Electrolyte and Nonelectrolyte Solutions Lab Report

Lab 7: Electrolyte and Nonelectrolyte SolutionsReport Form – for on campus students
Due Mon. Mar. 15 at 11:59 PM. Upload .pdf (preferred) or MS Word file to
Brightspace.
Group Member Names:
Part A
Question #1. Complete Report Table 1 below and classify each of the solutions as a stronger
electrolyte, weaker electrolyte, or nonelectrolyte.
Report Table 1 Conductivities and Electrolyte Classifications for Solutions Examined in Part A.
solution
(identity and concentration)
conductivity (𝝁S/cm)
Page 1 of 9
electrolyte
classification
Question #2. Compare the conductivity values of the solutions examined in Part A and explain
why the conductivities are different for each of the following pairs. Cite your data (i.e. refer to
your conductivity measurements) in your explanations.
•
deionized water and tap water
•
0.04 M acetic acid (CH3COOH) and 0.04 M hydrochloric acid (HCl)
•
0.04 M sodium hydroxide (NaOH) and 0.04 M ammonia (NH3)
•
0.04 M sucrose (C12H22O11) and 0.04 M acetic acid (CH3COOH)
Part B
Question #3. Calculate the sodium chloride (NaCl) concentration (in M) for each of the
solutions examined in Part B and enter the results in Report Table 2 below. Then, enter your
conductivity measurements in Report Table 2.
Report Table 2 Measured Conductivities and Electrolyte Classifications for Solutions Examined in Part B.
test tube #
conductivity, S/cm
[NaCl], M
1
2
3
4
5
Page 2 of 9
Calculation #1. Show a sample calculation (i.e., full work, including units) for the concentration
of NaCl in test tube #1.
Question #4. Graph your data from Question #3 (i.e., conductivity (μS/cm) versus NaCl
concentration (M)) using a scatter plot. (Conductivity should be on the y-axis and
concentration should be on the x-axis. If you need Excel help, please attend TA office hours.)
Label the axes and add a title. Paste a copy of your graph here.
Question #5. Summarize the effect of dilution on the conductivities of solutions. Describe any
numerical relationship between dilution and conductivity. Cite your data in your explanation
(i.e. refer to your conductivity measurements).
Page 3 of 9
Part C
Question #6. Write the balanced molecular equation for the reaction of HCl and NaOH.
Question #7. Write the balanced total ionic equation for the reaction of HCl and NaOH.
Question #8. Record your conductivity measurements in Report Table 3 below.
Report Table 3 Volumes of HCl Added and Measured Conductivities for Acid-Base Reaction Studied in Part C.
Total volume of 0.04 M HCl added (mL)
0 (initial)
Page 4 of 9
Conductivity (S/cm)
Question #9. Graph your data in Question #8 (conductivity (μS/cm) versus HCl volume (mL))
using a scatter plot. (Conductivity should be on the y-axis and concentration should be on the
x-axis.) You may connect the points with a line if you wish, but do not try to fit a line to your
data. Label the axes and add a title. Paste a copy of your graph here.
Under your graph, describe the trend shown.
Describe the trend shown in your graph above.
Calculation #2. Calculate the amount (moles) of each ion after the addition of 5 mL of HCl to
the NaOH solution and complete Report Table 4 below. See p. 7 of the lab for an example
calculation.
Total volume (mL) = ____________
Report Table 4 Moles of Ions Present in Solution After Addition of 5 mL of HCl to NaOH Solution.
NaOH(aq) +HCl(aq) → NaCl(aq) +H2O(l)
Amount (mol)
Na+
OH–
Initial
Addition
Reaction
Final
Page 5 of 9
H+
Cl–
Calculation #3. Show your work here, including units, for the total moles of ions in solution
after the addition of 5 mL of 0.04 M HCl solution.
Calculation #4. Show your work here, including units, for the concentration of ions in solution
(M) after the addition of 5 mL of 0.04 M HCl solution.
Calculation #5. Calculate the amount (moles) of each ion after the addition of 10 mL of HCl to
the NaOH solution and complete Report Table 5 below.
Total volume (mL) = ____________
Report Table 5 Moles of Ions Present in Solution After Addition of 10 mL of HCl to NaOH Solution.
NaOH(aq) +HCl(aq) → NaCl(aq) +H2O(l)
Amount (mol)
Na+
OH–
Initial
Addition
Reaction
Final
Page 6 of 9
H+
Cl–
Calculation #6. Calculate the amount (moles) of each ion after the addition of 15 mL of HCl to
the NaOH solution and complete Report Table 6 below.
Total volume (mL) = ____________
Report Table 6 Moles of Ions Present in Solution After Addition of 15 mL of HCl to NaOH Solution.
NaOH(aq) +HCl(aq) → NaCl(aq) +H2O(l)
Amount (mol)
Na+
OH–
H+
Cl–
Initial
Addition
Reaction
Final
Calculation #7. Calculate the amount (moles) of each ion after the addition of 20 mL of HCl to
the NaOH solution and complete Report Table 7 below.
Total volume (mL) = ____________
Report Table 7 Moles of Ions Present in Solution After Addition of 20 mL of HCl to NaOH Solution.
NaOH(aq) +HCl(aq) → NaCl(aq) +H2O(l)
Amount (mol)
Na+
OH–
Initial
Addition
Reaction
Final
Page 7 of 9
H+
Cl–
Calculation #8. Calculate the amount (moles) of each ion after the addition of 25 mL of HCl to
the NaOH solution and complete Report Table 8 below.
Total volume (mL) = ____________
Report Table 8 Moles of Ions Present in Solution After Addition of 25 mL of HCl to NaOH Solution.
NaOH(aq) +HCl(aq) → NaCl(aq) +H2O(l)
Amount (mol)
Na+
OH–
H+
Cl–
Initial
Addition
Reaction
Final
Question #10. Use the final moles of ions and volumes in Calculations #2, 5, 6, 7 and 8 to
calculate the total concentration of ions in solution after each addition of HCl. Enter your
results in Report Table 9. Transfer your conductivity data from Question #3 to Report Table 9.
Then, list the ions present in solution (after the reaction is complete) after each consecutive
addition of HCl solution.
Report Table 9 Ions Present, Total Ion Concentrations and Conductivities During Addition of HCl Solution to NaOH Solution.
total volume of HCl
solution added, mL
total concentration
of ions, M
Conductivity
(μS/cm)
0 (initial)
Page 8 of 9
ions present
in solution
Question #11. Consider your results in Question #10 and your graph of conductivity versus
volume of HCl added for the reaction of HCl with NaOH in Question #9. Explain the changes in
conductance in terms of the total ion concentration.
Question #12. Based on the ionic species present throughout the course of the acid-base
reaction, as listed in Question #10, answer the following question: at approximately what point
on your graph has all of the base (NaOH) been neutralized by the acid (HCl)? Explain your
reasoning.
Upload your completed report in MS Word or .pdf format to the Lab 7
assignment portal in Brightspace. It is due Mon. Mar. 15 by 11:59 PM.
Page 9 of 9
Tech_Lab 7: Electrolyte and Nonelectrolyte Solutions (CHM
11600 -Spring 2021)
Lab 7 (25 pts.) consists of the following:
1. A pre-lab quiz (found on Brightspace) (10 pts.)
2. A Lab Report (15 pts)
Record your group members’ names.
PRELAB PRACTICE QUESTIONS
As part of your individual preparation for lab, read the experiment and answer the following
questions. (Your answers will not be collected or graded.)
You will take a quiz on Brightspace related to these concepts.
1. What is/are the difference(s) between solutions that are “nonelectrolytes”, “weaker
electrolytes”, and “stronger electrolytes”? Explain your reasoning.
2. Which one of the following aqueous solutions would you expect to have the smallest
conductance: (a) 0.060 M NaCl, (b) 0.030 M NaCl, or (c) 0.015 M NaCl? Explain your
reasoning.
3. Which aqueous solution in each of the following pairs would you expect to have the highest
conductivity?
a. HF or HBr
b. NH3 or KOH
c. ethanol (CH3CH2OH) or tap water
4. Calculate the concentration of NaCl in a solution prepared by diluting 2 mL of 0.05 M NaCl
to a total volume of 10 mL with deionized (DI) water. What is the total concentration of
ions in the diluted solution?
5. Answer the following questions about the reaction of 5.00 mL of 0.040 M HCl with 45 mL of
0.013 M NaOH solution.
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
a. What ions are present in solution after the reaction is complete?
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b. Complete the reaction table below to find the moles of each ionic species in solution
when the reaction is complete.
NaOH(aq) +HCl(aq) → NaCl(aq) +H2O(l)
Amount (mol)
Na+
OH–
Initial
Addition
—
Reaction
—
H+
Cl–
—
—
—
Final
c. What is the total concentration of ions in solution after the reaction is complete? (Hint:
how many moles of each type of ion are present in solution? What is the total solution
volume?)
6. When equimolar amounts of HCl(aq) and NaOH(aq) are combined, what ions are present in
solution when the reaction is complete?
7. Write the total ionic and net ionic equations for the acid-base reaction you will perform in
this lab.
8. Before the conductivity probe can be used in this experiment, it must be calibrated. What
substance is used to calibrate the probe for no conductance (i.e., 0 μS)?
GOALS
The goals of this activity are to
•
•
•
measure the conductance of solutions that contain electrolytes and/or nonelectrolytes,
examine how the conductance of a solution changes as the concentration of ions, and the
identity of the ions present, changes, and
examine how conductance relates to chemical reactions.
LAB LEARNING OBJECTIVES
•
•
•
Classify aqueous solutions as being stronger electrolytes, weaker electrolytes, or
nonelectrolytes.
Calculate concentrations of ions in aqueous solutions containing stronger electrolytes.
Calculate concentrations of ions in aqueous solutions after an acid-base reaction has
occurred.
2
INTRODUCTION
Conduction of electricity requires movement of charged particles, such as ions. Substances that
produce ions when they dissolve in water (e.g., NaCl) can conduct electricity and are classified
as electrolytes. Substances that do not produce ions when dissolved in water (e.g., sucrose, or
“table sugar”) do not conduct electricity and are classified as nonelectrolytes.
Figure 1 Conductivity experiment using a light bulb to indicate conductance. (Image credit: BC Open Textbooks).
Electrolytic conduction is observed in solutions of strong and weak electrolytes, in molten salts
and in some ionic solids. Most electrolytes are ionic compounds. However, some non-ionic
compounds, such as acids and bases, form electrically conductive aqueous solutions as a result
of their reaction with water to produce ions. Conduction occurs in electrolytic conductors as
both positive ions and negative ions migrate toward electrodes (see Figure 1 above).
Electrolytic conduction involves a transport of ions from one part of the conductor to another.
The flow of current in an electrolytic conductor is accompanied by chemical changes at the
electrodes. Electrolytic conduction plays an important role in the function of electrochemical
cells, batteries, electrolysis and electroplating. A compound’s ability to ionize and conduct
electricity also plays a significant role in biological systems.
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Measuring Conductance
The conductance of an electrolytic or electronic conductor is the reciprocal of its resistance in
ohms. At one time, the unit of conductance was called mho (the reverse of ohm). This unit has
been renamed siemens after W. Siemens, a noted German physical scientist who did extensive
research into the behavior of electricity. Electrolyte solutions have conductance values that are
much lower than those of metals. Thus, conductance values of solutions are often reported in
units of microsiemens (1 μS = 10-6 S).
Solutions can be nonelectrolytes (nonconductors), weaker electrolytes (poor conductors) or
stronger electrolytes (good conductors). In Part A of the experiment, we will measure the
conductance of such solutions using a conductivity probe and examine how the conductance
changes as the concentration of ions, and the identities of the ions, changes. In Part B of the
experiment, we will examine the effect of dilution on solution conductivity.
When the conductivity probe is placed in a solution containing ions, an electrical circuit is
completed across the graphite electrodes that are on either side of the hole in the probe (see
Figure 2 below). A potential difference is applied to the two electrodes and results in a current
that is proportional to the conductance of the solution between the electrodes. The current
measured by the conductivity probe (µS/cm) is displayed on a LabQuest interface.
Though the conductivity probe measures conductance, we are often interested in the
conductivity of a solution. Conductivity, C, is found using the following formula:
C = G x kc
where G is the conductance, and kc is the cell constant, which depends on the area of the
electrodes and the distance between the two electrodes. Since the conductivity probe has a
cell constant of 1.0 cm–1, the conductivity and conductance have the same numerical value.
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Figure 2 Macroscopic (left) and atomic-scale (right) views of a conductivity probe.
Our conductivity probe uses an alternating current in order to prevent the complete ion
migration to the electrodes (a situation called polarization) and consumption of the solute by
reactions at the electrodes. With each half-cycle of the alternating current, the polarity (sign)
of the electrodes is reversed. This reverses the direction of the ion flow and reverses any
chemical reaction that may have occurred at the electrode in the previous half-cycle. Thus, the
solutions under study retain their identity and the electrodes are not contaminated by
reduction-oxidation reactions occurring on their surfaces.
Acid-Base Reactions
Strong acid – strong base reactions involve the reaction of ions in solution. Thus, conductivity
can be used to monitor the reaction progress. In Part C of this experiment, you will examine
the reaction of hydrochloric acid, HCl, a strong acid, with sodium hydroxide, NaOH, a strong
base, by measuring the conductance after the addition of consecutive aliquots of HCl solution
to the NaOH solution. You will explain the conductivity changes that occur during the reaction
using the identities and concentrations of ions in the combined solution. The molecular
equation (1), the total ionic equation (2), and the net ionic equation (3) for the reaction are
shown below.
HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l)
H+ (aq) + Cl- (aq) + Na+ (aq) + OH- (aq) → Na+ (aq) + Cl- (aq) + H2O (l)
H+ (aq) + OH- (aq) → H2O (l)
(1)
(2)
(3)
5
In Part C, you will need to identify the ions present in solution and calculate their
concentrations in order to explain the measured conductivity. As represented in the scheme
below, at the beginning, before you add HCl, you will have a solution of NaOH, which contains
Na+ and OH– ions (Figure 3; solution (1)). (Note: we will ignore the very small amounts of H3O+
and OH– present in the solution that are produced from the autodissociation of water.) When
you add the first aliquot of HCl, all of the HCl will react with NaOH to form NaCl and water. At
this point, there will be Na+, OH–, and Cl– ions (Figure 3; solution (2)) present in the solution.
(Note: not enough HCl has been added at this point to react with all of the NaOH, so OH – ions
will still be present in the solution.) When you add exactly enough HCl so that all of the NaOH
has reacted, there will be only Na+ and Cl– ions in solution (Figure 3; solution (3)). (In an acidbase titration, this is called the “equivalence point”.) As you continue to add HCl after all of the
NaOH has reacted, there will be H+, Na+, and Cl– ions in solution (Figure 3; solution (4)).
Figure 3 Solution compositions at several points during the addition of HCl to a solution containing NaOH.
Calculations
To explain your conductivity measurements, you must be able to calculate the total
concentration of ions in solution. In the example shown in the reaction table below, we start
with 0.10 mol of NaOH and add 0.02 mol of HCl in the first aliquot. All of the H+ ions produced
by HCl react with OH– ions present in solution. Not enough HCl is added to react with all of the
NaOH, so OH– ions remain in solution after the reaction is complete. The numbers of Na+ and
Cl– ions do not change because they are spectator ions and do not participate in the reaction.
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HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l)
Amount (mol) Na+ OHH+
ClInitial
0.10 0.10
Addition
+0.02 +0.02
Reaction
-0.02 -0.02
Final
0.10 0.08
0
0.02
For the example above, in order to calculate the concentration of ions in solution after the
reaction, you need to consider the total number of moles of ions and the total volume of the
solution. For example, if you started with 25 mL of NaOH solution and added 5 mL of HCl
solution, then your total volume of solution would be 30 mL (0.030 L). The total number of
moles of ions is 0.10 mol Na+ + 0.08 mol OH– + 0.02 mol Cl–.
𝑇𝑜𝑡𝑎𝑙 𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛 𝑜𝑜𝑓 𝑖𝑜𝑛𝑠 𝑖𝑛 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 =
(0.10 + 0.08 + 0.02 )𝑚𝑜𝑙
= 6.67 𝑀
0.030 𝐿
Now, say that you add a second 5 mL aliquot of HCl containing 0.02 mol HCl to the reaction
solution. As shown in the reaction table below, there would be a total of 0.20 moles of ions in
35 mL (0.035 L) of solution after the reaction. The total concentration of ions in solution after
the addition of the second aliquot of HCl is 0.20 mol/0.035 L = 5.71 M. Note that while the
total number of moles of ions in solution has not changed, the concentration decreases due
to the increase in volume.
NaOH (aq) + HCl (aq) → NaCl (aq) + H2O (l)
Amount (mol) Na+ OHH+
ClInitial
0.10 0.08
0.02
Addition
+0.02 +0.02
Reaction
-0.02 -0.02
Final
0.10 0.06
0
0.04
SAFETY
Wear goggles are all times while you are in the laboratory. Follow all instructions given by your
Graduate Instructor.
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GENERAL INFORMATION
You will work in groups of 4 for this experiment. Half of the groups should start with Part A and
the other half should start with Part B, and then switch. Your Graduate Instructor will direct
you which part to do first. All groups should do Part C last.
PROCEDURE
1.
Connect a Vernier conductivity probe to your iPad using the “Graphical” app (see
below).
Log in to the iPad and
Make sure to
Connect thd wireless
Blinking green
open the Graphical app turn on the
device and press DONE.
LED means
and select Sensor Data
wireless
You must select the device you are
Collection option.
conductivity
ID (listed on back of
connected.
probe. Red LED temperature probe) from
will blink.
those listed on the iPad.
Figure 4a
Figure 4b
Figure 4c
Figure 4d
You will perform a two-point calibration of the probe with air (0 μS) and a 10000 μS calibration
solution. To avoid spilling the jar of calibration solution, clamp and stabilize the conductivity
probe to a ring stand during the calibration process.
Note:
•
•
•
Use the calibration solution in the jar provided.
Do not transfer the calibration solution to another container.
Do not discard the calibration solution.
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2. Calibrate the conductivity probe.
•
On the bottom tool bar menu bar, tap the active sensor button to bring up the probe
menu.
•
•
Tap Zero.
Rinse the probe with DI water and gently pat dry with a Kimwipe. In the box for Enter
known value, enter a value of 0 µS/cm and click KEEP > APPLY to store the first
calibration point (i.e., air = 0 µS/cm). If the software does not accept a value of 0,
instead enter a value of 1 µS/cm.
Immerse the end of the probe into the jar containing the 10000 μS calibration solution
making sure to completely cover the opening at the end of the probe. There should be
no bubbles trapped in this opening.
•
9
•
Tap the active sensor button again and chose Calibrate. Swirl the jar gently for 10
seconds. Enter 10000 µS/cm as the value for the known value. Tap KEEP > APPLY to
store the second calibration point. Click Done.
The conductivity probe is now calibrated and should remain that way unless the probe
disconnects.
Remove the probe from the calibration solution, rinse it with DI water and gently pat dry with a
Kimwipe.
PART A: CONDUCTIVITY OF VARIOUS SOLUTIONS
Equipment Used
•
•
•
•
•
•
•
•
•
•
•
•
•
•
25-mL buret (1); lower cabinet
100-mL beaker (5); student drawer
10-mL graduated cylinder (2); student drawer
50-mL graduated cylinder (1); student drawer
stir bar (1); student drawer
6-inch (medium) test tube (12); student drawer
magnetic stirrer/hotplate (1); table
ring stand (2); table
buret clamp (1); table drawer
utility clamp (1); table drawer
wood support block (1); table drawer
conductivity probe (1); Vernier equipment box (lower cabinet)
LabQuest 2 (1); Vernier equipment box (lower cabinet)
power adapter (1); Vernier equipment box (lower cabinet)
Reagents Used
•
•
•
•
•
0.04 M acetic acid (CH3COOH) solution
0.04 M sucrose (C12H22O11) solution
0.04 M hydrochloric acid (HCl) solution
0.04 M sodium hydroxide (NaOH) solution
0.04 M ammonia (NH3) solution
10
Preparation of Solutions
1. Obtain 7 clean and dry medium-sized test tubes.
2. Mark each test tube at the same level, approximately 1/3 full, with a Sharpie marker and
label the test tubes #1-7. Place the test tubes in a test tube rack.
3. Take the rack with the 7 test tubes to the hood containing the Part A reagents. Using the
pipets attached to the reagent bottles, fill each of the 7 test tubes 1/3 full with the solutions
(i.e., one solution per tube) as described in Table 1 below. Use the pipet designated for
each reagent bottle so that you do not contaminate the reagents. Note, in particular, that
the concentrations of the solutions in test tubes #3-7 are the same. Obtain deionized (DI)
and tap water from the appropriate sink.
Table 1 Solutions for Part A Conductivity Measurements.
test tube #
solution
description
1
deionized water
2
tap water
3
0.04 M acetic acid (CH3COOH)
weak acid
4
0.04 M sucrose (C12H22O11)
“table sugar”
5
0.04 M hydrochloric acid (HCl)
strong acid
6
0.04 M sodium hydroxide (NaOH)
strong base
7
0.04 M ammonia (NH3)
weak base
Taking Measurements
1. At your bench, place the conductivity probe into test tube #1. Check that there are no air
bubbles trapped in the tip of the conductivity probe. If there is an air bubble, you can
remove it by very gently moving the probe up and down in the test tube. Rapid shaking will
result in spilling.
2. Wait about 10 seconds and record the conductivity. (What conductivity do you expect for
DI water? If your measurement doesn’t fit your expectations, repeat the calibration step
and/or consult your Graduate Instructor.)
3. Remove the probe from the solution, rinse it with DI water, and gently blot it dry with a
Kimwipe.
4. Using the same procedure, measure and record the conductivities for the solutions in the
remaining test tubes (#2 – 7).
WASTE DISPOSAL AND CLEANUP
Discard all solutions from Part A down the drain. Rinse all glassware three times with tap water
and three times with DI water.
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PART B: THE EFFECT OF DILUTION ON CONDUCTIVITY
Equipment Used
•
•
•
•
5 medium test tubes
2 small beakers
10-mL graduated cylinder
LabQuest 2 and conductivity probe
Reagents Used
•
0.05 M sodium chloride (NaCl) solution
Preparation of Solutions
1. Using a small beaker, obtain approximately 50 mL of 0.05 M sodium chloride (NaCl)
solution.
2. Obtain 5 clean and dry medium-sized test tubes.
3. Label the test tubes #1 – 5.
4. Add the specified volumes of 0.05 M NaCl solution and DI water described in Table 2 below
to each test tube. Stir the solutions with a glass stir rod.
Table 2 Solutions for Part B Conductivity Measurements.
test tube # volume of 0.05 M NaCl (mL) volume of deionized water (mL)
1
2
8
2
4
6
3
6
4
4
8
2
5
10
0
Taking Measurements
1. Place the conductivity probe into test tube #1. Check that there are no air bubbles trapped
in the tip of the conductivity probe. If there is an air bubble, you can remove it by very
gently moving the probe up and down in the test tube. Rapid shaking will result in spilling.
2. Wait about 10 seconds and record the conductivity.
3. Remove the probe from the solution, rinse it with DI water, and gently blot it dry with a
Kimwipe.
4. Using the same procedure, measure and record the conductivities for the solutions in the
remaining test tubes (#2 – 5).
WASTE DISPOSAL AND CLEANUP
Discard all solutions from Part B down the drain. Rinse all glassware three times with tap water
and three times with DI water.
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PART C: EFFECTS OF ACID-BASE REACTIONS ON CONDUCTIVITY
Equipment Used
•
•
•
•
•
•
2 100-mL beakers
50-mL graduated cylinder
25-mL buret
magnetic stir bar
stirrer/hotplate and power cord
LabQuest 2 and conductivity probe
Reagents Used
•
•
0.04 M sodium hydroxide (NaOH) solution
0.04 M hydrochloric acid (HCl) solution
In Part C, you will perform an acid-base reaction between HCl and NaOH and examine how the
conductivity of the solution changes over the course of the reaction.
Review Appendix C: Volumetric Measurement Techniques at Brightspace > Labs > Reference
Materials for detailed information on how to use burets.
Preparation
1. Obtain about 20 mL of 0.04 M sodium hydroxide (NaOH) solution in a small beaker.
2. Measure 15 mL of the NaOH solution using a 50-mL graduated cylinder into a 100-mL
beaker.
3. Add 30 mL of DI water, measured using a 50-mL graduated cylinder, to the NaOH solution in
the beaker.
4. Obtain about 30 mL of 0.04 M hydrochloric acid (HCl) solution in another small beaker.
5. Rinse a 25-mL buret and fill it with the HCl solution.
6. Place the diluted NaOH solution on a stirrer/hotplate and secure it with a 3-prong clamp.
Gently put a magnetic stir bar in the solution.
Taking Measurements
1. Place the conductivity probe into the solution so that the magnetic stir bar does not hit it.
Secure the conductivity probe with a utility clamp (which you may need to pad with a folded
paper towel).
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2. Place the buret above the beaker containing NaOH and secure it with a buret clamp (see
Figure 5 below).
Figure 5 Experimental set up for reaction of HCl with NaOH.
3. Record the initial conductivity of the NaOH solution in the beaker. There is no need to rinse
or clean the probe before or after consecutive additions of HCl (i.e., until Part C data
collection is complete).
4. Using the buret, add about 5 mL of the 0.04 M HCl solution to the beaker of NaOH, while
stirring with the magnetic stirrer. Record the exact volume of HCl solution added.
5. Wait 10 seconds and then record the conductivity of the solution.
6. Continue adding HCl in 5 mL aliquots until a total of 25 mL of HCl solution has been added
to the beaker. Record the conductivity and exact volume added after each addition of HCl.
In total you will have 6 measured conductivities. You may need to reduce the stir rate as
the beaker will be close to full.
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WASTE DISPOSAL AND CLEANUP
•
•
•
•
Discard all solutions from Part C down the drain. Rinse all glassware three times with tap
water and three times with DI water.
Return the 25-mL buret and magnetic stir bar to your Graduate Instructor.
Shut down LabQuest. Disconnect the conductivity probe and power adapter and return all
equipment in the Vernier equipment box to your Graduate Instructor.
Return the magnetic stirrer/hotplate and power cord to the side bench.
DATA ANALYSIS, RESULTS AND DISCUSSION
Answer the questions and complete the calculations found on the Lab 6 Report
Form. The assignment is in the Labs module on Brightspace. Make sure you use
the version of the report form for on campus students.
When you have completed your report, upload it to Brightspace in the Lab 7
portal. It is due Mon. Mar. 15 by 11:59 PM ET.
15
Recitation week 9
Lab 8: Acid-Base Equilibria Part 1 (Week 10, 3/23-3/26)
TA sample problem (Your TA will work out this problem for you):
1.In a titration experiment, 12.5 mL of 0.500 M H2SO4 neutralized 50.0 mL of NaOH. What is the concentration of the NaOH
solution?
A: 0.25 M
Student Practice Problems (You will work these problems on your own and then review the solutions as a class):
In the lab, you titrated 20.0 mL of NaOH with 0.100 M HCl. Calculate the concentration of NaOH using the titration plot shown below and assuming that the volume of HCl at the equivalence point
was 30.0 mL.
A: 0.150 M
3. A vinegar solution of unknown concentration was prepared by diluting 10.00 mL of vinegar to a total volume of 50.00 mL with
deionized water. A 25.00-mL sample of the diluted vinegar solution required 20.24 mL of 0.1073 M NaOH to reach the
equivalence point in the titration. Calculate the concentration of acetic acid (in M) in the original vinegar solution (i.e., before
dilution).
A: 0.4344 M
4. For each of the following solutions, state whether it is acidic, basic, or neutral. Explain.
1.0.1 M KCH3COO basic
2.0.1 M Fe(NO3)3 acidic
3.0.1 M LiNO3 neutral
5. In which choice are the 0.1 M aqueous solutions of salts below ranked according to decreasing pH?
NH4Br, NaBrO2, NaBr, NaClO2
1.NaBr > NaBrO2 > NaClO2 > NH4Br
2.NH4Br > NaBrO2 > NaBr > NaClO2
3.NaBrO2 > NaClO2 > NaBr > NH4Br
4.NaBrO2 > NH4Br > NaClO2 > NaBr
6. Pyridine is a weak base with a Kb of 1.7 x 10−9 at 25 °C. Calculate the initial molar concentration of an aqueous solution of
pyridine if the pH is 9.1.
A: 9.3 x 10−2 M
7. Which ionic compound forms a pH-neutral aqueous solution at 25 ºC?
1.Na2CO3
2.KNO3
3.LiF
4.NH4Cl
5.K2S
8. Calculate the pH of a 5.2 x 10−2 M solution of sodium benzoate (C6H5CO2Na) in otherwise pure water at 25oC. For benzoic acid
(C6H5CO2H), Ka = 6.5 x 10−5.
A: 8.45
9. Which of the following solutions is the most basic? Which is the most acidic?
1.0.10 M NaF – most basic
2.0.10 M NH4Cl – most acidic
3.0.10 M NaNO3
4.0.10 M NaClO4
5.0.10 M KBr
10. In chemical reactions, BF3 can act as
1.a Lewis acid AND a Brønsted-Lowry acid.
2.a Lewis base AND a Brønsted-Lowry base.
3.a Brønsted-Lowry acid BUT NOT a Lewis acid.
4.a Brønsted-Lowry base BUT NOT a Lewis base.
5.a Lewis acid BUT NOT a Brønsted-Lowry acid.
11. Sodium fluoride (NaF) is commonly added to the drinking water to prevent tooth decay in children. Calculate the pH of a
solution prepared by dissolving 0.4 mol of NaF in enough water to make 1.00 L of solution. (Ka of HF is 6.8 x 10−4)
A: 8.38
A student in lab adds 5.00 mL of 0.04 M HCl to 45 mL of 0.013 m NaOH solution.
Which one of the following reaction tables correctly shows the calculation of moles
of each of the ions in solution?
a)
Na+
NaOH(aq) + HCl(aq) + NaCl(aq) +
H20(1)
Amount
OH H+ ci
(mol)
Initial 0.0006 0.0006
Addition
0.0002 0.0002
Reaction -0.0002 -0.0002
Final 0.0006 0.0004 0 0.0002
b)
NaOH(aq) + HCl(aq) + NaCl(aq) +
H2O(1)
Amount
(mol)
Na+ OH H+
ci
Initial 0.0006 0.0006 0.0002 0.0002
Addition
0.0002 0.0002
Reaction
-0.0002 -0.0002
Final
0.0006 0.0004 0.0002 0.0004
H
NaOH(aq) + HCl(aq) + NaCl(aq) +
H2O(1)
Amount
(mol)
Nat OH
01
Initial 0.0006 0.0006
Addition
0.0002 0.0002
Reaction -0.0002 -0.0002 -0.0002
Final 0.0006 0.0004 0
0
Greetings,
Thanks for your email and interest in Purdue University.
Hope you and your family are doing good.
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Please have Elgin Community College send your official transcript directly to Purdue University West Lafayette, IN either electronically or can be
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Hope this helps.
Best,
Surya Dwadasi
Question 1 (2 points)
Calculate the total concentration of ions (Na+ and CI-) in a solution prepared by
diluting 4 mL of 0.05 M NaCl to a total volume of 10 mL with deionized (DI) water.
Enter your answer using 2 decimal places.
Your Answer:
Answer
units
Question 2 (2 points)
What is the molecular equation for the acid-base reaction that will be done in this
lab?
a) NH3(aq) + HCl(aq) – NH4+(aq) + CI+(1)
b) NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(1)
c) H2O(aq) + HCl(aq) → H30+(aq) + CI+(1)
d) 2NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2H2O(1)
Question 3 (2 points)
Match the following definitions
1. Compounds that do not dissociate
into ions in solution at all. Solutions
do not conduct a current.
nonelectrolyte
< weak electrolyte 2. Compounds that dissociate into ions in solution very little. Solutions conduct current poorly. strong electrolyte 3. Compounds that dissociate completely into ions in solution. Solutions conduct current well. Question 4 (2 points) One would expect NH3 to have a lower conductivity than KOH. True False Na+ H NaOH(aq) + HCl(aq) + NaCl(aq) + H2O(1) Amount OH (mol) Ci Initial 0.0006 0.0006 Addition 0.0002 0.0002 Reaction - -0.0002 -0.0002 -0.0002 Final 0.0006 0.0004 0 0 - d) Na+ NaOH(aq) + HCl(aq) + NaCl(aq) + H20(1) Amount OH н" ci (mol) Initial 0.0006 0.0006 Addition 0.0002 0.0002 Reaction -0.0002 -0.0002 Final 0.0004 0.0006 0.0002 0 Calculate [OH-] for a solution where [H,0+] = 0.00363 M. [OH-] = M Arrange the metal ions by their acidity in aqueous solution. Most acidic Least acidic Answer Bank In3+ Cd2+ Ga3+ A 0.0277 M solution of a monoprotic weak acid has an equilibrium [H,0+] of 8.39 x 10-² M. Determine the K, of the weak acid. 4.34 x 10-1 2.74 x 102 O 3.65 x 10-3 2.54 x 10-3 In a 1.0 x 10-4 M solution of HCN(aq), arrange the species by their relative molar amounts in solution. Greatest amount Least amount Answer Bank OH HO H0+ CN- HCN Look up the ionization constant (K2) values for the given acids. acetic (ethanoic) acid: K = formic (methanoic) acid: Ky = propanoic acid: K = Arrange the acids according to strength. Arrange the acids according to strength. Strongest acid Weakest acid Answer Bank formic acid acetic acid propanoic acid