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NAME ____________________________________________________ PERIOD ____________
Grade Level Indicators
 Show how atoms may be bonded together by losing, gaining or sharing electrons and that in a chemical reaction,
the number, type of atoms and total mass must be the same before and after the reaction (e.g., writing correct
chemical formulas and writing balanced chemical equations).
 Investigate the properties of pure substances and mixtures (e.g., density, conductivity, hardness, properties of
alloys, superconductors and semiconductors).
Lab 20 “EXPERIMENT in CONDUCTIVITY”
Pre-lab Questions:
1. Write the formulas for the following acids:
(a) phosphoric _________;
(b) perchloric _________;
(c) nitric _________;
(d) sulfuric _________;
(e) hydrochloric _________;
(f) hydrofluoric _________;
(g) carbonic _________;
(h) chromic _________.
2. Write the formulas for the following bases:
(a) calcium hydroxide _________;
(b) potassium hydroxide _________;
(c) sodium hydroxide _________;
(d) cupric hydroxide _________;
(e) aluminum hydroxide _________;
(f) ammonium hydroxide_________.
(g) barium hydroxide_________.
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3. Write the formulas for the following salts:
(a) potassium chromate ______________;
(b) potassium sulfate ______________;
(c) copper (II) chloride ______________;
(d) calcium carbonate ______________;
(e) potassium iodide ______________;
(f) lead (II) nitrate ______________;
(g) sodium chloride ______________;
(h) ammonium acetate ______________;
(i) cobalt (II) chloride ______________;
(i) iron (III) nitrate ______________;
4. Name the following compounds
(a) HNO2
____________________________________;
(b) Na2CO3 ____________________________________;
(c) HClO
____________________________________;
(d) HC2H3O2 ____________________________________;
(e) CaS
____________________________________;
(f) HBr
____________________________________;
(g) KI
(h) NH3
____________________________________;
____________________________________;
(i) NH4SCN
____________________________________.
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Water Molecules
Salt Crystal of Sodium Chloride
5. Pictured above are representations of a water molecule and a salt crystal of sodium chloride. Suppose a
student measures out 4.249 grams of sodium chloride and then adds the salt to a 100 mL (+/-0.16 mL)
volumetric flask containing distilled water. The solution is then transferred into a 250 mL beaker.
a. What is the concentration of the sodium chloride solution in moles per liter (M)? Show all work!
b. In the space provided below, show the interaction of the components of sodium chloride by
making a drawing that represents the different particles present in the solution in the 250 mL
beaker. Base the particles in your drawing on the particles shown in the reaction above. Include
one formula unit of sodium chloride and no more than ten molecules of water. Your drawing
must include the following details.
i. identity of all atoms and ions (including symbol and charge where appropriate).
ii. the arrangement and proper orientation of the particles in the solution.
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EXPERIMENT in CONDUCTIVITY
In this experiment you will use a piece of equipment called a well plate shown in Figure I. This piece of
equipment is made of clear plastic and contains wells used to hold solutions. Each well can be identified by a
number. In this experiment you will use a 12 or 16-well plate. The wells are small and only a few drops of
reagent will be needed. Doing experiments on a 'microscale' is very economical and considerably safer than
large scale experiments.
The conductivity apparatus used in this experiment using a 9-volt battery. A light-emitting-diode (LED) has
been wired to the battery. Two wires (electrodes) are attached to the LED so that if the electrodes are placed in a
solution which conducts electricity the LED will glow at a particular intensity.
PART I: Introduction to Conductivity
Check the conductivity apparatus using a test copper wire. Touch the electrodes of the conductivity apparatus to
the copper wire. Check that the LED (light emitting diode) glows brightly. When viewing the LED, look down
from above the LED rather than from the side. It may also help to slightly darken the room. If the conductivity
apparatus works continue with the lab, if it doesn’t see your instructor. It is important to wash and dry the
electrodes with distilled water following each measurement between each procedural step.
Using the spot well plate and a dropper, “fill” a well with tap water and just introduce the electrodes to the tap
water. Note whether the LED glows brilliantly, faintly, or not at all. NOTE: For each of the tests that follow,
you should immerse the electrodes to approximately the same depth.
Obs. #1:
Try immersing the electrodes more and more deeply into the tap water. Record your results.
Obs. #2:
Dry the electrodes. In a well add distilled water and test using the conductivity apparatus. Does the LED glow
brilliantly, moderately, faintly, or not at all?
Obs. #3:
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Test dry sucrose (C12H22O11) with the conductivity apparatus. Use a microspatula to quarter-fill a well with dry
sucrose. Be careful not to spill sucrose in the surrounding wells and be sure the electrodes are clean and dry
before testing the sample.
Obs. #4:
With the electrodes still in contact with the solid sucrose in the well, add a few drops of distilled water and
observe LED. Record your observations.
Obs. #5:
In another well, test dry sodium chloride (NaCl) with the conductivity apparatus. Use a dry microspatula to
place the NaCl into the well. Be sure the electrodes are dry before testing the sample.
Obs. #6:
With the electrodes still in contact with the dry sodium chloride, add a few drops of distilled water. Record your
observations.
Obs. #7:
Explain your observation of the conductivities of distilled water and tap water. (Obs. #’s 1-3)
Expl. #1:
Explain your observations of the conductivities of solid sucrose and of sucrose solution. (Obs. #’s 4-5)
Expl. #2:
Explain your observations of the conductivities of solid sodium chloride and of sodium chloride solution. Hint:
Write a solubility equation for the dissociation of sodium chloride. (Obs. #’s 6-7)
Expl. #3:
Explain the difference in the conductivities of the sodium chloride solution and sucrose solution.
(Obs. #’s 5, 7)
Expl. #4:
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Place a small sample of solid lead (II) nitrate on one well of a spot plate. In a second well, place a sample of
solid potassium iodide. Briefly describe the initial appearance of the dry potassium iodide and lead (II) nitrate
solids.
Obs. #8:
Test the small sample of solid lead (II) nitrate with the conductivity apparatus.
Obs. #9:
Test the small sample of solid potassium iodide with the conductivity apparatus.
Obs. #10:
Clean a Petri dish. Cover the bottom of the dish with a thin layer of distilled water. Use a spatula to carefully
place a few crystals of lead (II) nitrate into the water close to one side of the Petri dish. Try not to agitate the
water when adding the solid. Use a clean spatula to carefully add a few crystals of potassium iodide to the
opposite side of the Petri dish. Do not bump the Petri dish. It is important that the water not be agitated during
the experiment. Carefully touch the conductivity apparatus near each solid.
Obs. #11:
Explain why the two ionic solids were nonconductors, but the solution of each salt was. Hint: Write a solubility
equation for the dissociation of lead (II) nitrate and potassium iodide.
Expl. #5:
Continue to watch what happens. Carefully test the Petri dish in different areas.
Obs. #12:
Explain your observations for the pictures above. In your explanation include the use of the following terms:
anion, cation, electrolyte(non, weak or strong), precipitate, soluble, solvation, and electrical conductivity. A
compete and net ionic equation
Expl. #6 of (Obs. #9-12):
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PART II: Strong and Weak Electrolytes
Place four drops of the following three acids in separate wells: 0.1 M hydrochloric acid (HCl), 0.1 M acetic acid
(HC2H3O2), and 0.1 M sulfuric acid (H2SO4). Compare the conductivities of the three solutions by observing the
intensity of the LED, for example, no glow (non conductor), faint (poor conductor) or brilliant (good
conductor). Be sure to clean and dry the electrodes after each use. Classify each acid as a nonelectrolyte,
weak electrolyte or strong electrolyte. Identify the ions present in each solution that account for the
conductivity. If you can’t identify the ions, wait until you have written reactions in the next part.
Obs. #13:
Consider the following sample set of observations. Jennifer Redleg, an aspiring chemistry student, tested the
conductivity of a 0.1 M nitric acid solution (HNO3) and found that the LED glowed brightly. Jennifer concluded
that the HNO3 is a strong electrolyte. To demonstrate her knowledge of the ions formed in the solution Jennifer
wrote the following ionization equation:
HNO3(aq) + H2O(l)
H3O+(aq) + NO3-(aq) .
Write a similar chemical equation that indicates the species that are in solution for each of the compounds
whose conductivities were measured in Observations 13. (Jennifer also knew that a strong electrolyte can be
shown by an arrow pointing one way and a weak electrolyte with a reversible arrows.)
Equations #1A:
Place four drops of the following two bases in separate wells: 0.1 M sodium hydroxide (NaOH) and 0.1 M
ammonium hydroxide (NH4OH) and observe the intensity of the LED. Be sure to clean and dry the electrodes
after each use. Compare the conductivities of the two solutions by classifying each base solution as a
nonelectrolyte, weak electrolyte or strong electrolyte. Identify the ions that are present in each solution that
would account for the conductivity.
Obs. #14:
Equations #1B:
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Place four drops of the following four salt solutions in four separate wells: 0.1 M sodium acetate (NaC2H3O2),
0.1 M sodium chloride (NaCl), four drops of 0.1 M ammonium acetate (NH4C2H3O2) and 0.1 M ammonium
chloride (NH4Cl). Use the conductivity apparatus to check each solution. Be sure to clean and dry the
electrodes after each test. Record your results. Classify each solution as a nonelectrolyte, weak electrolyte or
strong electrolyte. Identify the ions that are present in each solution that would account for the conductivity.
Obs. #15:
Place four drops of methanol (CH3OH) in one well and four drops of ethanol (C2H5OH) in another well.
Compare the conductivities of the two solutions by observing the intensities of the LED, for example, no glow
(non conductor), faint (poor conductor) or brilliant (good conductors). Be sure to clean and dry the electrodes
after each use. Classify each compound as a nonelectrolyte, weak electrolyte or strong electrolyte. Identify the
ions that are present in each solution that would account for the conductivity.
Obs. #16:
Explain what is wrong with the statement: “All acids and bases are strong electrolytes.”
Expl. #7
Explain what is wrong with the statement: “All substances that have an “O-H group” (oxygen and hydrogen
bonded to each other) are bases that are strong electrolytes.”
Expl. #8
Explain why all the solutions in Observation #15 were strong electrolytes
Expl. #9
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PART III: Chemical Reactions
Make and observe the color of 100.0 mL of a 0.1 M solutions of the salts listed below and record the formula
and color for both the cation and the anion. (If a solution is colorless, the ions it contains must also be colorless.
The color of an ion is independent of the color of any other ions in the solution.)
Procedure to make _________________________ 0.1 M solution assigned to your team:
Place two drops of 0.1 M AgNO3 with four drops of 0.1 M NiCl2 in a well and observe what happens, then test
the conductivity of the solution. Record your visual observations and the conductivity results.
Obs. #17
Identify the precipitate (if any) and the ions present in solution. Explain how you arrived at your conclusions.
Expl. #10
Write a balanced chemical equation for the reaction.
Reaction 2
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Ions that remain unreacted in a solution are called spectator ions. A net ionic equation can be obtained by
algebraically canceling all the spectator ions. Write the net ionic equation for the above reaction.
Reaction 2A
Place two drops of 0.1 M AgNO3 and four drops of 0.1 M CoCl2 in a well. Observe what happens. Test the
conductivity of the solution. (Be sure to clean the electrodes of the conductivity device.)
Obs. #18
Write a balanced chemical equation and net ionic equations for the reaction in obs. #18.
Reaction 3A and 3B
Add a crystal of iron(III) nitrate, Fe(NO3)3 to one side of a Petri dish containing distilled water and then add a
few crystals of ammonium thiocyanate, NH4SCN to the other side. Describe what happens. Include a picture.
Obs. #19:
Explain what must be happening in the solution to account for your observations.
Expl. #11:
Write a net ionic equations for the reaction.
Reaction 4
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Post-lab Questions:
The answers to the following problems must accompany your laboratory report.
1. Write a balanced chemical equation and then a net ionic equation for each of the following
combinations.
(a) HNO3(aq) and NaOH(aq)
(b) H2SO4 (aq) and NaOH (aq)
(c) AgNO3(aq) and KCl (aq)
2. Why is it necessary to use distilled water when making and then testing the conductivity of aqueous
solutions you made in Part 1 of this lab?
3. Aqueous ammonia, NH3 (aq), and acetic acid, HC2H3O2(aq) solutions of equal concentrations, conduct
electric current the same, both do so weakly. Explain in words why the addition of one solution to the
other results in a substantial increase in electrical conductivity. (Hint: Write a balanced chemical
equation for the reaction that takes place.)
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4. Making predictions based on previous experimental evidence is an important goal for a chemist. Use
your classification system, to predict the conductivity of each of the following solutions. Predict whether
the substance is a strong electrolyte (SE), a weak electrolyte (WE), or a nonelectrolyte (NE).
(Hint: Fill in the first three columns first)
Substance
(a)
(b)
(c)
(d)
(e)
(f)
(g)
(h)
(i)
(j)
(k)
Ionic, Molecular
Soluble
Ionize or
Dissociate
Conductivity
HClO4
Ca(NO3)2
NH2CONH2 (urea)
HBr
H3PO4
(NH4)2CO3
PbCl2
KOH
C3H5(OH)3 (glycerol)
PbI2
CH3CH2CH2OH
5. The conductivity of a well mixed 0.1 M H2SO4 solution is measured with the slow addition of a 0.1 M
Ba(OH)2, solution. A graph shown above shows the conductivity measure on the y-axis and the amount
of Ba(OH)2 solution added is on the x-axis. Describe and explain the conductivity being shown by the
graph.
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