4. Determining the Chemical Formula
of an Ionic Compound
What you will accomplish in this experiment
You’ve learned that compounds are a chemical combination of elements, meaning that they’re created when two or
more elements chemically react with one another. The driving force for that reaction is that the elements are trying
to achieve the stability of a noble gas electron configuration (for most elements, this means an s2p6 “octet” of
electrons.
When the reacting elements are a metal and a nonmetal (elements with an electronegativity difference greater than
2.0), the metal atom achieves its octet when it loses one or more electrons (to become a positively-charged ion, a
“cation”), and the nonmetal atom gains those electrons (to become a negatively-charged ion, an “anion”).
The chemical formula for the ionic compound that’s formed describes the fixed proportion in which vast numbers
of the oppositely-charged ions combine in order to balance out the positive and negative charges.
This idea is illustrated below for the chemical combination of the elements Sodium and Chlorine to make the ionic
compound: Sodium Chloride, NaCl.
The ionic compound you’ll work with in this experiment is the chemical combination of the metal element Copper
(Cu) and the same nonmetal element as above, Chlorine (Cl). When these elements chemically combine, the fixed
proportion of ions in the compound can be expressed by the chemical formula, CuCl2, and by the compound’s
name, Copper (II) Chloride.
Your job will be to chemically separate these two elements, and then to determine the respective masses of Cu2+
and Cl- ions that were in your original sample of Copper (II) Chloride.
By converting these masses of each element into “moles” (using the molar masses of Copper and Chlorine), and
then determining the “mole ratio” of the two elements in the compound, you should be able to prove
experimentally that the chemical formula for Copper (II) Chloride is, in fact, CuCl2.
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Concepts you need to know to be prepared
A compound’s chemical formula is made up of the symbols (from the Periodic Table) that represent each of the
elements in that compound. And the numerical subscripts following these symbols indicate the fixed proportion in
which the elements combine when that compound is formed by a chemical reaction.
For an ionic compound (such as sodium chloride, NaCl), you
should NEVER think of that compound as just a PAIR of
oppositely-charged ions, as the chemical formula seems to imply.
If you look closely at the illustration of a sodium chloride (NaCl)
salt crystal to the right, you’ll see that it is definitely NOT just a
pair of Na+ and Cl- ions, but an extensive, orderly, threedimensional network of charged particles.
Each Na+ ion is attracted to the Cl- ions immediately above,
below, to the right, to the left, in front, and behind it. And each
Cl- ion is attracted to the Na+ ions surrounding it.
So it’s important to realize that the chemical formula for ANY
ionic compound does NOT describe a small grouping of ions, it
describes the fixed proportion in which vast numbers of the
oppositely-charged ions combine in order to appropriately
balance out the positive and negative charges.
So – one way to think about the chemical formula for Copper (II) Chloride, CuCl2, the ionic compound in this
week’s experiment, is that it expresses the “ION ratio” of the elements in the compound:
The compound CuCl2 contains 2 Cl- IONS for every 1 Cu2+ ION.
But a “bigger picture” way to think about that chemical formula is that it expresses the “MOLE ratio” of the
elements in the compound:
The compound CuCl2 contains 2 MOLES of Cl- ions for every 1 MOLE of Cu2+ ions.
One mole of any pure substance (element or compound) is really just an Avogadro’s number of “particles” of that
substance (where “particles” = atoms, ions, or molecules). For example:
1 mole of Copper metal = 6.02 x 1023 Cu atoms
Or in the case of Copper (II) ions:
1 mole of Copper (II) ions = 6.02 x 1023 Cu2+ ions
An Avogadro’s number of atoms for a particular element, one MOLE of the element, will ALWAYS have a mass in
GRAMS equal to the Atomic Weight listed for that element on the Periodic Table:
1 mole of Copper (Cu) = 6.02 x 1023 atoms of copper = 63.55 grams of Copper metal
Or for Copper (II) ions:
1 mole of Copper (II) ions = 6.02 x 1023 Cu2+ ions = 63.55 grams of Copper (II) ions
NOTE: The mass in grams of 1 mole of any pure substance is commonly referred to as the “Molar Mass.”
A similar unit relationship is easily determined for any other element on the Periodic Table. The first part of this
unit relationship is always the same:
1 mole of any element = 6.02 x 1023 atoms of that element
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The unit relationship is then completed by taking the Atomic Weight of the desired element from the Periodic
Table. For example:
1 mole of Chlorine (Cl) = 6.02 x 1023 atoms of chlorine = 35.45 grams of chlorine
Or for Chloride ions:
1 mole of Chloride ions = 6.02 x 1023 Cl- ions = 35.45 grams of chloride ions
These unit relationships can be used to convert the mass in grams of any substance to the number of moles
represented in that mass. And the unit conversion from grams of an element to moles of an element is essential to
this week’s experiment.
You’ll begin by obtaining a sample of copper (II) chloride and recording its mass to the maximum number of
significant figures provided by the balance.
You’ll then dissolve this compound in water. In dissolving, ionic compounds separate into their component ions.
So in water, copper (II) chloride “dissociates” into copper (II) ions (Cu2+) and chloride ions (Cl-). The “hydrated”
Cu2+ ions have a brilliant blue color, while the Cl- ions are colorless.
Into this solution, you’ll place a coiled piece of aluminum (Al) wire. This is aluminum metal: an element, not a
compound. The aluminum wire will spontaneous react with the Cu2+ ions in the solution, and the copper will be
converted into its elemental form: copper metal.
Think about what this means:
The Cu2+ ions will be gaining electrons from the Al atoms, and thus converted into atoms of Cu metal. The atoms
of Al metal are losing electrons to the Cu2+ ions, and are converted into colorless, water-soluble Al+3 ions.
And you’ll be able to observe this chemical transformation:
The brilliant blue color of the aqueous Cu2+ ions will eventually fade away, and brown copper metal will collect at
the bottom of your reaction beaker. The aluminum wire will become smaller as the Al atoms are converted into
aqueous Al+3 ions. This chemical reaction is slightly “exothermic,” meaning that heat is released as the reaction
progresses. You can “observe” this heat release in two ways: qualitatively – by simply noting the warmth of the
reaction container by touching it carefully with your fingertips, and quantitatively – by measuring the actual
temperature change of the aqueous mixture.
The assumption that you’ll make in this experiment is that every Cu2+ ion in solution will react with the aluminum
wire and be converted into an atom of copper metal. After the reaction goes to completion, you are to filter, dry,
and weigh this copper metal.
At this point, realize that you’ve chemically separated the two elements in the compound (copper and chlorine),
and you’ve determined the mass of copper that was in your original sample. By subtracting this mass of copper
from the mass of your original copper (II) chloride sample, you’ll know the mass of chlorine that was in that
sample.
Now remember your mole unit relationships:
You can use the molar masses of copper and chlorine from the Periodic Table to convert the mass in grams of each
element to the number of moles of that element.
So your final steps are to convert your experimental masses of copper and chlorine into moles, and then to
determine the mole ratio of chlorine to copper in the compound.
As discussed, the chemical formula of CuCl2 indicates that two moles of chloride ions combine with one mole of
copper (II) ions when the compound copper (II) chloride is formed: a mole ratio of two-to-one. Ideally, your
experimental result should be consistent with this fixed proportion.
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If it’s not, you’ll need to think back through the various steps of the procedure, identify the possible sources of
experimental error, and consider how each source of error might have contributed to making your experimental
result too high or too low.
You can see that it’s essential to do a good job of recording your data and observations, so that you’ll be able to
make a thorough assessment as to why your experimental result is higher or lower than the true value.
Procedure that you will follow
1. Measure a known mass of copper (II) chloride into a 250-mL beaker (approximately 3 grams, but weighed to
the maximum number of significant figures provided by the balance).
2. Add approximately 60 mL of distilled water to the beaker and stir until the copper (II) chloride has completely
dissolved. The resulting solution should be the brilliant blue color that is characteristic of Cu2+ ions in aqueous
solution.
3. Your lab instructor will provide you with an aluminum wire and demonstrate how to shape it into a coil. Take
a moment to measure the initial mass of the wire. You’ll then immerse the coiled end of the wire in the copper
solution. Maintain your grasp on the other end of the rod, so that you can use it to stir the reaction mixture.
4. Watch closely to observe the chemical reaction that occurs, and make detailed notes of your observations in
your lab notebook.
IMPORTANT: Be sure to review the Skills in Recording Observations provided in the Remedies document:

When recording observations of a substance, describe everything that you see (or perhaps smell): Is the
substance a solid, liquid, or gas? Is it a solution? What color is it? Is it opaque, translucent, or transparent? If
you are instructed to waft the vapors toward your nose, does the substance have an odor?

When recording observations of a reaction, first write down what you did (heat a solution, or add one chemical
to another and mix thoroughly). Then write down what you saw, heard, smelled, or felt with your fingertips:
Was there a color change? Did a gas evolve? Did a solid form? Was an odor emitted? Was there any sound?
Was any heat evolved? (If yes, monitor the temperature change with your thermometer, and record it as an
observation.) How long did it take for the reaction to occur?
A Note on Common Sense and Experimental Technique:

Because the chemical reaction between the Cu2+ ions and the Al atoms is taking place on the surface of the
aluminum wire, the copper metal that’s produced by the reaction will accumulate on that surface. So you’ll
need to use your metal spatula or glass stirring rod to scrape or knock the copper metal off the wire as the
reaction progresses. The goal is to keep exposing the aluminum surface to the solution so that the reaction can
continue all the way to completion (until all of the copper ions in the solution have been converted to copper
metal).
5. When the reaction appears to be complete, remove the aluminum wire, rinse and dry it thoroughly, determine
its final mass, and then return it to your lab instructor.
6. Now you’ll need to separate the copper metal from the reaction mixture by vacuum filtration. In preparation
for this procedure, pre-weigh a piece of filter paper (that fits in the Buchner funnel) with your watchglass. Set
up the vacuum filtration apparatus as illustrated in the LabCam video and described in the Remedies document.
Be sure to follow the steps carefully – especially securing the Buchner funnel to the filter flask with the rubber
adaptor, and clamping the neck of the filter flask to a ring stand. Ask your lab instructor to inspect your
vacuum filtration set-up before you begin filtration.
7. As shown in the video, you’ll need to place your pre-weighed filter paper in the Buchner funnel, pre-moisten it
with water, and make certain that the wet paper forms a seal against the funnel when the vacuum is turned on.
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Then pour the mixture into the center of the funnel. Remember that you can use small squirts of distilled water
from your plastic wash bottle to rinse the copper from the beaker into the funnel.
8. As the final step to the vacuum filtration, rinse the copper metal in the funnel with several squirts of acetone
from the plastic squeeze bottle provided in the hood. Acetone is a volatile organic solvent, so it will help the
copper dry more quickly. You should also allow the vacuum to pull air through the copper for several minutes,
to remove excess liquid and hasten drying.
9. Then carefully transfer the wet filter paper and copper metal from the funnel to the pre-weighed watch lass. Be
sure to spread the copper across the surface of the watchglass to enable its drying. Mark the watchglass with
your initials, and then place it in the heated oven to dry. When the drying and cooling is complete, weigh and
record the mass of the watchglass (containing the dried filter paper and copper). The mass of the copper metal
can be determined by difference.
10. Then by difference, determine the mass of chlorine in your original sample of copper (II) chloride.
11. Use Atomic Weights from the Periodic Table to convert your masses of copper and chlorine to moles, then
determine your experimental mole ratio of chlorine to copper.
IMPORTANT: You MUST dispose of all chemical waste as directed by your lab instructor. Do NOT put any
chemical waste in the laboratory sinks or garbage cans. Use the solid and liquid waste containers in the hood as
directed by your lab instructor.
ALSO IMPORTANT: Please thoroughly clean your laboratory glassware before replacing it in your equipment
drawer. Large Nalgene bottles of soap solution are provided near the laboratory sinks. You should notify your lab
instructor if the stock of soap solution is running low.
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Report Sheet 4: Chemical Formula of a Compound
Student ______________________________ Lab Partner__________________________ Date Lab Performed__________
Section #_________ Lab Instructor__________________________________________ Date Report Received ___________
Lab Notebook: Data and Observations
Experimental Data:
Mass of empty 250-mL beaker
_______________________
Mass of 250-mL beaker with copper (II) chloride sample
_______________________
Mass of copper (II) chloride sample (by subtraction)
_______________________
Mass of aluminum wire before reaction
_______________________
Mass of aluminum wire after reaction
_______________________
Mass of watchglass and filter paper (pre-weighed)
_______________________
Mass of watchglass, filter paper, and dried copper metal
_______________________
Additional Data:
Molar Mass of Copper
_______________________
Molar Mass of Chlorine
_______________________
Observations:



Initial physical appearances of the copper (II) chloride solid, aqueous solution, and the aluminum wire.
Chemical reaction of the aluminum wire with the copper solution, including the temperature change.
Final physical appearance of the copper metal, aluminum wire, and filtered solution.
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Formal Report: Results and Conclusions
Change in mass of aluminum rod
Calculation:
_______________________
Mass of copper metal produced by the reaction
Calculation:
_______________________
Mass of chlorine in original sample of copper (II) chloride
Calculation:
_______________________
Number of moles of copper metal
Calculation:
_______________________
Number of moles of chlorine
Calculation:
_______________________
Experimentally determined mole ratio of chlorine to copper
Calculation:
_______________________
Chemical formula of copper (II) chloride as determined experimentally
_______________________
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Additional Questions:
1. One way to express the fixed proportion of elements in the compound copper (II) chloride is the chemical
formula. Another way is the Mass Percent of each element in the compound. For example:
The “theoretical” Mass Percent of copper in the compound (as predicted by the true chemical formula, CuCl2),
would be:
Complete this calculation and determine the Theoretical Mass Percent of Copper in copper (II) chloride:
2. Now calculate the Mass Percent of Copper from your experimental data:
3. Compare your experimental mass percent of copper to the theoretical value by calculating the Percent Error.
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4. Thoroughly explain why your experimental mass percent is different from the theoretical value.
Think about what you observed: Was there any color remaining in the solution after your vacuum filtration? Was
any copper metal lost in transfer or filtration? Was the copper metal still moist when it was weighed?
Be specific as to how each possible source of error would have affected the outcome; that is, caused the mass
percentage of copper to be higher or lower than the theoretical value.
Finally, explain where you could improve your technique if you repeated the experiment.
5. Thought Question: How did the change in mass of the aluminum rod compare to the mass of copper produced
by the chemical reaction? Larger? Smaller? The same? And is this result consistent with your expectation?
Explain your reasoning.
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