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Na+
CO
3
2 –
Na+
Na+
CO
3
2 –
Na+
H
H
O
H
H
O
H
H
O
Na+
H
H
O
CO
3
2 –
H
H
H O
H
O
Experiment: Ionic Solutions (Electrolyte Solutions)
Introduction
Molecular compounds are made up of molecules, while ionic compounds are made up of ions. Ions are
different from molecules, as they have a charge. In an ionic compound, the number of positively
charged cations and negatively charged anions are such that charges are balanced. For example, in the
diagram below, note that there are two sodium cations (+1) to balance the charge of each carbonate
anion (-‐‑2).
Many ionic compounds dissolve in water; many do not. If an ionic compound dissolves in water, it
separates into individual charged ions. For example, when the soluble compound sodium carbonate
dissolves in water, the partial negatively charged side of the polar water molecules surround the
positively charged sodium ions, while the partial positively charged side of the polar water molecules
surround the negatively charged carbonate ions. The resulting solution is composed of separate
sodium ions and carbonate ions surrounded by water molecules.
The following chemical equation communicates how the soluble ionic compound, sodium carbonate,
separates into sodium ions, and carbonate ions. The notation “(aq)” means “aqueous” or that the ion is
dissolved in water. Note that water is not written as a reactant , but over the reaction arrow.
H
2
O
Na
CO
(s) à 2 Na
(aq) + CO
(aq)
Once ionic compounds are dissolved, the ions in solution may undergo further chemical reactions with
other substances, including neutralization, precipitation, oxidation-‐‑reduction, and other reactions.
Adapted with permission from Cascadia Community College
One technique that can be used to detect the presence of ions is conductivity, since charges in motion
conduct electricity. Soluble ionic compounds form solutions containing mobile ions that conduct
electricity and are therefore referred to as electrolytes. In contrast, insoluble ionic compounds do not
conduct electricity and are called nonelectrolytes because no separate ions are formed in solution.
Beyond being used to classify electrolytes and nonelectrolytes, conductivity is proportional to the
concentration of ions, so it can also be used to determine the actual concentration of ionic compounds
in water. Conductivity testing is simple, sensitive, and rugged/inexpensive equipment can be used.
For these reasons it is used for a wide variety of field and industrial analyses.
Molecular compounds are not made up of charged particles; therefore, they cannot conduct electricity
and are nonelectrolytes, like insoluble ionic compounds. However, there is an important class of
molecular compounds – even though not made up of ions – that can form ions via a chemical reaction
when they dissolve in water. If each molecule separates into ions, the compound is called a “strong
electrolyte”, but if the molecules of a compound produce only a few ions, it is called a “weak
electrolyte”. Soluble ionic compounds are also considered “strong electrolytes.”
For electrolytes, conductivity depends on concentration. In this lab you will measure the conductivity
of a solution with some initial concentration, and then you will dilute the solution by adding solvent.
The concentration of the original solution and diluted solution is determined by the following
equations:
Original solution: the initial concentration, C i
!
Diluted solution: the final concentration, C f
!
!
!
!
therefore, 𝐶 !
!
! !
! !
where C f and C i are the final and initial concentrations,
V i is the initial volume and V f is the final volume. (Notice that the units for V will cancel.)
Objectives
In this experiment, you will
ü Classify substances as strong, weak, or non-‐‑electrolytes.
ü Use conductivity to observe the process of dissolving an ionic compound.
ü Learn and practice the technique of dilution.
ü Observe the relationship between concentration of an ionic substance and conductivity.
Hazards
N Hydrochloric acid can cause chemical burns on the skin and damage eyes. Wear goggles and
wash your hands after using.
neighbors if you cannot find this.) Add the NaCl to the DI water in the beaker and watch the trace
on the screen while the NaCl dissolves. Sketch this trace in the data section of the report sheet.
This is now your original NaCl solution.
- When the conductivity becomes almost constant, record the final conductivity value. Click “Stop”,
remove the probe, and rinse it with distilled water into the waste beaker.
- Remove your NaCl solution from the stirring plate, remove the spin bar with tweezers and rinse
the bar with distilled water. SAVE your solution for the following steps:
- Pour between 20 and 25 mL of the original NaCl solution into a 50 mL graduated cylinder. Read
and record this volume as the “initial” volume (Vi) of the original NaCl solution to the nearest 0.
mL. Then add deionized water to the cylinder to a total volume of between 40 and 45 mL. Record
the “final” volume (Vf), to the nearest 0.1 mL. Pour the diluted solution into a new, dry 100 mL
beaker. (Why a dry beaker?)
This is now your diluted NaCl solution.
- Immerse the conductivity probe in the diluted solution and record the displayed conductivity
value.
- Discard the NaCl solutions and rinses in the sink. Rinse all of the glassware and the conductivity
probe and put the equipment away. Return the magnetic stir bar to your instructor.
Report Name _____________________Section______
Ionic Solutions
Lab Partner ____________________________
Data
Part A: Ionic Solutions
Compound
Conductivity Values (μS)
(Record data for all trials.)
Strong, Weak, or
Nonelectrolyte?
Many, few or no ions
produced in water?
NaCl
sodium chloride
CaCl
2
calcium chloride
HCl
hydrochloric acid
CH
3
COOH
(C
2
H
4
O
2
acetic acid
HOCH
2
CH
2
OH
(C
2
H
6
O
2
ethylene glycol
CH
3
OH
methanol
Part B: Conductivity Analyses
Data for the Original & Diluted NaCl Solutions (use significant figures!)
total volume of DI water for original NaCl solution, in mL
mL
total volume of DI water for original NaCl solution, in L
L
mass of dry NaCl used for original NaCl solution
g
calculated concentration of original NaCl solution (C
i
𝐶 !
=
𝑚𝑎𝑠𝑠 𝑜𝑓 𝑑𝑟𝑦 𝑁𝑎𝐶𝑙 (𝑖𝑛 𝑔𝑟𝑎𝑚𝑠)
𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 (𝑖𝑛 𝐿𝑖𝑡𝑒𝑟𝑠)
g/L
volume of original NaCl solution used (“initial” volume, V
i
mL
volume of diluted NaCl solution obtained (“final” volume, V
f
mL
calculated concentration of diluted NaCl solution (C
f
Since 𝐶
!
!
!
!
!
!
! !
! !
g/L
Post-lab Questions
- HCl is a covalent (molecular) compound in the gas phase. Does your data indicate that HCl
behaves as molecules or ions when dissolved in water? Explain your conclusion.
- Write a chemical equation that communicates what solid CaCl 2 forms when dissolved in water.
Note that water is not a reactant here. Use physical states in your equation as appropriate: (s),
(l),(aq)
- The following are beakers of water. Water molecules are already drawn in the beaker. Fill in the
ions or molecules present when each of the following substances is dissolved in water. Use spheres
to represent atoms/ions/molecules, and include a legend or labels with their chemical formula.
(Look at the example on pg. 1.) Use at least 4-‐‑5 spheres (molecules or ions) for each drawing.
HBr (a strong electrolyte) HF (a weak electrolyte) CH 3 OH (a non-‐‑electrolyte)
- Using your graph, estimate the conductivity of a 2.4 L solution that has 1.5 g of NaCl dissolved in
it?
a. Estimated conductivity: _______________________ include units!
b. Describe how you obtained the result.
- If you measured the conductivity of pond water and found it was 2000 μμS/cm, what concentration
of NaCl would you expect?
a. Estimated concentration: _______________________ include units!
b. Let’s say a 1%(m/v) NaCl solution is “salty” (2.5g NaCl in 250 mL water). Would the water
in the pond taste “salty”? State “yes” or “no”, and EXPLAIN. Show your work.
