that is accurate to four significant figures if the flask is
used at room temperature. The flask contains the stated volume (i.e., 100.0 mL) when it
is filled so that the bottom of the meniscus coincides with the line etched around the neck
of the flask. Rinse the flask with 25 mL of distilled water. If droplets form on the inner
surfaces, wash the flask with liquid detergent, rinse thoroughly with tap water, and then
rinse with three 25 mL aliqouts of distilled water.
2. Weigh 2 grams (1.9 to 2.1 g) of KHP, using a tared, clean dry beaker. Record the exact
weight of KHP in your notebook.
3. Add 20 mL of water to the beaker to begin to dissolve the KHP. Then transfer the KHP
to the volumetric flask, being careful not to spill any of the solution.
4. Rinse the beaker with three aliquots of 10 mL of distilled water. Add these washings to
the volumetric flask (again being careful not to spill any of the solution.) This ensures
that any KHP adhering to the walls of the beaker is transferred to the volumetric flask.
5. At this point, approximately 60 mL of distilled water should be present in the volumetric
flask (i.e., the water level should be about 3/4 of the way up the bulb of the volumetric
flask). Close the flask with the lid (or a rubber stopper or parafilm) and invert several
times to mix the contents. Continue this process until such time as the KHP is
completely dissolved.
6. Remove the lid (or rubber stopper or parafilm) and rinse with a small amount of distilled
water, again making sure that the water drops into the volumetric flask.
7. Add enough distilled water to nearly fill the flask to the etched line, cover, and invert
several times to mix the contents. Remove the flask cover, again rinsing to ensure that all
of the solution remains in the volumetric flask.
8. Add water with a clean medicine dropper (or pipette) to bring the bottom of the meniscus
to the etched line. Show your laboratory instructor your volumetric flask once you have
completed the preparation of the solution. If you have overfilled your flask, you will
need to begin again !
9. The laboratory instructor must sign the laboratory notebook next to the recorded volume
to verify that the volumetric flask was filled correctly.
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Chemistry 113.1
Introduction to Chemical Techniques
Experiments 9/10. Standard Solutions
(June 2012)
EXPERIMENT 9B. STANDARDIZATION OF SODIUM HYDROXIDE SOLUTION.
PERFORM THIS EXPERIMENT IN PAIRS. EACH PAIR SHOULD DO TWO TRIALS, WITH EACH
STUDENT IN THE PAIR DOING ONE OF THE TRIALS. RECORD IN THE NOTEBOOK WHICH TRIAL
IS DONE BY WHICH STUDENT. EVERY STUDENT SHOULD COMPLETE A SINGLE TITRATION IN
30 MINUTES OR LESS. IF TIME ALLOWS, PERFORM A THIRD TITRATION.
Acid/base reactions are neutralization reactions that generate a salt and (sometimes) water. For
instance, the reaction of the insoluble base Mg(OH)2 (s) with the strong aqueous acid HCl (aq)
gives water and the aqueous salt magnesium chloride. On the other hand, the reaction of aqueous
ammonia with hydrochloric acid gives the aqueous salt ammonium chloride (but no water). One
method to accurately investigate acid/base neutralization reactions is a titration.
Titrations are the precise measurement of the volume of one reactant (of known molarity)
required to completely react with another reactant (of unknown molarity) in a chemical reaction.
When the reaction is complete, the titration is said to have reached the end point, or equivalence
point, of the titration. Thus, titrations can be used to determine accurate concentrations of a
reactant in solution for many chemical reactions. In an acid/base titration, either a base of known
molarity is added to an acid of unknown concentration or an acid of known molarity is added to a
base of unknown concentration. The end point occurs when complete neutralization has
occurred, where neutralization implies that the acid has reacted with the base to generate water
and a salt. For example, the neutralization reaction of sodium hydroxide with potassium
hydrogen phthalate yields the aqueous salt sodium potassium phthalate and water, or
NaOH (aq) + KHP (aq) → KNaP (aq) + H2O (l) ,
where P stands for the phthalate ion. One way of determining the end point of an acid/base
titration is to use the weak acid phenolphthalein as an indicator. Phenolphthalein is colorless in
acidic solutions, a very pale pink in nearly neutral solutions, and fuchsia in basic solutions. In
this experiment, you will standardize the 0.1 M stock solution that your class made during
Experiment 11. This now standard basic solution will be used in Experiment 13A.
As an example of how to determine the molarity of a solution from a titration, let us assume that
28.50 mL of a freshly prepared aqueous solution of sodium hydroxide is required to titrate 50.00
mL of 0.0922 M potassium hydrogen phthalate. The molarity of the sodium hydroxide is easily
determined using the balanced chemical equation given above since at the equivalance point the
moles of KHP (H+) is equal to the moles of NaOH (OH-). Therefore:
VKHP x MKHP = VNaOH x MNaOH
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Chemistry 113.1
Introduction to Chemical Techniques
Experiments 9/10. Standard Solutions
(June 2012)
[Note: Although strictly speaking the formula above calls for V in liters (L), any volume unit can
be used as long as the same unit is used for both, because the unit cancels out.]
Therefore:
50.00 mL x 0.0922 MKHP = 28.50 mL x MNaOH
This gives MNaOH = 0.162 M
1. Two burettes are setup on the bench for each pair of students. Record the graduations of
the burettes in your notebook and determine the precision of volume measurements that
can be made with these burettes. Show this information to your laboratory instructor.
2. Put a few mL of water in each of the burettes, clamp them in a vertical position using a
ring stand and burette clamp and test the stopcock (i.e., the valve at the bottom of the
burette). You should be able to control the liquid flow to one drop at a time with
practice. If not, return the burette to the Stockroom and obtain another burette.
3. Put a few mL of distilled water into the burette and roll around to wet the inside walls of
the burette. If droplets form inside the burette, wash it with liquid detergent and a burette
brush (obtained from the Stockroom). Rinse with tap water (three 10 mL portions
allowing some of each portion to run through the tip) and then with three 10 mL aliquots
of distilled water.
4. Once the burette is clean, rinse the burette with three 10 mL portions of the solution that
the burette will hold. One burette will hold the approximately 0.10 M sodium hydroxide
solution [obtain 100 mL of ~0.10 M NaOH stock solution] and one will hold the standard
KHP solution. In each case, allow some of the solution to run through the tip. Discard
the rinsings in the appropriate waste bottle or in a waste beaker.
5. Make sure that the stopcock is closed. Hold the burette with its top below eye level and
fill it with solution to somewhere near the top mark. Run a little solution out of the
bottom and watch for air bubbles in the tip. (To remove a bubble, hold the partly filled
burette in a nearly horizontal position, open the stopcock part way and allow the slow
flow of liquid to push the bubble out of the tip.)
6. Clamp the filled burette in a vertical position and record the initial volume of liquid in
each burette. Always wait at least 15 seconds before taking a volume reading to ensure
that the liquid has drained down the inner surface of the burette.
7. Use the burette containing the standard KHP to place approximately 25 mL of the KHP
solution into a clean Erlenmeyer flask. Record the actual volume of KHP solution added
in your notebook. Add approximately 25 mL of distilled water and 2 drops of
phenolphthalein solution to the flask.
8. Place the Erlenmeyer flask on a sheet of white paper under the burette containing the
sodium hydroxide solution. (The white paper will help in seeing the end point.)
9. Begin adding sodium hydroxide to the KHP solution. Swirl the flask gently and steadily
to mix the acid and base solutions thoroughly. One can begin to add sodium hydroxide in
1 - 2 mL portions. Initially, a transient pink color will appear as sodium hydroxide is
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Chemistry 113.1
Introduction to Chemical Techniques
Experiments 9/10. Standard Solutions
(June 2012)
added. As the titration continues, the pink color will persist for longer intervals, but
continue to disappear. As this happens, begin to add sodium hydroxide in smaller
portions. The longer the pink color persists, the smaller the amount of sodium hydroxide
should be added. As you get closer to the equivalence point, sodium hydroxide should be
added dropwise.
10. The equivalence point is marked by the first permanent faint pink color in the KHP
solution. (Permanent is a color change that persists for at least 30 seconds. The pink
color may ultimately fade due to the absorption of CO2 from air.) If sodium hydroxide
has splashed onto the walls of the Erlenmeyer flask or any drop hangs from the burette
tip, when you are near the equivalence point, wash these down into solution using
distilled water from your wash bottle. Show your end point color (i.e. the flask) to your
instructor before continuing.
11. If your instructor indicates that you have overshot the equivalence point (i.e., the color of
the solution is dark pink or red/purple), add KHP solution from the acid burette dropwise
until the color disappears. Then add sodium hydroxide from the base burette dropwise
until you have reached the equivalence point. Be sure to record both the volume of KHP
added and the volume of sodium hydroxide added.
12. Obtain a new clean Erlenmeyer. Refill the burettes and repeat the titration with an
additional volume of KHP.
Once the two titrations have been completed, record your calculated value for the molarity (M)
of the sodium hydroxide solution in your laboratory notebook. The KHP solution can be
discarded in the appropriate waste container. The burettes should be emptied and rinsed with
three 10 mL portions of distilled water and returned to the Stockroom.
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Chemistry 113.1
Introduction to Chemical Techniques
Experiments 9/10. Standard Solutions
(June 2012)
EXPERIMENT 10A. PREPARATION OF A STANDARD FERROUS AMMONIUM
SULFATE (FAS) SOLUTION .
Ferrous ammonium sulfate hexahydrate (FAS, FeSO4(NH4)2SO4 • 6 H2O) is an iron(II) salt that
is more stable than iron(II) sulfate. Iron(II) is easily oxidized to iron(III) in the presence of
oxygen via
4 Fe2+ (aq) + O2 (g) + (4 + 2 x) H2O ↔ 2 Fe2O3 • x H2O (aq) + 8 H+ (aq) .
However, the ammonium ion in FAS is slightly acidic and, therefore, adds H+ to the solution.
The presence of H+ shifts the equilibrium to the left, stabilizing iron(II) in solution with respect
to oxidation. Similar to KHP, FAS has a high molar mass, which allows for high precision when
determining the number of moles in a solution of low volume. Thus, FAS is commonly used to
create standard solutions for oxidation/reduction reactions.
1. Obtain a 100.0 mL volumetric flask from the stockroom. Again, clean this flask
following the procedures given for the preparation of the KHP standard.
2. Weigh 4 grams (3.9 to 4.1 g) of FAS, using a tared, clean dry beaker. Record the exact
weight of FAS in your notebook.
3. Add 25 mL of distilled water to the beaker to begin to dissolve the FAS. Then transfer
the FAS to the volumetric flask, being careful not to spill any of the solution.
4. Rinse the beaker with two 10 mL aliquots of distilled water and a third 10 mL aliquot of
3M sulfuric acid to stabilize the solution. Add these washings to the volumetric flask
(again being careful not to spill any of the solution.) This ensures that any FAS adhering
to the walls of the beaker is transferred to the volumetric flask.
5. Close the flask with the lid (or a rubber stopper or parafilm) and invert several times to
mix the contents. Continue this process until such time as the FAS is completely
dissolved.
6. Remove the lid (or rubber stopper or parafilm) and rinse with a small amount of distilled
water, again making sure that the water drops into the volumetric flask.
7. Add enough distilled water to nearly fill the flask to the etched line, cover, and invert
several times to mix the contents. Remove the flask cover, again rinsing to ensure that all
of the solution remains in the volumetric flask.
8. Add water with a clean medicine dropper (or pipette) to bring the bottom of the meniscus
to the etched line. Show your laboratory instructor your volumetric flask once you have
completed the preparation of the solution. If you have overfilled your flask, you will
need to begin again !
9. The laboratory instructor must sign the laboratory notebook next to the recorded volume
to verify that the volumetric flask was filled correctly.
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Chemistry 113.1
Introduction to Chemical Techniques
Experiments 9/10. Standard Solutions
(June 2012)
EXPERIMENT 10B. STANDARDIZATION OF POTASSIUM PERMANGANATE.
PERFORM THIS EXPERIMENT IN PAIRS. EACH PAIR SHOULD DO TWO TRIALS, WITH EACH
STUDENT IN THE PAIR DOING ONE OF THE TRIALS. RECORD IN THE NOTEBOOK WHICH TRIAL
IS DONE BY WHICH STUDENT. EVERY STUDENT SHOULD COMPLETE A SINGLE TITRATION IN
30 MINUTES OR LESS. IF TIME ALLOWS PERFORM A THIRD TITRATION.
Oxidation/reduction reactions are chemical reactions that lead to a change in the oxidation
numbers of species in the reaction. In this experiment, you will perform a redox titration to
standardize the potassium permanganate solution. In this titration, aqueous iron(II) reacts with
the aqueous permanganate ion (dark red in water) in an acidic solution to give iron(III) and
aqueous manganese(II). You should balance this chemical reaction before coming to the
laboratory.
1. Carefully clean the two burettes and rinse several times with distilled water.
2. Once the burette is clean, rinse the burette with three 10 mL portions of the solution that
the burette will hold (one burette will hold the approximately 0.02 M stock potassium
permanganate solution [obtain about 100 mL ] to be standardized, and one will hold the
standard FAS solution), again allowing some of the solution to run through the tip.
Discard the rinsings in the appropriate waste bottle. Note: At least one of your
burettes will have white numbering and markers. This is the burette that must be
used for potassium permanganate.
3. Make sure that the stopcock is closed. Hold the burette with its top below eye level and
fill it with solution to somewhere near the top mark. Run a little solution out of the
bottom and watch for air bubbles in the tip. (To remove a bubble, hold the partly filled
burette in a nearly horizontal position, open the stopcock part way and allow the slow
flow of liquid to push the bubble out of the tip.)
4. Clamp the filled burette in a vertical position and record the initial volume of liquid in
each burette. Always wait at least 15 seconds before taking a volume reading to ensure
that the liquid has drained down the inner surface of the burette.
5. Use the burette containing the standard FAS solution to place approximately 25 mL of
the FAS solution into a clean Erlenmeyer flask. Record the actual volume of FAS
solution used in your laboratory notebook. Carefully and slowly add approximately 15
mL of 3 M sulfuric acid to the 25 mL FAS sample in the Ehrlenmeyer flask.. This acid
provides the H+ ions necessary for the oxdidation/reduction reaction.
6. Place the Erlenmeyer flask on a sheet of white paper under the burette containing the
potassium permanganate solution. (The white paper will help in seeing the end point.)
7. Begin adding potassium permanganate to the FAS solution, Swirling the flask gently and
steadily to mix the solution thoroughly. One can begin to add potassium permanganate in
1 - 2 mL portions. Initially, transient red streaks will appear as potassium permanganate
is added and reacts with iron(II). As the titration continues, the red color will persist for
longer intervals, but continue to disappear. As this happens, begin to add potassium
permanganate in smaller portions. The longer the red color persists, the smaller the
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Chemistry 113.1
Introduction to Chemical Techniques
Experiments 9/10. Standard Solutions
(June 2012)
amount of potassium permanganate should be added. As you get closer to the
equivalence point, potassium permanganate should be added dropwise.
13. The equivalence point is marked by the first permanent faint pink/red color in the FAS
solution. (Permanent is a color change that persists for at least 30 seconds.) If potassium
permanganate has splashed onto the walls of the Erlenmeyer flask or any drop hangs
from the burette tip, when you are near the equivalence point, wash these down into
solution using distilled water from your wash bottle. Show your end point color (i.e. the
flask) to your instructor before continuing.
8. If your instructor indicates that you have overshot the equivalence point (i.e., the color of
the solution is dark pink or red/purple), add FAS solution from the FAS burette dropwise
until the color disappears. Then add potassium permanganate from the other burette
dropwise until you have reached the equivalence point. Be sure to record both the
volume of FAS added and the volume of potassium permanganate added.
9. Obtain a new clean Erlenmeyer. Refill the burettes and repeat the titration with an
additional volume of FAS.
Once the two titrations have been completed, record your calculated value for the molarity (M)
of the potassium permanganate solution in your laboratory notebook.
The FAS solution can be discarded in the appropriate waste container. The burettes can also be
emptied and rinsed with three 10 mL portions of distilled water. Both volumetric flasks should
be cleaned with distilled water and returned to the Stockroom.
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Chemistry 113.1
Introduction to Chemical Techniques
Experiments 9/10. Standard Solutions
(June 2012)
HELPFUL TABLES:
The tables below no longer have table captions. You can still use these tables in your manual.
However, you will need to write the table captions for each table on your own.
Experiment 9A.
Molar mass of KHP:
Mass of KHP:
Volume of water:
Molarity of KHP:
Experiment 10A:
Molar mass of FAS:
Mass of FAS:
Volume of water:
Molarity of FAS:
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Chemistry 113.1
Introduction to Chemical Techniques
Experiments 9/10. Standard Solutions
(June 2012)
Experiment 9B:
Initial acid burette reading:
Final acid burette reading:
Volume of acid in flask:
Concentration of acid:
Initial base burette reading:
Final base burette reading:
Volume of base required for endpoint:
Concentration of base
(Note: Before the calculations in your
notebook to answer this.)
Concentration
Concentration
Concentration
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Concentration
Chemistry 113.1
Introduction to Chemical Techniques
Experiments 9/10. Standard Solutions
(June 2012)
Experiment 10B
Initial FAS burette reading:
Final FAS burette reading:
Volume of FAS in flask:
Concentration of FAS:
Initial permanganate burette reading:
Final permanganate burette reading:
Volume of permanganate required for
endpoint:
Concentration of permanganate
(Note: Before the calculations in your
notebook to answer this.)
Concentration
Concentration
Concentration
11
Concentration
Chemistry 113.1
Introduction to Chemical Techniques
Experiments 9/10. Standard Solutions
(June 2012)
POST-LABORATORY QUESTIONS
1. What is the difference between a stock solution and a standard solution? Why is there a
distinction between the two? (Hint: Think about the importance of precision when
performing a calculation.)
2. How would the volume of base needed to perform the KHP titration have changed if a
solution of 0.100 M barium hydroxide was used instead of the sodium hydroxide
solution?
3. What is the average concentration of the sodium hydroxide solution? What sources of
error could cause variation in your 2 (or 3) determinations of the concentration of sodium
hydroxide ?
4. Could a solution of iron(III) be titrated with potassium permanganate? Explain.
5. What is the average concentration of the potassium permanganate solution?
6. What sources of error would cause variation in your 2 (or 3) determinations of the
concentration of potassium permanganate ?
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