CHM 161: Ionization and Properties of Ionic Compounds
Introduction
In this experiment you will investigate the electrical conductance of aqueous solutions of
several ionic and covalent compounds and the properties of ionic compounds. Substances (or
solutions) that are electrical conductors contain charged particles that are free to move
throughout the substance or solution. For example, metals conduct an electric current because
they contain mobile electrons that carry a unit negative charge.
Pure water is a poor conductor of electricity. However, some substances, when dissolved in
water, dissociate into ions that are free to move throughout the solution. The ease with which
an electric current passes through an aqueous solution, as measured by the conductance of the
solution, depends on the concentration of mobile charges: high concentration of ions, high
conductance; low concentration of ions, low conductance; nearly zero concentration of ions,
nearly zero conductance.
You will perform experiments on several aqueous solutions and you will seek to learn
something about the nature of the dissolved species in each case: Are ions present? Is the
concentration of ions high or low? Do ions have characteristic properties? Is the dissolved
substance in a molecular rather than in an ionic form? You also will follow the conductance of a
solution during a chemical reaction.
The apparatus that you will use in several experiments is an ordinary light bulb in an open
circuit as shown in Figure 1. You will close the circuit by inserting the solution to be tested
between the electrodes attached to the bulb. If the test solution has mobile ions, the light bulb
will glow. The brighter the light, the higher the concentration of dissolved ions. Be careful in
handling the light-bulb apparatus. It should be plugged in only when making a determination
of the conductivity of the solution. The following procedures must be followed:
• Before adding a test solution to a beaker, it should be clean and rinsed with distilled water.
• Prior to the performance of an electrical conductance test, the electrodes of the conductivity
apparatus must be clean and dry. The electrodes are copper rods covered with insulating
material except for about one-half inch at the bottom. Since conductance is dependent on the
area of the electrodes in contact with the solution, the exposed copper should be completely
covered with the solution when performing an electrical conductance experiment. Most tests
will be done in a 50 mL beaker that should be filled with the test solution to a depth of
approximately two centimeters.
• After plugging in the electrical conductance apparatus, raise the beaker with the test
solution up to the electrodes and note in your lab notebook the brightness of the light bulb.
Withdraw the beaker and unplug the apparatus.
• The electrodes must be carefully cleaned after each test so that the next solution to be
investigated will not be contaminated by the previous one. Unplug the apparatus after each
experiment and use a wash bottle filled with distilled water to rinse the electrodes into a 400
mL beaker. The electrodes should then be wiped dry with a paper towel or tissue.
• The apparatus does not have an "on-off" switch. To turn it off you must remove the plug
from the electrical socket. Please be careful and do not pull on the cord!
Figure 1. Electrical Conductivity Apparatus
Procedure
The stations may be done in any order, but you should try to complete each station and
answer all the questions in your notebook before moving to another station. Ask for help
from the TA or instructor if you have questions about what to write in your notebook.
In writing your prelab outline, write a separate list of steps for each section. Please include
the questions interspersed throughout the experimental procedure in your outline. You
should answer them in your notebook as you go along during the experiment.
Work in pairs. Both students should enter all information in their own notebook. The goal of
today's experiment is for you to be able to write balanced ionic equations for substances that
dissolve in water and for chemicals that react in aqueous solution. This is a qualitative
experiment so you should be able to write explanations about what you see rather than do
numerical calculations. Enter all of your observations in your notebook as you go along.
Six stations for experiments have been set up around the lab with appropriate apparatus and
chemicals at each station. Two additional stations will be demonstrated by your instructor or
teaching assistant. You and your partner are eventually to visit each of the six stations and
perform the indicated operations. If a "traffic jam" develops at a particular station, bypass that
station and return to it later. While waiting, you can work on the postlab questions for stations
you have completed.
Before leaving each station, be sure that both you and your partner have recorded the complete
set of experimental observations for that station in each of your notebooks.
Station 1
The purpose of this station is to compare the electrical conductance of distilled and tap water.
You should complete the experiments at this station in a hood with the hood-light turned off.
Put some distilled water in a clean 50 mL beaker. Test the conductance of the water with the
light-bulb apparatus. Withdraw the beaker and unplug the apparatus. Add two drops of 0.1 M
AgNO3 (aq) to some distilled water in a test tube and record your observations. Finally, add two
drops of 1.0 M HCl (aq) to the distilled water containing the AgNO3 (aq) and record your
observations. Write a balanced net ionic equation in your notebook for any reaction that you
observe.
After discarding the waste solution in the container in the hood, clean the beaker and put some
tap water in it. Test the conductance of the tap water with the light-bulb apparatus. As before,
determine the effect of adding two drops of AgNO3 (aq) to some tap water in a test tube.
Comment on any difference in behavior that you saw between the distilled water and the tap
water.
Stations 2A and 2B
You will test the conductivity of the following twelve 1.0 M solutions:
HCl (aq)
HNO3 aq)
NaCl (aq)
H2SO4 (aq)
NH3 (aq)
MgCl2 (aq)
NaOH (aq)
CH2OHCH2OH (aq), ethylene glycol
CH3OH (aq), methyl alcohol
C6H12O6 (aq), sucrose (sugar)
NH4Cl (aq)
CH3COOH (aq), acetic acid
The solutions will have been already added to 50 mL beakers. All students will use the same
beakers. Be careful not to contaminate them. Rinse the electrodes thoroughly with distilled
water after each test.
Categorize each dissolved substance as a strong electrolyte (bright glow), a weak electrolyte
(feeble glow), or a non-electrolyte (no glow at all). Where appropriate, write an equation
showing the ions formed in the dissolving process; use a single arrow ( ⇒) for a strong
electrolyte and a double arrow ( ⇔ ) for a weak electrolyte.
Station 3 (Done in Hood)
Magnesium metal reacts with acidic aqueous solutions to liberate hydrogen:
Mg(s) + 2 H+ (aq) → Mg2+ (aq) + H2 (g).
The vigor of the reaction (the speed with which the Mg dissolves) depends on the
availability (i.e., the concentration) of hydrogen ions.
Put 5 mL of 6 M HCl (aq) in one graduated cylinder and 5 mL of 6 M CH3COOH (aq) in
another. At the same time, drop a piece of Mg ribbon, about 1 cm long, into each of the 50
mL beakers. Compare the vigor of the reaction in the two cases. Write a balanced
“molecular” equation (with appropriate phase subscripts) for the reaction in each beaker.
The resulting solutions are harmless to the environment and may be discarded in the sink.
Station 4 (Demonstration by TA)
When a weak electrolyte is diluted with water, the fraction of the molecules ionized
increases but the number of molecules per unit volume decreases. The net result upon
dilution can be an increase or decrease in the number of ions per unit volume, depending on
which effect "runs ahead" of the other.
Use a truly dry 250 mL beaker. Add to the beaker sufficient glacial acetic acid (99.6 w/w %)
to cover the bottom of the electrodes.
SAFETY NOTE: Avoid getting the acetic acid on your skin - it is a strong irritant.
While placing the light-bulb apparatus into the beaker (so that the electrodes are immersed
in the acid), steadily agitate the solution between the electrodes with a stirring bar. Dilute
the acetic acid by slowly adding distilled water. Observe what the bulb does as the water
dilutes the acetic acid. Discuss your observations in terms of the effect of dilution on the
percent ionization of the dissolved acetic acid. (See graph posted in the lab.)
Station 5 (Demonstration by TA)
At this station you will observe the change in conductivity during the course of a chemical
reaction between barium hydroxide and sulfuric acid.
Fill a 50 mL buret with 0.1 M H2SO4 (aq). Next, dilute about 10 mL of saturated Ba(OH)2 (aq)
solution in a 150 mL beaker with about 15 mL of distilled water. Put a stirring bar in the beaker
and place the beaker on a stirring plate. Place the light bulb apparatus next to the stirrer and put
the electrodes into the beaker. Record the brightness of the bulb when the apparatus is plugged
in. While the solution is stirred, adjust the stopcock on the dispenser to introduce steady
dripping of H2SO4 (aq) into the beaker (use a drip-rate a little faster than a drop a second). What
happens during the course of the reaction? Explain the behavior of the light bulb in terms of
what you think happens in the beaker as the H2SO4 (aq) solution is added. Use both words and
equations in constructing your explanation. Discard the test solution in the labeled waste jar.
Write a balanced “molecular” equation (with appropriate phase subscripts) for the reaction
which takes place.
Station 6 (Do this station after you have finished all the others, or if no other station you
need is available.)
Strong electrolytes contain both cations and anions in solution. Many ions are colored; others
are colorless. Although the anion can alter the color of the cation, and vice versa, most ions
exhibit characteristic colors (or are always colorless) in aqueous solution. In this experiment you
will observe the characteristic colors of twelve different ions: 6 cations and 6 anions, some of
which are colorless. Write the equation for the dissociation into ions of each of the compounds
(in 1.0 M aqueous solutions) below. Assuming that there is not any "mixing of colors" (for
example, yellow and blue to make green), deduce a color for each of the twelve different ions in
aqueous solution. Record in your notebook the color (not “clear” - all solutions are clear) of each
ion in each solution.
NiCl2 (aq)
Na2SO4 (aq)
KMnO4(aq)
K2SO4 (aq)
CuSO4 (aq)
KNO3 (aq)
K2Cr2O7 (aq)
ZnCl2 (aq)
KCl(aq)
CoCl2 (aq)
NaCl (aq)
K2CrO4 (aq)
Use your observations to predict the colors of 1.0 M aqueous solutions of the following and
explain your reasoning: NaNO3 (aq) Zn(MnO4)2 (aq) Na2Cr2O7 (aq) CoSO4(aq)
Station 7 (Done in Hood)
When strong electrolytes dissociate into cations and anions in aqueous solution, each can
undergo a chemical change independent of the other. Thus, sodium ions behave the same in
solution whether they result from the addition to water of sodium chloride, sodium bromide, or
sodium iodide. In this experiment you will investigate characteristic reactions of the copper(II)
cation.
In each of three 50 mL beakers, add approximately 2-3 mL of 0.5 M solutions of CuCl2 (aq),
CuSO4 (aq), and Cu(NO3)2 (aq), one solution per beaker. Why do you think the colors of the
solutions are so similar? (Refer to your analysis from Station 6). To each beaker add several
drops ofn a solution of the strong electrolyte, potassium hexacyanoferrate(II), K4Fe(CN)6
(aq), and swirl.
K4Fe(CN)6 (aq) → 4 K+ (aq) + Fe(CN)64– (aq)
Are the results the same for each solution? Why or why not? What is the formula for the
brownish precipitate?
Rinse the beakers and put in fresh 2-3 mL samples of the same copper-containing solutions.
Add 1-2 mL of Na2CO3 (aq) solution to each beaker and swirl. Are the results the same for
each solution? What is the formula and color of the precipitate formed?
Station 8
At this station you will perform a series of qualitative tests to observe the behavior of the
Ca2+ (aq) ion. You will investigate the behavior of the Ca2+ (aq) ion in the presence of the
SO42–(aq) ion, the PO43– (aq) ion, the CrO42– (aq) ion, the NO3– (aq) ion, and the OH– (aq) ion.
The source of OH–(aq) will be the reversible hydrolysis of aqueous ammonia.
NH3 (aq) + H2O (l) ⇔ NH4+ (aq) + OH– (aq)
In your notebook, reproduce the following (blank) data table:
Table X. Test Solution: CaCl2(aq)
Source Solution
Color of mixture
(“clear” is not a color)
Solubility
Formula of ppt.
(if any)
Na2SO4(aq)
XXXXXXXXX
XXXXXXXXX
XXXXXXXXX
Na3PO4(aq)
XXXXXXXXX
XXXXXXXXX
XXXXXXXXX
Na2CrO4(aq)
XXXXXXXXX
XXXXXXXXX
XXXXXXXXX
NaNO3(aq)
XXXXXXXXX
XXXXXXXXX
XXXXXXXXX
NH3(aq)
XXXXXXXXX
XXXXXXXXX
XXXXXXXXX
Take a clear-glass spot plate to a hood containing the following solutions in dropper bottles:
0.5 M CaCl2 (aq) (source of Ca2+), 0.5 M Na2SO4 (aq) (source of SO42–), 0.5 M Na3PO4 (aq)
(source of PO43–), 0.5 M Na2CrO4 (aq) (source of CrO42–), 0.5 M NaNO3 (aq) (source of NO3–),
and 1.0 M NH3 (aq) (source of OH–).
Put 3 drops of CaCl2 (aq) in each of any five wells. Next, add 3 drops of the source-solution
to the appropriate well.
For example, in the first well add 3 drops of CaCl2 and 3 drops of Na3PO4 (aq) and mix the
material in the well of the spot plate with a clean stirring rod. Rinse the stir rod with
distilled water before you use it in the next well.
The mixture in a well of your spot plate will remain clear if the product is soluble; will turn
slightly cloudy, with perhaps a small amount of solid settling out, if sparingly soluble; or
will become quite opaque and will show strong evidence of precipitate formation if
insoluble.
In your notebook describe the result of adding 3 drops of each source-solution to 3 drops of
CaCl2(aq). Give the color of the mixture in each well. Indicate whether a soluble, sparingly
soluble, or insoluble compound was formed. Give what you think to be the formula for any
sparingly soluble or insoluble substance that was formed. The test solution contains both the
Ca2+ (aq) ion and the Cl–(aq) ion. The latter does not form a precipitate with any of the
cations from the source solutions. Thus, if a reaction occurs with the formation of a
precipitate, it must be between the Ca2+ (aq) ion and an anion from the source solution.
CHM 161 Postlab Assignment for Ionization Experiment
You will need to refer to your notes from the previous two prelabs as well as to your lab
notebook.
Station 1. Write the net ionic equation for the reaction which takes place when HCl(aq) is
added to the solution of AgNO3(aq) in distilled water. What happens to the conductivity
when the HCl is added and why? What do you think the main source of ions in tap water
is?
Station 2. Classify each substance as a strong, weak, or non-electrolyte based on your
observations. Write a balanced chemical equation to show the dissociation into ions (or reaction
with water in the cases of weak acids and bases) that illustrates what happens when the
substance dissolves. Include phase subscripts on all species. For non-electrolytes, write:
A(aq) → N.R.
Station 3. Write an explanation of the observations you made of the reactions in the two
beakers given that the rate of reaction is proportional to the concentration of H3O+ in the
solution.
Station 4. How does the brightness of the bulb change as the acetic acid is diluted with
distilled water? Why does the conductivity increase when water is first added, and then
gradually decrease after a lot of water has been added?
Station 5. (At the beginning, some solid Ba(OH)2 may be present indicating that you really
have a saturated solution of it.) What ions are present in solution before any sulfuric acid is
added? What is present in the beaker when the light goes out? What ions are present in the
beaker when the light goes back on?
Station 6. Write the formulas of the ions into which each solute is dissociated and state the
color of each ion. In addition to the solutions listed in the instructions, predict and explain the
colors of aqueous solutions of: Ni(NO3)2, Ca(MnO4)2, Na2CrO4, and Co(NO3)2
Station 7. Explain why the solutions of the three copper salts are all the same color. Write the
formulas of the solid compounds formed when the K4Fe(CN)6 and the Na2CO3 solutions are
added to the solutions.
Station 8. Make a table of your observations for the five solutions tested with CaCl2(aq).
Write a balanced net ionic equation for the reaction of Ca2+ with each of the five negative ions
you tested. If no observable reaction occurred, write N.R. instead of products.
Extra Credit: What correlation can you see between the charge on the negative ion in the
reagents in station 8 and the amount (if any) of precipitate formed? Explain in terms of the
attractive forces among the ions in the crystal.