6. A Six-Bottle Study of Ionic Compounds
What you will accomplish in this experiment
In last week’s experiment, you observed a chemical change: a chemical reaction that resulted in a change in the
composition, and thus the identity, of the substances being observed.
That chemical reaction was the process of rearranging, removing, replacing or adding atoms (and/or electrons)
so that new substances were formed.
This week’s experiment will provide you with another opportunity to practice your skills in observing chemical
reactions. This time, however, you’ll need to be able to draw conclusions from your observations – you’ll need
to make a decision as to whether or not a chemical change has truly occurred (whether a new substance has
been formed).
In each case where a change in composition has occurred and a new substance has been produced, you’ll need
to write chemical formulas for those new substances, and then write a balanced chemical equation to fully
describe the chemical change.
Finally, you’ll use your observations of those chemical changes to deduce the identity of unknown substances
by carefully studying their chemical behavior.
Concepts you need to know to be prepared
It’s much easier to keep track of the chemical changes that matter is undergoing if a reaction is expressed in the
form of a chemical equation. Rather than writing out the names of the elements and compounds that are
participating in the reaction, the chemical formulas of these substances are used.
An arrow serves as the verb in this chemical sentence, indicating the transformation of the reacting substances
(the “reactants”) into the “products.”
For example, let’s write the chemical equation for the reaction between copper (II) chloride and aluminum
metal that you performed in last week’s experiment.
The English sentence that describes this reaction might read:
“Aqueous copper (II) chloride reacts with aluminum wire to form copper metal and aqueous aluminum chloride. ”
To express this reaction as a chemical equation, we would first need to write the chemical formulas of the
participating elements and compounds.
Reactants
Products
aqueous copper (II) chloride: CuCl2 (aq)
copper metal: Cu (s)
aluminum wire: Al (s)
aqueous aluminum chloride: AlCl3 (aq)
Notice that the subscripted symbol in parentheses included with each chemical formula indicates the physical
state in which that substance exists. Two of the substances in this chemical reaction are in the solid state (s),
and two are dissolved in water, or “aqueous,” (aq). Other common symbols you’ll encounter are (g) to indicate
a gas, and (l) to indicate a liquid.
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Our chemical equation describing the reaction between copper (II) chloride and aluminum wire can now be
written as:
CuCl2 (aq) + Al (s) → Cu (s) + AlCl3 (aq)
But clearly something is not right about the chemical equation as it’s written above. It doesn’t indicate that
matter is conserved during the reaction: the amount of matter that’s present at the start of the reaction is the
same as the amount of matter at the conclusion of the reaction, because no matter is created or destroyed.
The chemical formula for copper (II) chloride on the reactant side of the equation indicates two chloride ions;
yet the chemical formula for aluminum chloride on the product side of the equation shows three chloride ions.
The proportion of positive and negative ions in each of these compounds is fixed, so the chemical formulas of
the compounds CANNOT be changed so as to indicate the conservation of matter. But what CAN be changed
are the amounts of each participant as expressed by the chemical equation.
We can write numbers in front of the chemical formulas to indicate more than one of a reactant or product
“formula unit.”
A “formula unit” is simply the
small grouping of atoms or ions
expressed in the chemical formula
of an ionic or molecular substance.
For example, a formula unit of the
ionic compound sodium chloride
(NaCl) would consist of one Na+
ion and one Cl- ion.
And a formula unit of the
molecular compound methane
(CH4) would consist of one C atom
and four H atoms.
The numbers written in front of the chemical formulas in a chemical equation are called “coefficients,” and
adding them to the chemical equation is called “balancing” the equation.
The chemical equation for the reaction of copper (II) chloride with aluminum could be balanced by adding the
coefficients shown below:
3 CuCl2 (aq) + 2 Al (s) → 3 Cu (s) + 2 AlCl3 (aq)
Now the chemical equation shows that:
“3 formula units of aqueous copper (II) chloride react with 2 atoms of aluminum metal to produce 3 atoms of
copper metal and 2 formula units of aqueous aluminum chloride.”
3 (formula units) CuCl2 (aq) + 2 (atoms) Al (s) → 3 (atoms) Cu (s) + 2 (formula units) AlCl3 (aq)
If you count the number of each type of atom on the reactant side of the equation, you’ll find that it exactly
matches the number of each type of atom on the product side of the equation. The equation is “balanced.”
A balanced chemical equation MUST truly represent a chemical change: the transformation of one or more
substances into new substances, which have different chemical and physical properties. Evidence of this
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chemical change may be based on observations such as:
• The release of a gas, such as carbon dioxide gas. This can occur when an acid is added to an ionic
compound that contains the polyatomic anion carbonate (CO32-) or bicarbonate (HCO32-). The acid donates
H+ ions to the carbonate to create carbonic acid (H2CO3), which decomposes to water (H2O) and carbon
dioxide gas (CO2).
• The formation of a solid (called a “precipitate”) when aqueous solutions of ionic compounds are mixed, and
a product is created that is not water-soluble.
Chemical changes, whether they involve the release of gas or the formation of a precipitate, generally follow
one of a few simple patterns.
For example, the chemical change that you observed between copper (II) chloride and aluminum followed the
pattern of a “single-replacement” reaction, where aluminum replaced copper in an ionic compound.
AB + C → A + CB
3 CuCl2 (aq) + 2 Al (s) → 3 Cu (s) + 2 AlCl3 (aq)
A “double-replacement” reaction involves two compounds undergoing a “change of partners.” This often
means that two compounds react by exchanging ions, forming two new compounds:
AB + CD → AD + BC
One example is the reaction of an acid (hydrochloric acid, HCl) with a base (sodium hydroxide, NaOH) to
produce water and a salt, sodium chloride (NaCl).
HCl (aq) + NaOH (aq) → H2O (l) + NaCl (aq)
In the case of this particular example, it would be very difficult for you to discern that a chemical change had
actually taken place, as there would be no gas bubbles or solid precipitate for you to observe, indicating that
new substances had indeed been formed.
For this reason, all of the chemical changes that you’ll observe in this week’s experiment will involve either the
release of a gas or the formation of a solid precipitate. You’ll need to describe each one of these chemical
changes by writing a balanced chemical equation. The fact that all of these chemical changes will be doublereplacement reactions will simplify your task. And in the few cases where your observations indicate that NO
new substance is formed from a particular combination of reactants, you’ll simply write “NO REACTION” on
the products side of the arrow.
Here is the scenario for this week’s experiment:
The chemistry stockroom manager has brought six unlabeled bottles into the lab. These bottles contain aqueous
solutions, and they had once been properly labeled, but as the solutions are quite old, the labels have fallen off.
Some of the solutions are salts of heavy metal ions, so those will require proper disposal as hazardous waste.
It will be your job to determine which solution is in each bottle. From the labels, you can deduce which
compounds are possible, but you have to match each of the labels to the appropriate bottle.
There are three phases to this lab activity:
1. You’ll need to experiment with KNOWN samples of the six possible solutions, in order to determine how
each substance behaves when it is mixed with each of the others. the Observations table in the Report
Sheet will help you organize your observations for this step.
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2. You’ll follow the same procedure with the six UNKNOWN solutions, and deduce which of the unidentified
bottles contains which substance.
3. You’ll write balanced chemical equations describing each of the chemical changes that you observed when
the different pairs of solutions were mixed.
The six solutions are:
•
•
•
•
•
•
Lead nitrate, Pb(NO3)2
Silver nitrate, AgNO3
Sodium carbonate, Na2CO3
Potassium iodide, KI
Hydrochloric acid, HCl
Sodium hydroxide, NaOH
You’ll note that three of these compounds contain “polyatomic ions:” covalently-bonded groups of atoms that
have a net charge.
• The nitrate anion has the formula: NO3• The hydroxide anion has the formula" OH• And the carbonate anion has the formula: CO32You may recall that all of the metal cations you worked with in Lab #3 (the paper chromatography separation
experiment) were given to you as “nitrates:” aqueous solutions of the nitrate “salt” of each ion.
• The solubility table below will enable you to make predictions as to which combinations of the six
aqueous solutions will lead to a chemical change that produces a solid precipitate – a product that is not
water-soluble. (And note that the table shows that almost all nitrate salts are water-soluble, which
made it possible for you to “spot” aqueous solutions of those metal ions on your chromatography paper
in Lab #3.)
Procedure that you will follow
Part A: Observations of Known Solutions
1. Obtain a set of dropper bottles labeled with the names (and/or chemical formulas) of the six known
reference solutions.
2. In separate wells in your newly-CLEANED spot plate, mix five drops of any known solution with five drops
of another known solution, and record your observations in the appropriate box in the Part A Observations
Table on your Report Sheet.
3. Repeat this process until you’ve made all possible combinations of the six solutions, and recorded all
observations in your table. It’s essential that you are quite certain of what you’ll observe when you mix any
one of the known reference solutions with any other of the solutions.
4. When emptying your spot plate into your waste beaker, make sure you rinse it completely with distilled
water from your wash bottle before conducting more tests.
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Part B: Observations of Unlabeled Solutions
5. Obtain a set of dropper bottles labeled as the six unknown solutions. You already know that these are the
same six solutions you worked with in Part A, but you need to determine which is which.
6. In separate wells in your newly-CLEANED spot plate, mix five drops of any unknown solution with five
drops of another unknown solution, and record your observations in the appropriate box in the Part B
Observations Table on your Report Sheet.
7. Repeat this process until you’ve made all possible combinations of the six solutions, and recorded all
observations in your table.
8. In the last row of the Part B Table, write the correct name and chemical formula for each of the solutions
you identify. Circle the chemical formula for those solutions which contain heavy metal ions, and so will
require special hazardous waste treatment.
9. When emptying your spot plate into your waste beaker, make sure you rinse it completely with distilled
water from your wash bottle before conducting more tests.
Part C: Writing a Balanced Chemical Equation for each Reaction Observed
10. For each combination of solutions in which you observed that a chemical reaction took place, you’ll need to
write out the chemical equation for that reaction by:
•
Determining the chemical formulas for the products of the reaction.
•
Balancing the chemical equation.
•
Indicating the physical state of all reactants and products in the reaction.
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 6: A Six-Bottle Study of Ionic Compounds
Student ____________________________ Lab Partner_________________________ Date Lab Performed__________
Section #_________ Lab Instructor_______________________________________ Date Report Received ___________
Lab Notebook: Data and Observations
Observations Table for Part A:
Sodium
hydroxide
Hydrochloric
acid
Potassium
iodide
Sodium
carbonate
Silver
nitrate
Lead
nitrate
NaOH
HCl
KI
Na2CO3
AgNO3
Pb(NO3)2
Lead
nitrate
Pb(NO3)2
Silver
nitrate
AgNO3
Sdium
carbonate
Na2CO3
Potassium
iodide
KI
Hydrochloric
acid
HCl
Sodium
hydroxide
NaOH
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Observations Table for Part B:
Unknown 1
Unknown 2
Unknown 3
Unknown 4
Unknown 5
Unknown 6
Unknown 1 =
Unknown 2 =
Unknown 3 =
Unknown 4 =
Unknown 5 =
Unknown 6 =
Unknown 6
Unknown 5
Unknown 4
Unknown 3
Unknown 2
Unknown 1
Identity of
Unknown
(name &
chemical
formula)
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Formal Report: Results and Conclusions
Part C: Balanced Chemical Equations for Solution Combinations where a Reaction Occurred
For each combination of solutions in which you observed that a chemical reaction took place, write out the
chemical equation for that reaction by:
• Determining the chemical formulas for the products of the reaction.
• Balancing the chemical equation.
• Indicating the physical state of all reactants and products in the reaction.
Additional Questions:
Translate each of the following descriptions of a chemical reaction into a balanced chemical equation. Use
symbols to indicate the physical state of each participant in the reaction as described in the sentence.
a) Copper metal reacts with oxygen gas to form the solid copper (I) oxide.
b) The salt iron (III) oxide reacts with hydrogen gas to form iron metal and liquid water.
c) Aluminum metal reacts with liquid bromine to form crystals of aluminum bromide.
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