ELECTROLYTES AND NONELECTROLYTES
CHEM 151
Fall 2009
Fill-in
Name______________________
Pre-Lab attached (p 14)
Stamp Here
Lecture Instructor ________________
Partner ____________________
Date ______________
LEARNING OBJECTIVES: After completing this experiment, you should:
•
Be able to use a conductivity meter to determine the conductivity of a substance.
•
Make accurate predictions about the conductivity and electrolytic capacities of various
substances.
•
Be able to write balanced dissolution reactions.
•
Be able to write balanced precipitation reactions.
•
Be comfortable setting up the apparatus for and performing a titration.
TO EARN YOUR FINAL STAMP: The following items must be completed in lab. You may
complete the entire assignment in the lab; this reflects the minimum required to earn your final
stamp.
 Complete the data table (the first table) on p. 5.
 Answer all questions on pp. 6-9.
INTRODUCTION
Electrolytes are substances that consist of charged particles called ions. When electrolytes are
dissolved in water (or other polar solvents) they ionize into positive (cation) and negative (anion)
ions. In this experiment, you will explore what types of compounds can become electrolytes, what
determines electrolyte strength, and how electrolytes are involved in the conduction of electricity.
Discussion
Before we can discuss the properties that make a compound an electrolyte, we must first understand
something about the properties of electricity.
Electricity results from the movement of charged particles through a conductor. The charged
particles can be either electrons or ions (positive or negative). In some cases, both types of particles
can be involved. When the movement of electricity is through a metal, the electrons move from one
metal atom to another which serves as the means for carrying the charge in the electrical circuit. If a
liquid is included as part of the electric circuit, something must carry the charge through this solution
otherwise no electrical current will flow. When a non-electrolyte is added, no ions are formed in
solution and therefore, no current flows. If we instead put a light bulb in the electrical circuit** (a
beaker containing a polar solvent, such as water), it is possible to tell whether the compound being
added to the solvent is an electrolyte or a non-electrolyte by whether or not the light bulb lights up.
(**Note the set-up in the fume hood. The bulb assembly clamps to a ring stand.)
#4 Electrolytes and Nonelectrolytes
Rev F09NF Fall 2009
Page 1 of 14
Determining the type of bonds present (Ionic, Polar Covalent or Nonpolar Covalent)
The percent ionic character of a bond is dependent on the differences in electronegativities (see sec.
9.6, Tro, 1st ed. and Figure 9.10) of the atoms present, and the arrangement of these atoms in the
molecule. When electronegativity differences are very large (>2.0) a bond is considered ionic; the
electrons are completely transferred from one atom to another. When the electronegativity
difference is intermediate (2.0 > x > 0.5) the bond is considered polar covalent, and regions of
partial positive and partial negative charges are formed. When the electronegativity difference is
small (< 0.5) then the molecule is considered nonpolar covalent. Although we have not covered
these concepts in class yet, these trends can be estimated by the position of the two atoms on the
Periodic Table. When a compound is formed from elements that are on opposite sides of the
Periodic Table, such as from the s-block (metal) and the p-block (nonmetal), that compound is
typically ionic. When a compound is formed from elements all within the p-block (all nonmetals),
that compound is typically covalent. Hydrogen is often an exception – it is often considered a
nonmetal in bonding. Use the flowchart below to help you decide how to classify individual
molecules as ionic, polar covalent, or nonpolar covalent.
Is a metal, or
NH4+ present?
Yes, then the
compound is Ionic
No, then the compound
is Covalent
Polar covalent; contains ANY two
different non-metals bonded
together
(except C-H which are defined as
non-polar)
Non-polar covalent; contains
only C-H bonds, or all the same
element (e.g. O2)
#4 Electrolytes and Nonelectrolytes
Rev F09NF Fall 2009
Page 2 of 14
Determining/Predicting Electrolyte Behavior
There are three types of electrolytes: strong, weak and non-electrolytes. A polar solvent is
necessary for an electrolyte to function:
1. Ionic compounds that are soluble are strong electrolytes. (Review solubility rules
on p. 5 of the lab, or in your text). Ionic compounds that are insoluble are nonelectrolytes.
2. Polar covalent compounds can be strong, weak or non-electrolytes. Most acids
have weak polar covalent bonds to the hydrogen, allowing water to pull the H+ off
of the molecule – thus acids will be an electrolyte – either strong or weak. (strong
acids are strong electrolytes, weak acids are weak electrolytes)
3. Non-polar covalent compounds form non-electrolytes.
For example, in Figure 2 the ionic compound NaCl, when dissolved in a polar solvent, dissociates
(splits apart) completely to form sodium ions and chloride ions. These ions form because the
electrostatic forces present in the polar solvent help to stabilize the ions.
Figure 2: The dissolution of NaOH
NaOH
Polar
Solvent
Na+
δ-
Na+(aq) + Cl-(aq)
H
O
δ+
H
δ+
δ-
H
δ+
Cl-
O
H
δ+
The following flow-chart helps you classify water-soluble compounds for electrolyte
behavior:
Water-soluble
compound?
YES
Ionic?
NO
YES
Probably
strong
electrolyte
NH3 or other
molecular
base?
Acid?
NO
Strong Acid?
YES
Strong
electrolyte
#4 Electrolytes and Nonelectrolytes
YES
NO
Weak
electrolyte
Weak
electrolyte
Rev F09NF Fall 2009
NO
Probably
nonelectrolyte
Page 3 of 14
Ionic compounds are strong electrolytes in polar solvents if they are soluble; they form large
numbers of ions that can carry a large electrical current.
Polar covalent compounds may or may not form ions when dissolved in a polar solvent, depending
on the compound.
Strong acids (polar covalent), (e.g. HNO3, H2SO4, HClO4, HCl, HBr, HI), when dissolved in polar
solvents, will form large numbers of ions and therefore form strong electrolytes in water solution.
H2SO4 (aq) → H+1 (aq) + HSO4-1(aq) → 2 H+ (aq) + SO42- (aq)
100% ionized
Weak acids and weak bases (i.e., ammonia, carbonic acid) will only partially ionize (partially
dissociate or react to form ions) in a polar solvent. This partial dissociation is denoted by
“equilibrium” arrows: . Most of the compound remains in solution as non-ionized molecules.
Since there are very few ions in solution, only a small amount of the electrical current can flow. The
light bulb will light up, but only dimly, and is very dependent on the concentration of the
compounds. These compounds are weak electrolytes.
H2CO3 (aq)
HCO3-1 (aq) + H+1(aq)
99% molecules
1% ions
All nonpolar covalent and many weakly polar covalent compounds stay as molecules in solution
and do not form ions. They are unable to carry a charge and therefore are non-electrolytes.
If the solvent is nonpolar, it does not have the electrostatic forces (ion/dipole interactions) to
stabilize ions and so none form. Therefore in a non-polar solvent, all compounds which are added to
it are non-electrolytes.
Also, a solvent may be very polar, but if there is no compound added to it to provide ions or support
the formation of ions, the solvent itself will not conduct electricity.
#4 Electrolytes and Nonelectrolytes
Rev F09NF Fall 2009
Page 4 of 14
Summary of Bond Types
1. Ionic: Contains both metals and nonmetals, or the ammonium (NH4+) ion. Elements are
greatly separated on the Periodic Table.
2. Covalent: Contains only nonmetals (including hydrogen).
a. Polar Covalent: typically contains H-O or H-N bonds, strong acids. Elements have
intermediate separation on the periodic table (separated by two or three other
elements).
b. Nonpolar Covalent: typically contains lots of H-C bonds, or all atoms are the same
element. Elements are right next to each other on the Periodic Table.
Summary of Electrolytes in Water
1. Strong electrolytes: a. soluble ionic compounds that dissociate (almost) completely into
ions.
b. strong bases (NaOH, LiOH, KOH, RbOH, CsOH, Ca(OH)2,
Ba(OH)2, and Sr(OH)2 dissociate almost completely into ions.
c. strong acids (polar covalent molecules) that dissociate almost
completely into ions.
2. Weak electrolytes: a. polar molecules that dissociate partially into ions, eg., weak acids,
weak bases.
b. slightly soluble salts that dissociate partially into ions.
3. Non-electrolytes:
a. nonpolar (covalent) compounds, that do not dissociate into ions.
b. very weakly polar compounds, that do not dissociate into ions.
Solubility Guidelines
Ion(s)
Li+, Na+, K+, NH4+
C2H3O2-, NO3-, ClO3-,
ClO4Cl-, Br-, I-
Rule
Group IA and ammonium salts are all
soluble.
Acetates and nitrates are soluble.
Exceptions
none
Most halides are soluble.
Salts containing Ag+, Hg22+, or Pb2+ are
insoluble.
CaSO4, SrSO4, BaSO4, Ag2SO4, Hg2SO4,
PbSO4 are insoluble.
Group IA and ammonium carbonates are
soluble.
Group IA and ammonium phosphates are
soluble.
Group IA and ammonium sulfides, CaS,
SrS and BaS are soluble.
Group IA and ammonium hydroxides,
Ca(OH)2, Sr(OH)2, Ba(OH)2 are soluble.
SO42-
Most sulfates are soluble.
CO32-
Most carbonates are insoluble.
PO43-
Most phosphates are insoluble.
S2-
Most sulfides are insoluble.
OH-
Most hydroxides are insoluble.
#4 Electrolytes and Nonelectrolytes
Rev F09NF Fall 2009
none
Page 5 of 14
Molecular, Ionic and Net Ionic Equations
When we write a reaction, how can we represent the reactants and products? Are there some ions
that do not participate in the actual reaction? If so, can we write an equation that reflects only the
reacting species? The example below shows a variety of different ways to represent a reaction.
Let’s consider the reaction between lead (II) nitrate and sodium hydroxide, both of which are soluble
in water and are thus electrolytes. When these two compounds react, they form insoluble lead (II)
hydroxide and sodium nitrate. We can write the molecular equation that is described by this
process:
Molecular Equation: Pb(NO3)2(aq) + 2 NaOH(aq) → Pb(OH)2(s) + 2 NaNO3(aq)
We can also express this reaction in words, where ‘+’ changes to ‘reacts with’ or simply ‘and’, and
the arrow becomes ‘yields’ or ‘produces’:
Word Equation:
lead (II) nitrate reacts with sodium hydroxide to produce lead (II) hydroxide and
sodium nitrate.
Note that phases of matter are not noted in the word equation, nor is reaction stoichiometry. It is
assumed that the reader knows nomenclature and solubility rules, and will be able to translate the
names of the compounds, will be able to determine solubility, and will be able to balance the
equation.
To focus in on the specific chemicals that directly react, the net ionic equation can be useful. To
generate a net ionic equation, an ionic equation is generated by systematically splitting each aqueous
compound into it s component ions. Note how the superscripts and stoichiometric coefficients are
treated. Any solids, liquids or gases remain in their given form – only aqueous species split! Any
ion that show up on both sides of the equations can be cancelled out as spectator ions, and the
remaining ions and compounds go into the net ionic equation.
Molecular:
Pb(NO3)2(aq) + 2 NaOH(aq) → Pb(OH)2(s) + 2 NaNO3(aq)
Ionic:
Pb+2(aq) + 2 NO3-(aq) + 2 Na+(aq) + 2 OH-(aq) → Pb(OH)2(s) + 2 Na+(aq) + 2 NO3-(aq)
Net Ionic:
Pb+2(aq) + 2 OH-(aq) → Pb(OH)2(s)
In Part 6 of the experiment, you will observe what happens when the electrolyte, Ba(OH)2, is reacted
with a second electrolyte, H2SO4 The reaction between these two compounds is an acid/base
reaction and a precipitation reaction that produces water and the insoluble salt, barium sulfate.
word equation:
barium hydroxide + sulfuric acid  barium sulfate + water
molecular equation: Ba(OH)2 (aq)
+ H2SO4 (aq)
→
BaSO4 (s)
+ 2 H2O (l)
strong base
strong acid
insoluble salt
water
strong electrolyte
strong electrolyte
non-electrolyte non electrolyte
+2
+
-2
ionic equation:
Ba (aq) + 2 OH (aq) + 2 H (aq) + SO4 (aq) → BaSO4 (s) + 2 H2O (l)
net ionic equation: Ba+2 (aq) + 2 OH- (aq) + 2 H+ (aq) + SO4-2 (aq) → BaSO4 (s) + 2 H2O (l)
Water is polar covalent and forms very few ions in solution. Also, since BaSO4 is insoluble in
water, it is a non-electrolyte. Therefore, there is a point in the titration when equal stoichiometric
amounts of H2SO4 and Ba(OH)2 will be added and there are no ions to carry any charge. At this
point the light will go out. Addition of more H2SO4 after this point will supply more ions. Since
there is no more Ba(OH)2 available to react with these ions, the H+ and HSO4- ions are available to
carry the current.
#4 Electrolytes and Nonelectrolytes
Rev F09NF Fall 2009
Page 6 of 14
Experimental Procedure
CAUTION: 1) It is extremely important to use safety goggles for this experiment and to use
caution in the handling of acids and bases.
2) It is also important to unplug the conductivity tester when cleaning the
electrodes or when not using the tester so that your body does not succeed in
completing the circuit which will result in a nice jolt of electricity coursing through
your body!
Technique tip: Rinse the electrodes between solutions by unplugging the conductivity tester and
using a water wash bottle to rinse the electrodes into a waste beaker.
1.
Obtain a set of electrolyte solution wide-mouth jars. You will use these solutions as-is,
there is no need to transfer them to another container. Do not throw out the solutions
when you are finished, return them to the supply bench. Please share these solutions
with other groups! Do not cross-contaminate these solutions! Obtain a water wash
bottle and a medium size beaker large enough to fit under the conductivity tester, to use
as a ‘Waste Beaker’. Set up the conductivity tester at your bench (shown in Figure1), to
determine if the compound behaves as an electrolyte or a non-electrolyte in that solution.
Make sure to completely submerge the electrodes into the solutions before plugging in
the conductivity tester. Also, using your observations and rules in Appendix A attached
and/or the Laboratory Handbook, Appendix B, Tables 3 and 4, determine whether each
compound is ionic, polar covalent or non-polar covalent. The structures of ammonia,
toluene, acetic acid and methanol are:
CH3
N
H
H
H
Ammonia
O
H3C
Toluene
H
O
H
Acetic Acid
H3C O
Methanol
Data:
Solution Tested
(name)
Distilled water
Formula
Conductivity
(strong, weak
or non)
Type of Bonds in Compound
(ionic, polar covalent, or
nonpolar covalent) *refer to chart, p.
2
H2 O
0.1 M sodium chloride
0.1 M methanol (CH3OH)
0.1 M sodium hydroxide
0.1 M barium hydroxide
0.1 M ammonia (NH3 / NH4OH)
0.1 M hydrochloric acid
0.2 M acetic acid (CH3COOH)
0.1 M sulfuric acid
#4 Electrolytes and Nonelectrolytes
Rev F09NF Fall 2009
Page 7 of 14
For each solution tested, write a chemical equation representing the dissociation behavior of each
compound in water solution. If no dissociation occurs, write NR for No Reaction. Include charges
and phase labels. The first two are started for you as an example of how to structure your equations.
Water
H2O (l) →
sodium chloride
NaCl(aq) 
Methanol
CH3OH (aq) 
sodium hydroxide
barium hydroxide
ammonia(a weak base)
NH3 (aq) + H2O (l)
NH4+1(aq) + OH-1 (aq)
hydrochloric acid
acetic acid
CH3COOH (aq)
sulfuric acid
2.
In the fume hood**, there are two beakers with solid salts in them. Beaker A contains solid
sodium chloride. Beaker B contains solid calcium phosphate. Using these solids, and the
accompanying light-bulb set-ups, perform a conductivity test on them.
a.
Is the solid sodium chloride a good conductor?
b. Is the solid calcium phosphate a good conductor?
Yes
No
Yes
No
Using your answers above, explain why an ionic solid is (or is not) a good conductor of electricity:
_________________________________________________________________________________
_________________________________________________________________________________
_________________________________________________________________________________
**
note: the location of the solid salts might have to change depending on space limitations. Ask the lab instructor if you cannot find
them in the hood.
3.
Return to your bench area. Obtain two 100-mL beakers and fill each with ~ 50 mL of distilled
water. Weigh out 0.5 grams of each solid. Into one beaker place the 0.5 grams of sodium
chloride. Into the other beaker, place the 0.5 grams of calcium phosphate. Record your
observations of each solid in the water. Now, test each solution with your conductivity set-up.
a.
Is the sodium chloride in water a good conductor?
b. Is the calcium phosphate in water a good conductor?
#4 Electrolytes and Nonelectrolytes
Rev F09NF Fall 2009
Yes
No
Yes
No
Page 8 of 14
Briefly explain what is happening in solution for both solids that allows them to conduct or not
conduct electricity:
_________________________________________________________________________________
_________________________________________________________________________________
4.
In the fume hood, ~20 mL of glacial acetic acid (100% hydrogen acetate) is already in a dry
250 ml beaker. Using this solution, and the accompanying lightbulb set-up, perform the
conductivity test. Caution: Glacial acetic acid has harmful fumes. Avoid breathing vapors
and avoid contact with skin. Obtain your own 100-mL beaker and fill it with ~ 25 mL
distilled water. Now pour about 10 mL of the glacial acetic acid (from the stock bottle) into
your beaker and perform the conductivity test. Using a squirt bottle add more distilled water to
the beaker while testing the conductivity. Note what happens as you add more water. Dispose
of in the appropriate electrolytes waste container.
a. Is glacial acetic acid (i.e., 100% acetic acid) a good conductor?
Yes
No
b. Does the diluted acetic acid give a greater conductivity than the pure compound? (circle
one)
Greater conductivity
Less conductivity
About the same
c. Explain, on the molecular level, what happens to the glacial acetic acid when water is
added.
_____
_____
______
5. The solution is already made up and is in a labeled container. In the fume hood, test the
conductivity of hydrogen chloride dissolved in toluene. Do not dump out or contaminate
this solution. Leave the solution for the next group.
a. Is the HCl in toluene an electrolyte?
Yes
b. Is the HCl in water (part 1 of the lab) an electrolyte?
No
Yes
No
c. Therefore, what happens to the HCl molecules when it is dissolved in water?
_________________________________________________________________
d. In toluene, does the HCl exist as ions or as molecules – how do you know?
_____________________________________________________________________
_____________________________________________________________________
_____________________________________________________________________
#4 Electrolytes and Nonelectrolytes
Rev F09NF Fall 2009
Page 9 of 14
6. Obtain a magnetic stirrer (under the hoods), magnetic stirbar (the stirbar looks like a giant white
“pill” and is in a beaker on the stock bench – if you can’t find any there is a stir bar “lifeguard”
which looks like a long white wand hanging by the fume hood. Obtain the “lifeguard”, stick it in
the titration waste container, swirl it around, and you will rescue a stir bar from drowning! –
rinse it off into the waste container with distilled water), buret, buret clamp, conductivity tester
and a 250 mL beaker to perform this part of the experiment. Set up the apparatus with the beaker
on the magnetic stirrer, with the stir bar inside (Figure 2). You will be using reagents from large
stock bottles located on the lab bench. Do NOT use the same solutions that were used for the
conductivity tests. Place in the beaker about 30 mL of ~ 0.1 M barium hydroxide solution (from
the large stock bottle, record the actual amount below) and dilute it with 50 mL of distilled
water. Prepare the buret first by rinsing the buret with water and 2 small (5-7 mL!!) portions of
the sulfuric acid solutions before titratin. Suspend the buret using the buret clamp and ringstand
so that you can fill the buret. Fill the buret full of dilute, ~0.1 M, sulfuric acid (again, the sulfuric
acid is in a large stock bottle). Open the buret, allow some to drain, as this fills the tip of the
buret. Turn on your stirplate so that a mild swirling action occurs. (your stir-bar should not be
popping around like popcorn!) Immerse the electrodes of the conductivity tester into the barium
hydroxide solution. You will be titrating the acid into the base once the experiment begins
(answer questions 1 and 2 below before beginning!). And you will be testing the conductivity
continuously while titrating. Record your observations about the solution and conductivity
in each section below.
Figure 2
Questions:
Describe in words and write net ionic reactions to demonstrate the electrolyte behavior, include
all ions and molecules present. Show any dissociation reactions. Be sure to include phases!!
Volume of Ba(OH)2 _______ mL in beaker.
1. Write a chemical equation representing the dissociation behavior of Ba(OH)2 in water
_____________________________________________
Would you expect the Ba(OH)2 solution to conduct electricity – why or why not?:
______________________________________________________________________________
______________________________________________________________________________
#4 Electrolytes and Nonelectrolytes
Rev F09NF Fall 2009
Page 10 of 14
Write a chemical equation representing the dissociation behavior of H2SO4 in water
____________________________________________
Would you expect the H2SO4 solution to conduct electricity- why or why not?:
______________________________________________________________________________
______________________________________________________________________________
2. Write balanced molecular, ionic, and net-ionic equations for the reaction of sulfuric acid with
barium hydroxide
Molecular: ____________________________________________________________________
Ionic: ________________________________________________________________________
Net: ionic: ____________________________________________________________________
3. Titrate approximately 17 – 18 mL of H2SO4 to the solution in the beaker. Describe what
happens to the solution, and to the light bulb, after 17-18 mL of H2SO4 is titrated into the
beaker.
Visually – how does the solution in the beaker change in appearance? _____________________
______________________________________________________________________________
Visually – how does the brightness of the lightbulb change? _____________________________
What ion(s) are present in solution? _____________________________________
Explain from a chemical perspective (accounting for the species in the solution as well as their
“relative” amounts) why the lightbulb has changed in brightness.
______________________________________________________________________________
______________________________________________________________________________
______________________________________________________________________________
#4 Electrolytes and Nonelectrolytes
Rev F09NF Fall 2009
Page 11 of 14
4. Add additional sulfuric acid until the light bulb goes out. This is the “endpoint” of the
titration.
Volume of H2SO4 when the light goes out: ______________
Visually – how does the solution in the beaker change in appearance? _____________________
______________________________________________________________________________
What ion(s) are present in solution? _____________________________________
Explain from a chemical perspective (accounting for the species in the solution as well as their
“relative” amounts) why the lightbulb has gone out.
5. Add more sulfuric acid until the light bulb comes back on brightly.
Volume H2SO4 when light is bright: _______________
What ion(s) are present in solution? _____________________________________
Explain from a chemical perspective (accounting for the species in the solution as well as their
“relative” amounts) why the lightbulb has turned back on.
#4 Electrolytes and Nonelectrolytes
Rev F09NF Fall 2009
Page 12 of 14
Homework Problems – these can be done at home – they do not need to be completed to get a
final stamp!
1. A solution of acetic acid is made by dissolving 0.100 mole of CH3COOH, in enough water to
make one liter of solution. Conductivity measurements show that the acetic acid is 1.34%
ionized. Remember: Avogadro's number is 6.022 x 1023 molecules (or ions, or atoms, or
students, or anything)/mole. Write the dissociation equation for acetic acid in water and then
calculate a and b.
Equation: ____________________________________________________________________
a. If there is 0.100 mole of CH3COOH dissolved in 1 L of total solution, please write the
concentration (in M) of the acetic acid
answer: _______________
b. If 1.34% of the acetic acid is ionized, what % was not ionized and remains as molecules?
answer. _______________
c. Calculate, using stoichiometry and your knowledge of %, the number of acetate ions,
CH3COO-1 (aq)
answer. _______________
d. Using your answer in part b, calculate the number of hydrogen acetate molecules,
CH3COOH, in the solution.
answer. _______________
2. How many mL of 0.200 M H2SO4 solution must be added in order to react completely with 25.0
mL of 0.100 M NaOH?
Balanced Equation:
________________________________________________________
Calculations:
answer. ______________mL
#4 Electrolytes and Nonelectrolytes
Rev F09NF Fall 2009
Page 13 of 14
Stamp:
Pre-Lab Assignment
1. Name the following compounds. Classify the following compounds as containing polar
covalent, nonpolar covalent, or ionic bonds. What is the conductivity in solution (strong, weak,
or nonelectrolyte)? See the flowchart on p. 2 of the lab for help!
Name
Type of bond
Conductivity
NaBr
______________________________
___________________ ___________________
LiOH
______________________________
___________________ ___________________
HBr (aq)
______________________________
___________________ ___________________
HF (aq)
______________________________
___________________ ___________________
N2
______________________________
___________________ ___________________
2. Name the following ionic compounds. Classify them as soluble or insoluble in water using your
solubility guidelines. What is the conductivity in solution (strong or non-electrolyte)?
Name
Soluble/Insoluble
Conductivity
NaBr
______________________________ ___________________ ___________________
Ba(OH)2
______________________________ ___________________ ___________________
Ca3(PO4)2 ______________________________ ___________________ ___________________
MgCl2
______________________________ ___________________ ___________________
3. Consider the following titration reaction: HCl (aq) + AgNO3 (aq) → AgCl (s) + HNO3 (aq)
a. Write the ionic equation for the reaction.
b. Write the net-ionic equation for the reaction.
c. At the equivalence point, when exactly 0.050 mol of HCl (aq) has reacted with 0.050 mol of
AgNO3 (aq), would you expect the solution to conduct electricity? Explain your choice.
4. Why is it important to unplug the conductivity tester while cleaning the electrodes?
#4 Electrolytes and Nonelectrolytes
Rev F09NF Fall 2009
Page 14 of 14