Electrolytic Titration
Patricia A. Riley Lincoln Park High School
2001 N. Orchard St. Mall
Chicago IL 60614
(312) 534-8130 ext. 148
Objectives:
For students in sixth, seventh, and eighth grades, to show that:
1. Ions are present during this titration.
2. Acids and bases are electrolytes.
3. No ions are present at the end point of this titration.
4. A precipitate and water are made from ions in this reaction.
5. Electrical conductivity depends on the presence of ions in a solution.
Materials Needed:
For class demonstration:
buret conductivity tester with clamp
goggles and apron ringstand and buret clamp
phenolphthalein solution 1000-mL beaker
magnetic stirrer and stir bar saturated solution of Ba(OH)2
(or glass stirring rod) 1.0 M H2SO4 solution
At each team work station:
2 eyedroppers 3 dixie cups
saturated solution of Ba(OH)2 1.0 M H2SO4 solution
litmus paper pH paper
phenolphthalein solution simple conductivity tester
distilled water
Strategy:
1. Ahead of time
a. Prepare the two solutions needed as follows:
1) Saturated barium hydroxide solution
12 grams Ba(OH)2 + enough distilled water to make 800 mL of
solution
2) 1.0 M sulfuric acid solution
14 mL concentrated H2SO4 + enough distilled water to make 250 mL of
solution
b. Fill the buret with 50 mL of 1.0 M H2SO4 and position the buret in the
clamp attached to the ringstand.
c. Place the magnetic stirrer with a 1000-mL beaker on it under the buret.
d. Attach the conductivity tester to the ringstand and position its
electrodes to hang inside the beaker. Place the stir bar in the
beaker. If a magnetic stirrer is not available, use a glass stirring
rod to manually stir the beaker's contents.
e. Fill the beaker with enough of the Ba(OH)2 solution to submerge
partially the conductivity tester's electrodes.
f. Position the buret above the beaker.
g. Draw a labeled diagram of the apparatus on the chalk board.
2. At the start of class draw attention to the apparatus and advise the
students to observe closely everything that happens. Plug in the
conductivity tester and turn on the magnetic stirrer. Ask a student to
describe what he or she notices.
a. Add 3 to 5 drops of phenolphthalein to the beaker. What happens? Ask
for suggestions as to why this occurs.
b. Open the buret stopcock and adjust the flow to a steady drip of
individual drops. What changes now occur? Why might they be occurring?
Why does the beaker grow cloudy and its color change? Why does the
light get dimmer? Why does the light go out and then come back on? Why
was the phenolphthalein added? What is the precipitate that forms?
Lead the discussion into the concepts of acids, bases, neutralization,
indicators, and electrolytes.
3. How can the answers to the above questions be checked? Suggest that
experiments be done. Divide the class into teams of 3 or 4 students and
give each team a handout with directions for experiments to do. These
include the following:
a. Using litmus paper and pH paper, test the Ba(OH)2 and H2SO4 solutions
and distilled water. What does this suggest? Now use the student
conductivity tester on them. What does this reveal?
b. Use an eyedropper to put 3 eyedroppersful of Ba(OH)2 into a dixie cup
and add a drop of phenolphthalein to it. Test the contents with the
conductivity tester. What happens? Why? Rinse the electrodes of the
tester in distilled water. Save this cup and its contents for part 3d.
c. Repeat step 3b with the H2SO4. Be sure to rinse the tester in
distilled water to clean it.
d. To the dixie cup from step 3b, add fresh H2SO4 (Don't use the H2SO4 from
step 3c.) drop by drop with an eyedropper, while checking constantly
with the conductivity tester. Stop when the light goes out. Check the
cup at this point with litmus and pH papers. What are the results?
What do they mean? What about the color of the solution? Clean the
tester with distilled water.
e. Students return to their seats with their results.
4. Discuss the results. Be sure the students understand the following points:
a. Litmus paper is red in acids, blue in bases. Acids have pH's less than
7, bases greater than 7. H2SO4 is thus an acid; Ba(OH)2 a base.
b. Only solutions that contain ions conduct electrical current. Therefore,
Ba(OH)2 and H2SO4 solutions must contain ions, but distilled water does
not.
c. Phenolphthalein is pink in the presence of a base but colorless in the
presence of an acid. This suggests that the solution in the beaker
changes color just when no more base is present. Be sure to identify
this as the End Point and stress that there is no acid present either
when this happens.
d. The light grows dimmer as more and more of the ions of the base are
paired with ions from the acid, according to the following equation:
Ba+2 + 2 OH- + 2 H+ + SO4+2 = 2 H2O + BaSO4
BASE ACID WATER WHITE SOLID
SALT
At the End Point, there are only water and the white solid salt, BaSO4,
left in the beaker and NO IONS! The light goes out. As we continue to
add more H2SO4, we add more ions and so the light gradually comes back
on.
e. The solution turned cloudy as the white solid BaSO4 formed. This solid
is an example of a precipitate and is insoluble in water. Its ions are
locked up and not able to carry electrical current.
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