Acid Base Chemistry
Electrical Conductivity as an Endpoint Detector in an Acid-Base Reaction
Vernier Software & Technology; “Chemistry with Computers”
13979 Millikan Way, Beaverton, OR 97005-2886
http://faculty.une.edu/cas/jvesenka/modeling/technology/vernier/sensorspdf/cond.pdf
Adapted for the Mobile Chemistry Laboratory by Barbara B. Bunn, Mike R. Johnson, and Hunter Clayton.(MCL Staff),
Kelly Rowland, Kelly Randall, Erin Hopkins, Tara Setliff In Mrs. Teresa Setliff’s Chemistry Class, Bassett High
School, Bassett, VA and Mrs. Mary Lou Hearn’s Chemistry Class, Nottoway High School, Nottoway Court House, VA
Introduction:
The Arrhenius theory of acids and bases states that an acid produces hydrogen ions (H+) in water and a base
produces hydroxide ions (OH-) in water. Strong acids and bases are considered to be completely ionized
(dissociated), that is, the reaction lies far to the right and there is virtually no starting material left. For example:
where (aq) means dissolved in water.
Solutions that are ionized are capable of conducting a current. If the substance is completely ionized, it is called a
“strong electrolyte”. Some substances, such as aqueous acetic acid, are only partially ionized and are called “weak
electrolytes”. A “non-electrolyte” dissolves in water as a molecule and does not conduct a current. Sugar is one
example of a nonelectrolyte. Non-electrolytes are generally molecules that are made up of nonmetal elements.
The process of an acid base titration is normally monitored with respect to the pH of the solution. The pH scale runs
from 1 to 14 where 1-7 is acidic,7 is neutral and 7-14 is basic. At the equivalence point, the moles of acid equal the
moles of base, and the solution is neutral with a pH of 7. Recall water is a molecule, and when it is pure does not
conduct a current, and also has a pH of 7.
The titration of barium hydroxide with sulfuric acid can be studied using a conductivity probe. The equation shows
the reaction that takes place. Since the products consist of an insoluble precipitate of BaSO4 (Ksp = 1.1 X 10-10) and
the molecule water, we can find the equivalence point with the conductivity probe.
At the equivalence point only two species remain. The reactants sulfuric acid (H2SO4) and barium hydroxide
(Ba(OH)2, which are both strong electrolytes have been consumed. Since there are virtually no electrolytes in
solution at the equivalence point, the conductivity should be zero.
The end point is defined as the point in a titration where the indicator changes color; the equivalence point is the
point in a titration where the moles of acid equal the moles of base. Because we are using conductivity as the
indicator, the end point is the same as the equivalence point.
In this experiment, we will monitor the conductivity during the reaction between H2SO4 with a known concentration
and an unknown concentration of Ba(OH)2. The equivalence point (which in this case is the same as the endpoint)
gives the number of moles of Ba(OH)2, and since we know the volume of Ba(OH)2, used, we can determine the
concentration (molarity) of the barium hydroxide solution. We can also see the effect that ions and their
concentration have on conductivity.
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Endpoint Determination by Conductivity
Acid Base Chemistry
As H2SO4 is slowly added to Ba(OH)2 of unknown concentration, changes in the conductivity of the solution are
monitored using a Conductivity Probe. When the probe is placed in a solution that contains ions, and thus has the
ability to conduct electricity, an electrical circuit is completed across the electrodes that are located on either side of
the hole near the bottom of the probe body. This results in a conductivity value that can be read by the interface.
The unit of conductivity used in this experiment is the microsiemens, or µS.
Goals:
• To demonstrate an acid-base titration using conductivity as the endpoint detector
• To introduce the concepts of endpoint and equivalence point of a titration
• To demonstrate differences between strong electrolytes non electrolytes
Prelab Questions:
1. What would you predict the conductivity reading would be for the following
situations? High or Low; Increasing or Decreasing.
• When the Conductivity Probe is placed in Ba(OH)2, prior to the addition of H2SO4.
• As H2SO4 is slowly added, producing BaSO4 and H2O.
• When the moles of H2SO4 added equal the moles of BaSO4 originally present.
• As excess H2SO4 is added beyond the equivalence point.
2. A solution (50.00mL) of NaOH is titrated with 0.0213 M HCl. At the equivalence
point. The buret reads 27.55 mL. (Titration was begun with the buret reading 0.00.)
•
What volume of HCl was used in the titration?
•
How many moles of HCL were used to arrive at the equivalence point?
•
How many moles of NaOH were consumed to reach the equivalence point?
•
What is the molarity (concentration in moles/Liter) of the NaOH solution?
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Endpoint Determination by Conductivity
Acid Base Chemistry
Safety: Eye protection must be worn.
Figure 1
Set Up:
Procedure: Note: Steps 1, 2 and 3 may be done for you.
1.
Prepare the computer for data collection by opening the file Conductivity.
The vertical axis of the graph has conductivity scaled from 0 to 2000 µS. The
horizontal axis has volume scaled from 0 to 50 mL. Check to be sure the x
axis has units of mL. If not, double click on “volume” in the table window
and change the units to mL. Click ”okay”.
2. Measure out approximately 60 mL of 0.080 M H2SO4 into a 250-mL beaker.
Record the precise H2SO4 concentration in your data table. CAUTION:
H2SO4 is a strong acid, and should be handled with care.
3. Obtain a 50-mL buret and rinse the buret with a few mL of the H2SO4
solution. Use two utility clamps to attach the buret to the ring stand as shown
in Figure 1. Fill the buret a little above the 0.00-mL level. Drain a small
amount of H2SO4 solution so it fills the buret tip and leaves the H2SO4 at the
0.00-mL level of the buret. Dispose of the waste solution from this step as directed by the MCL staff.
Students: START HERE
4. Measure out exactly 50.0 mL of Ba(OH)2 of unknown concentration using a graduated cylinder. Place the
solution into a clean, dry 250-mL beaker. Then add approximately 50 mL of distilled water to the beaker.
CAUTION: Ba(OH)2 is toxic. Handle it with care.
5. Arrange the buret containing H2SO4, conductivity probe, beaker containing Ba(OH)2, and stirring bar as shown
in Figure 1. The Conductivity Probe should extend down into the Ba(OH)2 solution to just above the stirring
bar. Set the selection switch on the amplifier box of the probe to the 0-2000 µS range.
6. Before adding H2SO4 titrant, click Collect and monitor the displayed conductivity value (in µS). Once the
conductivity has stabilized, click Keep. In the edit box, type “0”, the current buret reading in mL. Press
ENTER to store the first data pair for this experiment.
7. You are now ready to begin the titration. This process goes faster if one person manipulates and reads the buret
while another person operates the computer and enters volumes.
a. Add 1.0 mL of 0.020 M H2SO4 to the beaker. When the conductivity stabilizes, again click Keep. In the
edit box, enter the amount of acid used as shown on the buret. Press ENTER. You have now saved the
second data pair for the experiment.
b. Continue adding 1.0 mL increments of H2SO4, each time entering the buret reading, until the conductivity
has dropped to or below 100 µS.
c. After the conductivity has dropped below 100 µS, add one 0.5-mL increment and enter the buret reading.
d. After this, use 2-drop increments (~0.1 mL) until the conductivity reaches 25 µS. Now you must be careful!
Add ½ to 1 drop until the conductivity reaches a minimum (Zero µS at the equivalence point). Enter the
volume after each addition. When you have passed the equivalence point, continue using 2-drop increments
until the conductivity is greater than 50 µS again.
e. Now use 1.0-mL increments until the conductivity reaches about 2000 µS, or 25 mL of H2SO4 solution has
been added, whichever comes first.
When
you have finished collecting data, click Stop.
8.
Printing Your Graph
9. Print a copy of the Graph window. Highlight the graph window. Click on the printer icon. Enter your name(s)
and the name of the experiment in the edit box. Click Okay; in the next box that appears, indicate the number
of copies (one for each person in your group) and click Okay.
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Endpoint Determination by Conductivity
Acid Base Chemistry
10. Print one copy of the table window to share with members of your group. Highlight the table window and
follow the steps in # 9.
PROCESSING THE DATA:
1. From the data table and graph that you printed, determine the volume of H2SO4 added at the equivalence point.
The graph should give you the approximate volume at this point. The precise volume of H2SO4 added can be
determined by further examination of the data table for the minimum voltage. Record the volume of H2SO4.
2. Calculate moles of H2SO4 added at the equivalence point. Use the molarity, M, of the H2SO4 and its volume, in
Liters, L.
3. Calculate the moles of Ba(OH)2 at the equivalence point. Use your answer in the previous step and the ratio of
–
moles of Ba(OH)2 and H2SO4 in the balanced equation (or use the 1:1 ratio of moles of H+ to moles of OH
from the equation).
4. From the moles and volume of Ba(OH)2, calculate the concentration of Ba(OH)2, in mol/L.
Data Table
Molarity of H2SO4
Volume of H2SO4 , closest to zero point
Moles of H2SO4
Volume of Ba(OH)2
Moles of Ba(OH)2
0.080 M
Post Lab Exercises
1. If you used HCl to titrate NaOH to the equivalence point, would the conductance be zero?
2. Why or why not?
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Endpoint Determination by Conductivity