EXPERIMENT 15
USING CONDUCTIVITY TO LOOK AT SOLUTIONS:
DO WE HAVE CHARGED IONS OR NEUTRAL MOLECULES?
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GOAL
After you complete this experiment, you should have a better understanding of aqueous solutions and the
forms that solutes may be in: charged ions or neutral intact molecules. In the Preliminary Activity, you will
learn about measuring conductivity and what conductivity tells us about a solution. In Investigation 1,
you’ll choose a researchable question, design and conduct your experiment, and then share this with the
class. In Investigation 2, you will apply conductivity to a titration.
INTRODUCTION
Many substances dissolve in water. Some substances dissolve so that we get individual neutral molecules
surrounded by H2O molecules. Other substances break apart into ions so that we get separate charged ions
surrounded by H2O molecules. We can easily distinguish between the two in lab. Solutions that contain
many ions will conduct electricity well–we call these electrolytes or strong electrolytes. Solutions that
contain only neutral molecules will not conduct electricity–we call these non-electrolytes. Solutions that
contain only a few ions will conduct electricity weakly–we call these weak electrolytes.
Figure 1 below shows how O2 dissolves in water. Notice that in the pure sample shown at left, we have
neutral molecules composed of two oxygen atoms each. When O2 dissolves in water, each molecule
remains intact.
Water molecules keep the O2
molecules apart. Since everything in our beaker is
neutral, this solution will not conduct electricity and
will be a non-electrolyte. Most molecular substances,
that is, substances that exist as neutral molecules when
they are pure, will also exist as neutral molecules when
they dissolve in water.
Figure 2 below shows how NaCl dissolves in water.
Notice that in the pure sample shown at left, we have
charged ions: Na1+ and Cl1-. NaCl does not have
molecules. When NaCl dissolves in water, each ion goes
its separate way. In the solution, water molecules keep
the ions apart from each other. This solution contains
charged ions that will conduct electricity well, making
this solution a strong electrolyte. Ionic salts are
compounds like NaCl that are composed of charged ions.
You can recognize the formula of a salt by looking for a
metal paired with a non-metal or polyatomic ion.
H
H
H
O
Na1+
Cl1-
Cl1-
Na1+
H2O
O
H
Na1+
O
1+
Na
H
H
Cl1-
H
O
O
H
H
H
Cl1-
Occasionally, a substance that exists as neutral molecules
when pure will react with water when it dissolves. The
most common examples are acids and bases. HNO3,
nitric acid, is a good example. As shown in the figure
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below, pure HNO3 exists as neutral molecules. [Note that while the formula for HNO3 contains a negatively
charged polyatomic ion, it does not contain a positively charged metal ion or polyatomic ion. It is NOT a
salt like NaCl.] When HNO3 dissolves in water, however, the water molecules pull H1+ off of the acid,
leaving NO31- behind.
Ions form in solution and thus this solution will conduct electricity. HNO3 is an example of a strong acid,
so named because 100% of its molecules fall apart to ions and result in a strong electrolyte solution. We can
write this as
HNO3(l) H2O H1+(aq) + NO31-(aq)
Eqn 1
Unlike HNO3, most acids are considered weak acids because only a small percentage of their molecules
will be broken apart into ions by the water. H2S is a weak acid. Typically less than 1% of its molecules are
broken into ions when placed in water. Since this results in a solution with only a very few charged ions and
mostly neutral molecules, it will conduct electricity weakly and be known as a weak electrolyte. Since most
of the H2S molecules remain intact, not broken into ions, we most often represent its solution as intact
H2S(aq), not as H1+(aq) and HS1-(aq). We usually want to show how most of the solute is present rather than
emphasizing the few ions that form.
Of course, not everything dissolves in water. These insoluble substances may be either molecular or ionic,
but in either case we leave their formulas intact and include the appropriate state of matter. For example, we
might have CH4(g) or PbI2(s). Their presence will not cause a solution to conduct electricity.
Beware the confusing OH group! To predict conductivity, you need to decide if a substance is ionic or
covalent. If you see “OH” in a formula, think carefully. When OH1- is paired with a metal ion, as in NaOH
or Fe(OH)2, you have an ionic compound. When OH is attached to non-metals, as in HOH or CH3OH, you
have a covalent (molecular) compound.
In Investigation 2, you will use conductivity to monitor a titration. The reaction you will study is between
aqueous NaOH, sodium hydroxide, and aqueous HCl, hydrochloric acid.
NaOH (aq) + HCl (aq) → NaCl (aq) + H2O (l)
Eqn 2
At the start of your titration, your beaker will contain only HCl, so the conductivity you measure will allow
you to decide if HCl exists as ions or molecules when dissolved in water. As you add NaOH from the buret,
it will be immediately consumed by reaction with the HCl to form our reaction products. At the equivalence
point of the reaction, you will have only the products of Eqn 2 in your beaker. Thus, the conductivity at this
point will allow you to decide if the products are intact or broken apart into ions. As you continue past the
equivalence point, the excess NaOH will also be present in the beaker and influence the conductivity we
observe.
HAZARDS
The solutions from Investigation 1 include acids. Be careful not to spill these on your skin. Be careful not
to dilute or cross-contaminate solutions. Rinse your conductivity probe between samples and blot it dry
with a tissue. In Investigation 2 you will use NaOH, which is a strong base and HCl which is a strong acid.
Both are corrosive and severe irritants, so be careful when handling. The solutions used in this experiment
can go down the drain.
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LABORATORY OBSERVATIONS AND DATA
Your instructor will assign you a partner to work with in lab. Record your partner’s name in your lab
notebook. You will each write up your own lab report, however, so be sure that you both have a complete
set of notebook entries and lab data before leaving lab. As always, include what you do and what you
observe in your lab notes.
PRELIMINARY INVESTIGATION
Procedure:
1. Set the switch on the Conductivity Probe to the
0–20000 µS/cm conductivity range.
2. Connect the Conductivity Probe and the data-collection interface.
3. Rinse the electrode with distilled water and the blot it dry with a tissue.
Repeat this cleaning before each new solution. Zero the probe by
pressing Ctrl 0. You shouldn’t need to re-zero again unless you have
problems.
4. Determine the conductivity of a distilled water sample.
a. Add 30 mL of distilled water to a clean 50 mL beaker.
b. Place the tip of the electrode into the distilled water. The hole near
the tip of the probe should be completely covered by the water.
c. Start data collection.
d. Stop data collection after about 15 seconds.
e. Use the Statistics option to determine the mean conductivity value.
4. Determine the conductivity of a sodium chloride solution.
a.
b.
c.
d.
Add about 0.25 g of NaCl to the distilled water.
Stir to dissolve the NaCl.
Repeat b–e of Step 4 above to determine the conductivity of the NaCl solution.
When done, flush the NaCl solution down the drain with excess water.
Discussion:
Return to the prelab side of the room with your lab partner. Working as a pair, answer these questions in
your lab notebook. Consult the Introduction for help.
1. List 2 factors that affect the conductivity of a solution.
2. Define and give an example of strong electrolyte, weak electrolyte, and non-electrolyte.
3. Consult the List of Available Materials posted near the instructor’s bench. List 2 Researchable
Questions that you could explore using these materials.
Participate in the discussion led by your instructor. As directed by the instructor, choose a Researchable
Question from the list generated by the class.
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INVESTIGATION 1
1. Prepare your Research Plan in your notebook. The sections should be
a. Researchable Question
b. Hypothesis
c. Variables (Manipulated, Responding, and Controlled variable(s))
d. Materials List
e. Safety Issues
f. Procedure
g. Data Table
2. Get your plan approved by the instructor, make any needed changes, and then conduct your
Investigation 1.
3. After you complete Investigation 1, return to the prelab side of the room and complete an
Investigation Report form (available at the Instructor’s bench).
4. Prepare to share the key points orally with the rest of the class.
5. During the class discussion of Investigation 1, you will be asked to briefly share your
investigation with the rest of the class. You may modify or expand upon your conclusions based
upon class discussion. Take notes on the investigations of others in your lab notebook since the
lab report will include questions on the work of other groups. (All the Investigation Report forms
will be available for viewing on Inquire the day after you complete this experiment.)
INVESTIGATION 2: MONITORING A TITRATION BY CONDUCTIVITY
1.
Obtain about 100 mL of 0.100 M NaOH in a clean beaker
2.
Obtain a 50-mL buret and rinse the buret with a few milliliters of the NaOH solution. Repeat this
two more times. Fill the buret to above the 0.00-mL level of the buret. Drain a small amount of
NaOH solution so it fills the buret tip and leaves the NaOH at the 0.00-mL level of the buret.
Although we don’t usually worry about starting at exactly 0.00 mL, in this procedure doing so will
make data recording simpler.
3.
Use a clean 250-mL beaker to get 10.0 mL of HCl of unknown concentration from the dispenser.
Add 90 mL of distilled water to the beaker.
4.
Add a stir bar to your HCl beaker and position this on the stirring plate. Position the conductivity
probe so that it is in the solution but away from the spinning stir bar. Position the buret above your
beaker so that solution from the beaker won’t hit the conductivity probe. Set the selection switch on
the amplifier box of the probe to the 0-20000 μS range.
5.
Your computer should already be in the LoggerPro program, with the file “Exp 3 Electrolytes” in
the folder Roanoke Experiments open. The vertical axis of the graph has conductivity scaled from 0
to 2000 μS. The horizontal axis has volume scaled from 0 to 25 mL. Left click on the “1000” mark
on the y-axis and drag down until “4000” comes into view.
6.
Before adding NaOH 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 will do 2 trials—a more general one and a detailed one. You are now ready to begin the first
titration. This process goes faster if one person manipulates and reads the buret while another
person operates the computer and enters volumes.
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a.
8.
Add 1.0 mL of 0.100 M NaOH to the beaker. When the conductivity stabilizes, again click
“Keep.” In the edit box, type the current buret reading read to the nearest 0.00 mL. Press
ENTER. You have now saved the second data pair for the experiment.
b.
Continue adding 1.0 mL increments of NaOH. With each increment, let the conductivity
stabilize, click “Keep,” and enter the buret reading. Continue until you have added 25-mL
of NaOH.
c.
When you have finished collecting data, click “Stop.” Note approximately where your
equivalence point is (to the nearest 0.0-mL). This will be important in the next trial.
d.
Print a copy of the Graph window using the popup menu File, Print Graph. Enter your
name(s). Print a copy for each team member. Print a copy of the Table window using the
popup menu File, Print Data Table. Enter your name(s). Print a copy for each team
member
After printing your graphs and data tables from the first titration, you are now ready to begin your
second titration.
9.
Proceed as you did before, using a clean 250-mL beaker to get 10.0 mL of HCl of unknown
concentration from the dispenser. Add 90 mL of distilled water to the beaker.
10.
Refill the buret to the 0.00-mL with the NaOH solution you prepared in Step 1.
11.
Add a stir bar to your HCl beaker and position this on the stirring plate. Position the conductivity
probe so that it is in the solution but away from the spinning stir bar. Position the buret above your
beaker so that solution from the beaker won’t hit the conductivity probe. Set the selection switch on
the amplifier box of the probe to the 0-20000 μS range.
a.
b.
c.
e.
f.
g.
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Begin as you did in the first titration, clicking “Collect.” A box will come up and there will
be several saving options. Click “Erase and Continue” to begin a new trial. Monitor the
displayed conductivity value. 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.
Add 1.0-mL of 0.100 M NaOH to the beaker. When the conductivity stabilizes, again click
“Keep.” In the edit box, type the current buret reading read to the nearest 0.00 mL. Press
ENTER. You have now saved the second data pair for the experiment.
Continue adding NaOH in 1.00-mL increments until you are within 1.0-mL of the
equivalence point you located in your first titration.
After this, use 2-drop increments (~0.1 mL) until the minimum conductivity has been
reached at the equivalence point. Enter the buret reading after each 2-drop addition. When
you have passed the equivalence point, continue using 2-drop increments until you are 1.0mL past your equivalence point.
Now use 1.0-mL increments until 25 mL of NaOH solution have been added.
When you have finished collecting data, click “Stop.” Note approximately where your
endpoint is. Print a copy of the Graph window using the popup menu File, Print Graph.
Enter your name(s). Print a copy for each team member. Print a copy of the Table window
using the popup menu File, Print Data Table. Enter your name(s). Print a copy for each
team member
Flush all of your used and unused solutions down the drain with excess water. Rinse your buret
with distilled water and return it to the supply bench.
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RESULTS AND DISCUSSION
For Investigation 1, write several paragraphs that describe your investigation, including a summary of each
of the points in your research plan. Include a clean table that clearly shows your data, including units. Now
describe what you can conclude from your investigation. Make sure that you refer to your data as the basis
for your conclusions. If your data is inconclusive or if you had experimental difficulties, describe these.
For Investigation 2, examine your computer generated graph for the second titration. The equivalence point
of the titration is at the minimum conductivity. You can find the approximate volume from the graph. Now
examine the data table closely at this approximate volume. Find the lowest value for conductivity. The
volume at this data point is the volume of NaOH needed to reach equivalence.
Now prepare a data table based on your second titration with the following 6 rows: Volume of HCl;
Molarity of NaOH; Volume of NaOH; Moles NaOH; Moles HCl; and Molarity HCl. Note that the first two
rows, the volume of HCl and molarity of NaOH are given in the procedure. The volume of NaOH is the
equivalence volume that you just determined from your graph and data table. Calculate the moles of NaOH
from the volume and molarity of NaOH. Since our titration is a 1:1 reaction, the moles of HCl at
equivalence is equal to the moles of NaOH. Finally, calculate the molarity of the HCl solution from its
moles and volume. Be sure to include sample calculations in your report. Attach your computer printouts.
QUESTIONS
1. Write several paragraphs, summarizing things your class collectively learned about conductivity and
types of electrolytes from Investigation 1. Use examples from this experiment to illustrate your points.
Your writing must be clear, grammatically correct, and chemically correct. Consult the Investigation Report
Forms on Inquire to find the information you need.
(a) Find an investigation that compares several different compounds. Describe the experiment, what
the data showed, and what you can conclude from it.
(b) Find an investigation that compares solutions with different concentrations. Describe the
experiment, what the data showed, and what you can conclude from it.
(c) Find one more investigation that you found interesting. Describe the experiment, what the data
showed, and what you can conclude from it.
2. In Investigation 2, you studied the reaction NaOH (aq) + HCl (aq) → NaCl (aq) + H2O (l)
(a) Which of these four substances is actually in the beaker at the equivalence point? [See the
Introduction for help.] Are there ions present in the solution at the equivalence point? Why or why not?
If you have ions, give their formulas with charges.
(b) Prior to reaching the equivalence point, which of the reactants was in excess (not the limiting
reagent)? Does the conductivity prior to equivalence suggest that this exists as ions or intact molecules?
If you have ions, give their formulas with charges.
(c) After equivalence, the other reactant is in excess. Does the conductivity after the equivalence point
suggest that this reactant exists as ions or intact molecules?
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Investigation Report Form
Experiment Number:
Team Members:
Researchable Question:
Variables
Procedure Method Summary
Data
What was learned/concluded
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