The N-Bottle
Introduction
The purpose of this experiment is to use your knowledge about acid-base and precipitation
reactions that occur in aqueous solutions to identify the ionic substances present in a set of solutions.
These solutions will be labeled A through G, and you will be given the list of cations and anions that are
present. All the solutions that you will be given are clear; hence you can already conclude that only
soluble cation/anion combinations have been selected.
As a simple example, suppose that you are given three clear solutions X, Y and Z and that the
solutions may contain the actions Na+, K+ and Ag+ and the anions CO32-, NO3- and Br-. Since the
solutions are clear, you may conclude that one of the solutions does not contain AgBr since AgBr is
insoluble in water and would form a precipitate. Likewise, one of the solutions does not contain Ag2CO3
since this compound is also insoluble. Thus, one solution must be AgNO3 and the others must be
Na2CO3 and KBr or K2CO3 and NaBr. You could identify which is which by systematically mixing ~0.5
mL pairs of solutions and observing what happens. Suppose you mix X with Y and observe the
formation of a pale yellow precipitate. Since Na+ and K+ compounds are soluble, either X or Y must
contain AgNO3, and the reaction which occurs must be one of the following:
2 Ag+ (aq) + 2 NO3- (aq) + 2 Na+ (aq) + CO32- (aq)  Ag2CO3 (s) + 2 Na+ (aq) + 2 NO3- (aq) (4-1)
Ag+ (aq) + NO3- (aq) + K+ (aq) + Br- (aq)  AgBr (s) + K+ (aq) + NO3- (aq)
(4-2)
2 Ag+ (aq) + 2 NO3- (aq) + 2 K+ (aq) + CO32- (aq)  Ag2CO3 (s) + 2 K+ (aq) 2 NO3- (aq) (4-3)
Ag+ (aq) + NO3- (aq) + Na+ (aq) + Br- (aq)  AgBr (s) + Na+ (aq) + NO3- (aq) (4-4)
Suppose you now mix X with Z and also observe a yellow precipitate. Then, one of the above
four reactions most have occurred; this identifies X as AgNO3 (aq). The identity of Y and Z is still
unclear. If you mix Y and Z, no reaction occurs since all possible products are soluble. You must now
try to distinguish what is in these solutions. Obviously, a precipitation reaction will not work. If some
concentrated HCl is available, you could add several drops to samples Y and Z. Suppose that no reaction
occurs with Y, but that a gas is released from Z. the reaction is evidently
CO32- (aq) + 2 H3O+ (aq)  H2CO3 (aq)  CO2 (g) + H2O (l)
(4-5)
Thus, Z must contain carbonate, and Y must contain bromide. We have no evidence at this point
about Na+ and K+. A flame test could be used to differentiate between these ions. Na gives a bright
yellow flame test. However, since glass vessels contain sodium silicates, almost every aqueous solution
gives a positive test for Na. Potassium gives a purple flame that is often masked by the yellow of Na.
Viewing the test through blue cobalt glass can filter out this yellow color. Suppose solution Z shows a
positive flame test for K. Z must contain K2CO3 and, by elimination, Y is NaBr.
Among the pieces of information you need for this experiment is a set of solubility rules. Several
statements of the rules are possible. However the set of rules given below is particularly useful because
the rules are listed in hierarchical order; that is, rule 1 takes precedence over rule 2 and so on.
1.
2.
3.
4.
5.
6.
All Li+, Na+, K+, and NH4+ salts are soluble.
NO3-, C2H3O2- and ClO4- salts are soluble.
Ag+, Pb2+ and Hg22+ salts are insoluble.
Cl-, Br- and I- salts are soluble.
CO32-, S2-, O2-, PO43- and OH- salts are insoluble.
SO42- salts are soluble except for those of Ca2+, Sr2+ and Ba2+.
Thus, for example, using these rules we can predict that NaOH is soluble (rule 1 takes precedence
over rule 5), that Pb(C2H3O2)2 is soluble (rule 2 takes precedence over rule 3) and that AgBr is insoluble
(rule 3 takes precedence over rule 4).
Technique
For mixing solutions, you can use small (micro) test tubes and plastic disposable Pasteur pipets with
graduated volume markings on the barrel. The precise amounts of solutions you mix are not crucial; a
good idea would be to use about 0.5 mL of each. You can hold the tube flick it with your finger to
achieve mixing. It is never a good idea to put your finger over the mouth of the test tube and shake
it! After mixing solutions and observing any reaction, you may discard the contents of the test tube into
the designated waste container.
You may work in pairs and discuss the experiments and interpretations with our partner. However,
each person must submit an independent lab report. Alternatively, you may work alone if you so choose.
Equipment Needed
burner
small graduated cylinder
micro test tubes
glass stirring rod
red and blue litmus paper
nichrome wire
disposable plastic pipets (7/pair of students)
test tubes (6”)
test tube holders for micro and 6” test tubes
Chemicals Needed
Conc. HCl
Unknown bottles labeled A-G (These unknowns contain the following cations and anions: Ag+,
Ba2+, Co2+, H3O+, Li+, Na+, Sr2+ and Cl-, CO32-, OH-, NO3-, SO42-. Each solution contains a
different cation, but more than one solution can contain the same anion. Recall that a salt
solution must contain both a cation and an anion.)
Procedure
Obtain about ten mL of each solution A-G in separate small test tubes. Label the test tubes
carefully. If you pour out more solution than this, do not return any of it to the reagent bottle; just
discard what is left over into the inorganic waste container at the end of your work. Place a clean plastic
pipet in each test tube. Test the acidity or basicity of each solution by dipping a clean glass stirring rod
into the solution and placing a drop of the solution onto strips of red and blue litmus paper. An acidic
solution (H3O+ present) will turn blue litmus to red and a basic solution (OH- or CO32- present) will turn
red litmus paper to blue.
Flame Tests
Obtain a piece of nichrome wire attached to a glass stirring rod and 2-3 mL of concentrated HCl in
your small graduated cylinder. The wire should have a loop on the end; if it does not, wrap the end
around a pen or pencil. You will want to be sure that the wire itself is free or any material that could
color the flame. You can clean the wire by dipping the loop into the concentrated HCl and then holding
it in the hottest portion of the burner flame. It will be necessary to clean the wire each time before you
test a new solution. (Why?) The purpose of the HCl is to convert any metal ions into their chloride salts,
which are easily volatilized.
For each sample, dip the clean loop of the wire into the solution and thrust it into the flame. Note:
the results will be more satisfactory if the burner flame is ‘cool’; this can be achieved by turning the
burner barrel clockwise until the blue inner cone disappears. Record the results, noting not only the color
of the flame but also the intensity and duration. The colors of flames are as follows; Li—red; Ba—
green; Na—yellow; Sr—red. The other ions give no characteristic flame test color. As mentioned above,
every solution that has been stored in a glass vessel will give a positive Na test due to the leaching of
sodium from the glass. Thus, a positive flame test due to a dissolved unknown Na salt should exhibit a
persistent and very intense yellow color. As a corollary, the colors due to other cations are visible briefly
after the yellow of the leached-in Na disappears. Because you have other criteria besides flame tests for
determining the ions present, no “known” solutions will be provided for comparison. Make some
tentative hypotheses about which solution contains which cations based on above results.
Mixing solutions
Now, begin to mix pairs of the solutions together systematically and observe the results. In order to
help you interpret the results, you should prepare a chart detailing what will happen when various
solutions are mixed before you begin testing any of the solutions. The idea, of course, is to identify both
the cation and the anion present in each solution by giving the proper formula for the compound made
from these ions—for example, MgCl2. To illustrate, the chart for the example mentioned in the
introduction would look as follows:
CO32-
Br-
NO3-
Na+
no ppt.
no ppt.
no ppt.
K+
no ppt
no ppt.
no ppt.
Ag+
AgCO3 (s)
AgBr (s)
no ppt.
Write your char, which will be substantially larger than this example, in your notebook. Turn in the
copy of this chart to your professor, who will grade it as a portion of your lab score for this experiment
and return it to you before the end of the lab period.
Be sure you have reached your conclusions about which ions are present in which solutions before
you leave the lab. Extra working time will not be allowed for this experiment, so be sure that you have
done all the tests needed during your own lab period.
Results
Construct data tables in which you present your observations from the litmus, flame and
precipitation tests.
Discussion
Present a table in which you identify the compound (cation and anion) present in each solution, A-G.
For example,
A
B
NaCl (aq)
Mg(NO3)2 (aq),
etc.
For each solution, explain on the basis of your test results how you reached the conclusion you did.
For each reaction that occurred (either precipitation or acid/base) write out the net ionic equation.
For example, for Equation (4.1), the net ionic equation is:
2 Ag+ (aq) + CO32- (aq)  Ag2CO3 (s)
(4-6)