Experiment 24
Qualitative Analysis I
GOAL:
In this experiment you will explore characteristic reactions of a number of ions. The reactions you
observe will be used to identify unknown solutions next week. We will use this study to develop logical
schemes for analyzing unknowns, and also to learn and review some descriptive chemistry of the
elements.
INTRODUCTION:
You will be studying seven cations (Na1+, Mg2+, Ni2+, Cr3+, Zn2+, Ag1+, Pb2+) and four anions (NO31-,Cl1-,
I1-, SO42-). These ions may be separated from each other and identified using simple chemical and
physical properties. Characteristic colors of aqueous ions as well as solid compounds can be very useful
in identification. In general, we expect main group cations to be colorless in solution and to form white
solids. Transition metal cations often have characteristic colors both in the aqueous solution and in solid
compounds. Some ions that generally form colorless compounds can give brilliant colors if mixed with
the right partner. For example, both Pb(NO3)2 and KI dissolve to give colorless solutions, but when
mixed, a bright yellow precipitate of PbI2 is produced.
I. Reactions of cations with NaOH and NH3
Most metal ions react with aqueous OH1-, hydroxide ion, to form gelatinous precipitates of the metal
hydroxides. For example, Al3+(aq) reacts to form Al(OH)3(s).
Al3+(aq) + 3 OH1-(aq) → Al(OH)3(s)
Eqn 1
Ag1+ is a rare exception, forming Ag2O(s) rather than AgOH. All other hydroxide precipitates you
encounter in this experiment are neutral hydroxides. Adding more OH1- causes some insoluble
hydroxides to re-dissolve. For example, Al(OH)3(s) reacts to form the complex ion Al(OH)4-(aq). In lab,
if you observe a precipitate re-dissolving upon addition of more OH1-, it is doing a reaction like Eqn 2.
Al(OH)3(s) + OH1-(aq) → Al(OH)41-(aq)
Eqn 2
In this experiment metal hydroxide precipitates that re-dissolve in excess OH1- form complex ions with
the general formula M(OH)4n-(aq).
Metals ions may react with NH3 to form either insoluble precipitates or soluble complex ions. If a
precipitate forms, it is the neutral hydroxide. For example, Al3+(aq) reacts with NH3(aq) to form
Al(OH)3(s). In lab, if you observe an ion reacting with NH3 to form a precipitate, it is doing a reaction
like that in Eqn 3.
Al3+(aq) + 3 NH3(aq) + 3 H2O(l) → Al(OH)3(s) + 3 NH41+(aq)
Eqn 3
If no precipitate forms or is only briefly present, a complex ion with NH3 has formed, for example,
Zn(NH3)42+(aq), Ni(NH3)42+(aq), and Ag(NH3)21+(aq). See Eqn 4.
Zn2+(aq) + 4 NH3(aq) → Zn(NH3)42+(aq)
Eqn 4
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II. Anion Precipitation Reactions
Of the cations we are studying, only Ag1+ and Pb2+ react with any of our four study anions to form
precipitates. Ag1+ should form precipitates with two of the anions; while Pb2+ should form precipitates
with three of them. The colors and solubilities of these precipitates will allow us to determine the
presence of these anions in unknown mixtures.
Any precipitates that form in these reactions are simple insoluble salts of the ions. Occasionally when a
large excess of one ion, especially the anion, is present, anomalous results may be noted. Ag2SO4, which
is moderately soluble, normally remains dissolved under our conditions, but may form a fine white
precipitate when concentrated. PbCl2 and PbI2, which are normally insoluble under our conditions, may
re-dissolve by forming complex ions with excess anion. The solubility of PbCl2 is also highly
temperature dependent. A PbCl2 precipitate will re-dissolve if it is heated in water. This fact may be
used to identify the precipitate since the other Pb2+ precipitates will not re-dissolve when heated.
III. Stability Sequences
Occasionally individual precipitation reactions are insufficient to allow identification of an unknown
mixture. A stability series such as those shown below may be helpful. A stability series is a ranking of
salts and ions for a given cation with the most stable species on the right. Addition of an appropriate
reagent to a compound can convert that compound into a compound further to the right in the series.
However, the equilibria will not allow you to go from right to left.
For Ag+:
Ag2O(s) < AgCl(s) < Ag(NH3)2+(aq) < AgI(s)
For Pb2+: PbCl2(s) < PbSO4(s) < PbI2(s) < Pb(OH)2(s) < Pb(OH)42-(aq) in xs OHFor Zn2+: Zn(OH)2(s) < Zn(NH3)42+(aq) < Zn(OH)42-(aq) in xs OHFor Ni2+: Ni(OH)2(s) < Ni(NH3)42+(aq) < Ni(OH)2(s) in xs OHFor Cr3+: Cr(OH)3(s) < Cr(OH)41-(aq) in xs OH1From the lead stability series, for example, we can see that addition of SO42-(aq) to solid PbCl2 (or
aqueous Pb2+) would cause PbSO4(s) to form. You could not, however, go the other direction. Adding
Cl-(aq) to PbSO4(s) will not form PbCl2(s). We could, however, add I-(aq) to Pb2+(aq), PbCl2(s) or
PbSO4(s) and get PbI2(s). Addition of excess OH1- to any of species in the lead series would cause the
formation of the complex ion Pb(OH)42-. Remember that any species shown as an ion in the stability
series above will be soluble in water. The neutral compounds will precipitate.
IV. Nitrate Ions
Nitrate ions, NO31-, will not precipitate, regardless of the ions mixed with them. The brown ring "spot
test" has been developed to identify this ion.
V. Sodium Ions
Sodium ions, Na1+, will not precipitate, regardless of the ions mixed with it. A flame test is the easiest
way to identify Na1+.
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WASTES AND HAZARDS:
All waste from qualitative analysis experiments should be placed in the combined qual waste beaker in
the hood. Acid and base solutions are corrosive. Use care when handling these solutions. Although the
ion solutions should be handled with respect, low concentrations minimize hazards that might be
associated with solid or concentrated forms of the same compounds. One of the bases used in this
experiment is aqueous ammonia. Look up its MSDS. Record in the Hazards section of your notebook
entries,
 appearance and odor
 concentration (found under Composition)
 potential health effects and the appropriate first aid for
o inhalation
o skin contact
o eye contact
 Given that our 6M NH3 is about 10%, do you think the hazards of our solutions are quite similar
to or quite different from what you found on the MSDS? Explain.
LABORATORY OBSERVATIONS AND DATA:
In the work that follows, keep clear, complete laboratory notes in your notebook. Be sure to fully
describe each reagent before mixing, and each final solution after mixing. Include colors,
precipitates, etc. Do not simply record “no change” upon mixing. Give the color and state of the
system, e.g. “green precipitate remains.” Some of the precipitates will be gelatinous. Any cloudiness is a
precipitate, even if the solid doesn’t settle to the bottom. Be sure to mix all reactions thoroughly.
PROCEDURE:
Stock solutions of the seven cations (Na1+, Mg2+, Ni2+, Cr3+, Zn2+, Ag1+, Pb2+) are supplied as 0.1 M
solutions of the nitrate salts. Stock solutions of NO31-, I1-, and SO42- are 0.1 M sodium salt solutions.
The Cl1- solution is 0.5 M NaCl and is slightly acidified. Other reagents available include 6M solutions of
NaOH, NH3, HCl, and HNO3.
Record your observations in the table provided.
I. Reactions of cations with NaOH and NH3
All Na1+ salts are soluble, including NaOH. Consequently, Na1+ will not form a precipitate with NaOH or
NH3. We will save a little time, and not perform these tests. You should still, however, record in your
table of results that Na1+ gives no precipitate with any of the reagents.
Obtain a well plate that has six columns of four wells each. Add 10 drops of Mg2+(aq) to each of two
wells in the first column. Similarly, prepare two wells each with 10 drops of the remaining five cations:
Ni2+, Cr3+, Zn2+, Ag1+, and Pb2+ . Add one drop of 6M NaOH to one of the wells for each cation. Tap the
well plate to mix. Note any reactions that occur. Record your observations. Any precipitates (even just
cloudy ones) are insoluble hydroxide salts.
Now add 10 more drops of 6M NaOH to the same wells. Tap the well plate to mix. Note any reactions
that occur. Any precipitates that re-dissolve are reacting with aqueous OH1- to form complex ions.
Formulas for these complex ions can be found in the introduction. If a precipitate does not re-dissolve, no
reaction has occurred. Record colors and the presence of any precipitates (even cloudiness).
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Finally, add 10 drops of NH3(aq) to each of the wells in the second row. Mix and observe any reactions.
When precipitates form, the metal cations have reacted with NH3 and H2O to form insoluble hydroxide
salts. If no precipitate forms or forms briefly and immediately re-dissolves, the metal cation has reacted
with NH3 to form a soluble complex ion with the ammonia. Formulas for these complex ions can be
found in the introduction. Record colors and the presence of any precipitates (even cloudiness).
Dump your used solutions in the designated waste container. Rinse the well plate thoroughly with
distilled water and shake it briefly to remove any excess water.
II. Anion Precipitation Reactions
Put 10 drops of Ag1+ into each of four wells, and 10 drops of Pb2+ into another four wells. Add 1 drop of
Cl1- solution to the first Ag1+ and Pb2+ wells. Note the colors of any precipitates. Add an additional 10
drops of Cl1- to each of these wells. Note any changes.
Repeat this procedure using the remaining three anions (I1-, SO42-, and NO31-) and the remaining unused
wells of Ag1+ and Pb2+. If you do not get the proper number of precipitates (two for Ag1+, and three for
Pb2+), then repeat.
III. A Stability Sequence
The following four reactions illustrate the stability sequence for Ag1+. Record observations for each of
the reactions, noting precipitates, colors, and the re-dissolution of precipitates.
Rxn 1: Combine 1 mL AgNO3(aq) and 1 mL NaOH(aq) in a test tube.
Rxn 2: Centrifuge and discard the liquid from Rxn 1. Add several droppersful of NaCl(aq) to the
precipitate from Rxn 1. Mix thoroughly.
Rxn 3: Centrifuge and discard the liquid from Rxn 2. Add NH3(aq) to the solid until it dissolves.
Rxn 4: Add NaI dropwise to the solution from Rxn 3 until a reaction is seen.
IV. Sodium Ions
For this experiment, test a solution of NaNO3 to see the characteristic test for Na1+.
Sodium flame test procedure:
Dip a cotton swab into the solution to be tested and then heat end of the swab in a Bunsen burner
flame. A very intense yellow flame is observed when Na1+ is present. Douse the swab in water to
extinguish the flame and dispose of it in the container provided.
When using a flame test to determine the presence of Na1+ in an unknown, you must always use a new
swab to avoid contamination. Despite even your best efforts, you may still see a slight yellow flame
when no Na1+ is present in the unknown. This happens because nearly everything is contaminated with
small amounts of Na1+.
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V. Nitrate Ions
Perform the brown ring test for nitrates on a sample of aqueous NaNO3. Record your observations. Be
careful! It is very easy to get a false negative result from the brown ring test.
Nitrate brown ring test.
Put 20 drops of the solution to be tested into a test tube, and then, carefully and slowly, add 20
drops of concentrated H2SO4. (The test tube will get hot, and may splatter if the acid is added too
quickly.) If necessary, cool this mixture by allowing cold tap water to run over the outside of the
test tube. In a second test tube, dissolve about 0.1 g FeSO4.7H2O in 1 mL of water. Now, hold the
test tube containing the H2SO4 solution at a 45o angle while you allow 5 drops of the FeSO4
solution to run slowly down the inside of the tube. Do NOT mix. The aqueous FeSO4 should
form a layer above the acid. If nitrate ion is present, a smoky brown ring will form at the solution
interface. This ring may take several minutes to form, and will disappear eventually. Look at the
interface between the two layers from several angles. If you were doing this test on a real
unknown sample, and there were a chance that the sample also contained I1-, this would interfere
with the test. To remove I1-, add several drops of saturated Ag2SO4 solution until no more AgI
precipitates.
RESULTS AND DISCUSSION:
At the end of this handout, a blank table is provided for you to record your observations for when the
cationic solutions are mixed with the following: OH1-, xs OH1-, NH3, and anion precipitation. (You may
not have mixed some of the combinations because we knew nothing would happen. Record this
information in your table anyway.) In the last column record your initial observations, noting especially
the color of each ion solution. Some additional notes are provided in this column as well. In each box, list
solution colors, formulas for any complex ions formed, formulas of any precipitates formed, and
precipitate colors. See the Introduction for help with formulas. DO NOT write something like “no
change.” Record the color and state of the chemicals present. The anion precipitation column will have
entries for only Ag1+ and Pb2+. The other cations will not precipitate with any of the anions. Make a copy
of this table to keep for yourself after you turn in your report. You will need this table next week to help
identify unknown solutions.
AN INTRODUCTION TO FLOWCHARTS:
The reactions you observed this week can be used to separate and identify unknowns. Flowcharts are a
useful way to track these reactions. You will learn much more about flowcharts next week when you
work with solutions of unknown ions. This week, you'll learn just enough to answer the question below
and get you ready for next week.
Let’s imagine we were testing a solution that could contain any of our seven cations. Would adding I1- to
this mixture let us decide what was present? Let’s answer this by drawing a flowchart to show the
possibilities. Start by listing all the ions that could be present. The vertical line with I1- next to it is our
way of showing that we add some I1-.
Na1+(aq), Mg2+(aq), Zn2+(aq), Ag1+(aq), Pb2+(aq) all colorless
Ni2+(aq) green,
Cr3+(aq) purple/blue
I1-
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Look at your results from this week and consider what would happen as each ion reacts with I1-. As you
saw in lab, Ag1+ and Pb2+ will precipitate as AgI(s) and PbI2(s), while all other metal ions remain
dissolved in solution. We represent this with the flow chart below.
Na1+(aq), Mg2+(aq), Zn2+(aq), Ag1+(aq), Pb2+(aq) all colorless
Ni2+(aq) green,
Cr3+(aq) purple/blue
I1-
AgI(s) pale yellow
PbI2(s) bright yellow
Na1+(aq), Mg2+(aq), Zn2+(aq) all colorless
Cr3+(aq) purple/blue, Ni2+(aq) green
Notice that the chart separates precipitates (to the left) from solutions (to the right) and that the formulas
of precipitates and complex ions are used. In the lab we can accomplish a similar separation. By
centrifuging and decanting, the precipitated AgI and PbI2 can be physically separated from the other ions
that are still in solution. Also notice that when precipitates form, we use a double drop line, rather than a
single line, in the flowchart.
Note that the flowchart above is being used to show all the possible things you might see. If you were
actually identifying an unknown in lab, you would include only the things you actually see. How would
we record our in-lab work? Let’s say you were given an unknown and told that it contained only one of
the seven cations. What is the first thing you would observe? The color! Your unknown is a clear and
colorless solution. You would write
Unknown #1, only 1 cation, colorless
Could be Na1+(aq), Mg2+(aq), Zn2+(aq), Ag1+(aq) or Pb2+(aq)
You have already eliminated Ni2+ and Cr3+ based upon color, so there is no reason to include them in your
flowchart. Now which test do you run? Several good choices are available, but let’s stick with the
addition of I1- since we used that above. In lab, you add some I1- to a bit of your unknown and still have a
clear, colorless solution with no precipitate. You should record this as
Unknown #1, only 1 cation, colorless
Could be Na1+(aq), Mg2+(aq), Zn2+(aq), Ag1+(aq) or Pb2+(aq)
I1No ppt., so don’t have
Pb2+ or Ag1+
Colorless, could be
Na1+ or Mg2+ or Zn2+
Since you didn’t get a precipitate, do you really need to record “no ppt”, or can you just write nothing
under the double drop line? It is safer to write “no ppt.” When you come back to this later, your
observations will be clear. If you had simply left it blank, you might think that you had forgotten to
record a precipitate that formed.
Was this addition of I1- a useful way to identify your unknown? In this case, we narrowed the
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possibilities but will need to do additional steps to choose from among the remaining possibilities. We
will need to do a second test of the original solution, or we might add tests to the separated precipitate and
solution. We would then add these steps to the flowchart. Next week you will do such multi-step
flowcharts.
Does that mean that adding I1- isn’t a good way to start our unknown identification? It’s certainly not the
only step we might have chosen, but a different unknown might make things seem quite different. Let’s
say you get a second unknown, it is also clear and colorless and contains only one of our seven cations.
Again you choose to add I1-, but this time you get a pale cream precipitate. You should record this as
Unknown #2, only 1 cation, colorless
Could be Na1+(aq), Mg2+(aq), Zn2+(aq), Ag1+(aq) or Pb2+(aq)
I1-
pale cream ppt must be AgI(s)
so unknown contained Ag1+
colorless soln
In one step we have identified this unknown. We don’t need to do any additional tests or try to list
possibilities that we didn’t actually see. Since you now know that the unknown contained Ag1+, you no
longer list any other possibilities.
QUESTIONS:
1.
Give balanced chemical equations. Check your lab results and the Introduction to decide what
happens in each case. Some may be N.R., no reaction.
a.
Cr3+(aq) + xs OH1-(aq)
b.
Ni2+(aq) + NH3(aq)
c.
Na1+(aq) + NH3(aq)
d.
Cr3+(aq) + NH3(aq)
e.
Pb2+(aq) + xs OH1-(aq)
f.
Pb2+(aq) + Cl1-(aq)
g.
Ni2+(aq) + Cl1-(aq)
2.
Draw a flow chart, similar to the first one in the Introduction to Flowcharts section, to show what
would happen if you did each of the tests below. Start each flowchart with the specified ions in
the mixture listed at top. Include the color of each of those ions in aqueous solution. After the
specified reactant has been added, you will have both a precipitate and a supernatant solution.
Give correct formulas and colors for the precipitates. Give correct formulas and colors for the
ions in the supernatant solution. Be careful; you may get some complex ion formation.
a.
excess OH1- added to a mixture of all seven metal ions
b.
NH3 added to a mixture of aqueous Na1+, Mg2+, Ni2+, Cr3+, and Zn2+
c.
I1- added to a mixture of Ag1+, Zn2+, and Ni2+
3.
Carefully re-read the Introduction Section III Stability Sequences. Write balanced equations.
a.
PbCl2(s) + excess OH1-(aq) [see the stability sequence for Pb2+]
b.
AgCl(s) + excess OH1-(aq) [see the stability sequence for Ag1+]
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4.
While again watching those stability sequences, draw flow charts showing what would happen if
you added
a.
excess OH1- to a mixture of AgCl(s) and PbCl2(s) [see your answers to #3]
b.
NH3 to a mixture of AgCl(s) and PbCl2(s) [see stability sequences!]
5.
You are given a test tube containing an aqueous solution of a nitrate salt of one of our cations
(Na1+, Mg2+, Ni2+, Cr3+, Zn2+, Ag1+, or Pb2+) but you don’t know which it is. You must use the
reactions of this experiment to identify the single cation in the test tube. If you could do only one
test, which would be most helpful in identifying the unknown or at least narrowing the
possibilities? Consider these options: add 1 drop OH1-, add excess OH1-, add NH3, add Cl1-, add
I1-, do a flame test, or do a brown ring test. Write several well-organized paragraphs in which you
consider these possibilities, discuss the merits of each option, and then finally draw a conclusion
about the best choice or choices.
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Characteristic Reactions for Identifying Unknown Ions
Test Ion
OH-
XS OH-
NH3
Anions
(Cl-, I-, SO42-)
Miscellaneous
Na+
Yellow flame
Mg2+
Slow pption
Ni2+
Cr3+
Zn2+
Ag+
Stains black;
Ag2SO4 may
ppt at high
conc
Pb2+
PbCl2 will
dissolve in
HOT water
NO3-
All nitrates are
soluble; Brown
ring test
Stability sequences
Some precipitates are less soluble than others. You can only move left to right.
For Ag+:
For Pb2+:
For Zn2+:
For Ni2+:
For Cr3+:
Ag2O(s) < AgCl(s) < Ag(NH3)2+(aq) < AgI(s)
PbCl2(s) < PbSO4(s) < PbI2(s) < Pb(OH)2(s) < Pb(OH)42-(aq) in XS OHZn(OH)2(s) < Zn(NH3)42+(aq) < Zn(OH)42-(aq) in XS OHNi(OH)2(s) < Ni(NH3)42+(aq) < Ni(OH)2(s) in XS OHCr(OH)3(s) < Cr(OH)4-(aq) in XS OH9