EXPERIMENT 8:
IONIZATION AND ELECTROLYTE BEHAVIOR
Introduction: You will observe the conductivity of several pure substances and of aqueous
solutions formed from a number of substances. You will decide, based on the observed
conductivity, if a substance is a strong electrolyte, a weak electrolyte or a nonelectrolyte. You
will use observed changes in conductivity as a way to monitor the progress of a chemical
reaction and to predict the products of the reaction.
Background: Electricity can be viewed as a flow of charged particles, like electrons or ions.
Very few pure substances, even ionic compounds, conduct electricity. The exceptions, the metallic
elements and a few rare other elements and compounds, are important industrial materials.
However, several types of substances conduct electricity when they are dissolved
in water to make an aqueous solution. These substances are called electrolytes and include acids,
bases and salts (ionic compounds).
Salts and metal hydroxide bases already have ions present. Those ions separate, or dissociate, from
one another as the solid salt dissolves.
Ba(OH)2 (s)
Ba(OH)2 (aq)
Ba2+ (aq) + 2 OH¯ (aq)
Molecular acids and bases consist of neutral molecules when pure. However, during dissolving,
certain covalent bonds break to create ions. Thus electrolytes release free ions
into the solvent. Those ions are free to move about in the solution and to carry current through the
solution. Electrolytes are classed by how many ions they release into the
solution. Some substances are 100% dissociated in solution; they are strong electrolytes. Soluble
ionic compounds and a few acids like hydrochloric, nitric and sulfuric acids are
strong electrolytes.
HCl (aq)
H+ (aq) + Cl¯ (aq)
Other substances release only a few ions into solution with the bulk of the substance remaining
associated; these are weak electrolytes. All slightly soluble salts and many molecular acids and
bases (except the strong ones) are weak electrolytes.
HC2H3O2 (aq)
H+ (aq) + C2H3O2¯ (aq)
These “weakly ionized” substances undergo a process of ionization and recombination until
equilibrium is established, favoring the reaction to the left. This is the subject of later chapters in the
text. Suffice it to say that only a few ions are produced. Your instructor may demonstrate
conductivity, using an apparatus consisting of electrodes that are placed into the test sample and a
light bulb that lights only if electrical current can pass through the substance being tested. The light
bulb may shine brightly if there are many ions free to carry a lot of current or it may light only dimly
if there are few ions available to conduct. In fact, it may not light up at all if there are almost no ions
free to conduct electricity. Based on the behavior of the light bulb, you will classify the substance
being tested as a strong electrolyte, a weak electrolyte or a nonelectrolyte. Alternatively – students
may use individual conductivity meters that are connected to batteries.
After observing a number of pure substances and solutions with a single compound mixed into
water, you will observe a chemical reaction. Two dissolved compounds, H2SO4 and Ba(OH)2,
will be mixed together while the conductivity apparatus continually tests the conductance of the
mixture. By observing the light bulb, you will see evidence that a chemical change is occurring,
and you can predict the identity of the products based on their electrolyte nature.
Materials
A list of substances to be tested is given on the report sheet. Your instructor may do
this experiment as a demonstration or individual conductivity meters may be given.
Procedure
1. Observe the brightness of the light bulb as your instructor inserts the test apparatus into each
test sample. Decide if the bulb is bright, dim or not lit at all. Interpret these observations and
decide if the test sample is a strong electrolyte, a weak electrolyte or a nonelectrolyte.
Record your observations and conclusions in your notebook.
2. Observe and record the brightness of the light bulb when testing each of the two reactants,
H2SO4 and Ba(OH)2, separately and during the mixing of the two compound solutions. Be
alert to small gradual changes in the bulb’s light output. Write a brief explanation of these
changes. Write the complete balanced equation for the reaction that takes place.
Safety
Wear your safety goggles during the demonstration.
Do not touch the electrodes at any time. Do not plug in or unplug the conductivity
apparatus yourself. If you touch both electrodes at the same time while the apparatus is
plugged in, a severe electrical shock may result.
Svante Arrhenius, in the 1880’s , determined with great
precision that ions contained charges which were specific
for each element in solution, and varied between elements
in exact multiples. So – Mg always had precisely the amount
of positive charge as Na, etc.
The Theory of Ionization contributed to the understanding of
the true nature of electric charge.
He also investigated kinetics and developed the formula for the
energy of activation in reactions, and proposed that meteorites
could have carried simple life forms to start life on Earth.
EXPERIMENT 8: REPORT
Ionization and Electrolyte Behavior
Section ___________
Electrolyte Behavior
Sample
Substance
#
1
Deionized water
Conductivity (bright,
dim or none)
2
Tap water
3
Solid NaCl
4
1.0 M NaCl(aq)
5
Solid sugar, C6H12O6
6
1.0 M sugar, C6H12O6(aq)
7
HC2H3O2, glacial
8
1.0 M HC2H3O2(aq)
9
1.0 M HCl(aq)
10
1.0 M NaOH(aq)
11
1.0 M NH3(aq)
12
1.0M NH4Cl (aq)
13
1.0 M Ba(OH)2(aq)
14
1.0 M H2SO4(aq)
15
Saturated BaSO4(aq)
observations from
procedure 2:
Ba(OH)2 + H2SO4
?
Electrolyte class (strong,
weak or nonelectrolyte)
Questions:
1. Explain your observations as Ba(OH)2 solution was added gradually to the H2SO4 solution.
Identify the products of the reaction and write a complete balanced equation.
2. Why is there a difference in the conductivity of pure solid NaCl and of the 1.0 M NaCl
solution?
3. Pure (glacial) acetic acid doesn’t conduct electricity while 1.0 M aqueous acetic acid does,
somewhat. Write an equation to show how acetic acid reacts with water to produce ions. (One
of those ions is the hydrated H+ ion, written as H3O+.)
4. The 1.0 M acetic acid and 1.0 M hydrochloric acid solutions had the same concentration of
molecules when the solutions were made. Do they have the same concentration of ions in the
solutions that we tested? What is the evidence?
5. Did you see similar conductivity of the solutions of the two bases, sodium hydroxide and
ammonia? Explain