Inorganic Qualitative Analysis1,2,3
Authors: B. K. Kramer* and J. M. McCormick
Last Update: January 15, 2013
Qualitative analysis is the identification a
sample's component(s). Unlike a quantitative analysis, we are not
concerned with the amount of a substance present in a sample but only with its
identity. In this exercise we will focus on identifying the cations and anions
that make up ionic compounds, both solid and in solution. Ideally there would be
chemical tests that could be used to identify individual ions without
interference by any other ions. Unfortunately, there are often complications.
For example, the formation of a yellow precipitate upon addition of aqueous S2-
confirms the presence of Cd2+ in a solution. The color of this
compound, however, will be hidden if any Pb2+ or Cu2+ are
present in solution since they will form a black precipitate with added S2-.
In order to test for cadmium, then, any interfering ions must first be removed.
This will be the case for most ions in a mixture: before their identities can be
confirmed, they must be isolated from the remaining solution.
The separation of ions in solution can be accomplished by
the addition of a precipitating agent that will selectively react with an ion in
the solution and not with others that may be present. The solid that is produced
can then be removed from the liquid by centrifugation and decanting. Because
many ions may behave similarly, separation of individual ions from a complex
mixture is not usually possible. Instead, a group of ions with similar
reactivity may be separated by precipitation from a larger mixture. After they
are isolated in a solid, they must be further separated and reacted to confirm
each one’s identity.
There are several types of reactions that can be used to
confirm the presence of ions in solution. The most common are precipitation
and complexation. In a precipitation reaction, an ion in solution reacts
with an added reagent to form a solid. Whether a solid will form from a given
reaction can be predicted by the solubility product constant (Ksp)
of the solid under the given conditions. Solubility product constants are the
equilibrium constants for the dissolution of an "insoluble" ionic solid in
water. A low Ksp implies that the compound does not dissolve
to an appreciable degree in water. If the two ions are mixed in solution, a
precipitate will tend to form. If steps have been taken to remove ions that form
competing precipitates, the presence of a properly colored solid can be used to
confirm the presence of a given ion. If several different precipitates remain,
the conditions of the solution can be manipulated to selectively redissolve one
or more of the solids. When the equilibria involved are well understood,
selective precipitation can be a powerful tool in the identification of unknown
Complexation can also be used to determine the presence of
an ion in solution. In a complexation reaction, a cation (typically a metal)
forms covalent bonds with one or more ligands (More
Info). A ligand is a neutral or negatively charged species that donates
electrons to the positively charged metal to form a coordinate covalent,
or dative, bond. These complexes may either be neutral or charged,
depending on the charge on the metal and on the ligand. When a complex forms it
may not precipitate (charged complexes are often quite soluble in water, for
example), and the formation of a complex is one way in which an insoluble metal
ion can be forced to dissolve. Similarly, complex formation can also be used to
separate a mixture of ions by keeping one or more in solution while others
precipitate. If the complex formed between a metal ion and a specific ligand has
a distinct color, complex formation can be used to demonstrate the presence of a
specific metal ion by simply adding the ligand to the solution. They are useful
in confirming the presence of a single ion after separation has been achieved.
The tendency to form a complex can be determined by the formation constant (Kf)
of the reaction. Formation constants are defined as the equilibrium constant for
the reaction of the metal ion with the ligand(s) to form a complex. A large Kf
implies a strong tendency for complex formation.
Qualitative analysis schemes have been performed by
introductory chemistry students for many years. They are used to help students
understand reactivity and to develop problem solving skills. There are several
different approaches to these experiments. In one case, students are given a
step-by-step procedure, often in the form of a flow chart, which they can use to
isolate and identify unknown ions in solution. In another, students first
analyze known solutions to determine how different anions will behave when
reacted with various reagents. They then compile these results into their own
flow chart that they apply to their unknowns. The experiments can be carried out
on solutions containing mixtures of cations (same anion), mixtures of anions
(same cation) or on salt mixtures.
The creation of the flow chart from scratch is very
valuable but is also a very time consuming process. Your experiment will involve
a modification of the flow chart procedure. The reactivity of the different ions
with precipitating agents can be predicted based on the Ksp of
the salt formed if the two were to react. You will use the Ksp’s
of several salts to determine the best way to cause their separation. You will
then prepare a flowchart to separate and identify these components. You will
test this flowchart in the lab. Ideally you would create a flowchart for the
separation of the entire mixture. Because of time constraints, the flowcharts
for the remaining species you need to separate will be given to you.
The overall experiment has three parts. In the first part
you will analyze known mixtures of cations using your predetermined procedure
and procedures that are given to you. This part is expected to take one to one
and a half weeks. In the second part you will apply similar procedures to known
mixtures of anions in solution. This part should take less than a full
laboratory period. In the final week you will be given three unknown mixtures: a
solution of three unknown cations, a solution of three unknown anions and a
solid salt consisting of a single cation and a single anion. You will need to
use the procedures you learned the previous weeks to identify the components of
The separation flow chart for the cations and anions
encountered in this exercise are shown in Fig. 1 and Fig. 2, respectively.
These flow charts show the steps required to separate and identify the cations
and anions that you may find in your known and unknown mixtures. These charts
have been prepared based on theoretical information about the ions and
experimental observations. The flow charts can help you understand the order in
which separations must take place in order to isolate ions that may behave
Figure 1. An example cation separation scheme for
this exercise. See Fig. 3 for an explanation of the symbols used. You will
need to prepare a flow chart for your instructor's approval for the chloride
group (Ag+, Hg22+, Pb2+) before
starting this exercise. Click
to obtain these flow charts in PDF format suitable for printing.
Figure 2. An example anion separation
scheme. See Fig. 3 for an explanation of the symbols used. Click
obtain this file in PDF format for printing.
Throughout the flow charts, reagent additions and other
procedures are indicated along the connecting lines; these are explained in more
detail below. The formula for each species, along with any identifying physical
characteristics (such as color), is given in the box. The symbols and formalism
used in the flow charts are given in Fig. 3.
Figure 3. Key for the ion separation flow charts
given in Fig. 1 and Fig. 2.
Since this analysis is qualitative and not quantitative,
it is not necessary that exact amounts of reagent be added at each step, but it
may be useful to know that there are approximately 20 drops in 1 mL. Each
procedure should be performed on approximately 0.5 mL of a fresh sample of
solution. IMPORTANT! Since some ions removed early in a given procedure
may mask those determined later, it is essential that the entire chart is
followed in order.
You will be using microcentrifuge tubes throughout this
procedure. The tubes can hold either 1.5 or 2.0 mL (listed on the tube). If
the volume of your solution exceeds that of the tube, separate the solution into
two tubes and treat each one according to the flowchart.
Throughout this series of experiments, you will be
expected to follow the directions that are presented in the flow charts included
in this lab. The reactions shown in the charts are described in the accompanying
text. There are several procedural steps that are indicated in the flow chart
that are described here.
- Precipitation After the addition of a
precipitating agent, it is important to mix the solution thoroughly by
shaking or by stirring with a clean glass stir rod. Be sure not to
add more of a precipitating agent than indicated in the procedure as this
may cause undesired side reactions.
- Separation After the addition of a
precipitating agent to a solution, a solid (precipitate) and liquid
(supernatant) will result. These often must be isolated and treated
separately. The most common method for separating a supernatant and
precipitate is to centrifuge the mixture to cause the solid to compact at
the bottom of the tube. The procedure for centrifuging will be demonstrated
by your instructor. It is essential that the centrifuge is balanced with a
tube containing the same volume of liquid as your sample to prevent it from
"walking" off the table!!
After centrifuging, the supernatant can be decanted by simply pouring from
one test tube into another or by careful removal with a clean pipet.
- Washing When a supernatant is removed from a
solid, it is almost certain that some of the liquid has been left behind.
This liquid can be removed by the addition of a clean solvent (usually cold
water, but indicated in the procedure if not) which is thoroughly mixed with
the precipitate. After centrifuging and decanting, the solid is now ready
for further reaction as dictated in the procedure. IMPORTANT!
Improper separation and washing of precipitates is the most common source
of error in this exercise. So, be sure that you learn how to do this
properly using the known solutions.
- pH Adjustment Often it is necessary to adjust
the pH of a solution until it is just alkaline or just acidic.
This is usually accomplished by the dropwise addition of a strong acid or
base. In order to make sure that the solution does not become too acidic or
basic, the pH of the solution must be monitored. You will use universal
indicator paper to determine the pH of your solutions. The proper method for
finding the pH of a solution involves stirring the solution with a clean
glass rod and then touching the tip of the rod to a piece of indicator
paper. Do NOT place the indicator paper directly in your solution!
You should check the pH of the solution after each addition of a drop of
acid or base. If you are to acidify, stop the addition as soon as the paper
registers a pH of just less than 7. If the solution is to be made basic, add
base until the paper registers just more than 7. If you are to neutralize
the solution, add the appropriate acid or base until the paper reads very
close to 7.
- Heating Solutions There are two different
heating methods that will be used in this exercise. If a solution needs to
maintain a near boiling temperature for a few minutes, the tube should be
placed in a boiling water bath. Microcentrifuge tubes can be heated by
placing test tubes full of water in a water bath. After allowing the water
to boil, place the microcentrifuge tube in the top of the test tube, making
sure that the liquid contained is fully immersed in hot water. Make sure
that the test tube is anchored so that it does not spill into the bath.
- If the solution needs to be heated directly, the
solution should be transferred to a glass test tube which
should be held in a test tube clamp facing away from you and your fellow
students and passed back and forth through the flame of a Bunsen burner.
Be careful not to let the solution bump and jump out of the test tube by
keeping the flame near the surface of the solution rather than at the
bottom. Stirring the solution with a glass rod may also help. The procedures
will indicate which heating method is necessary for each step. DO NOT
place a plastic microcentrifuge tube in or near a flame.
- Flame Tests In many cases, the color emitted
when a cation is heated directly in a flame can help identify the element.
If a procedure calls for a flame test, follow the directions below.
- 1) Clean a wire loop by first dipping it in 6 M
HCl and then heating it long enough to drive off any contaminants from
- 2) Place the loop in the solution or solid to be
tested, making sure a drop of liquid or crystal remains in the loop.
- 3) Place the wire in the flame and observe the
color emitted. If instructed, view the flame through a piece of cobalt
Throughout the experiment, you will use a flow chart (Fig.
1) to help you separate and identify the cations in your system. The first part
of the chart has been left blank. Using the information in the paragraphs below,
propose the steps to fill in the flow chart to isolate the chloride group (Ag+,
Hg22+ and Pb2+) from a mixture, separate each
ion from the others and confirm the presence of each ion. You will not be
allowed to begin your experiment until your instructor has confirmed that your
flow chart is prepared correctly. Your instructor will then give you a completed
chart that includes the amounts of each solution to be added in each step.
Silver, mercury(I) and lead(II) are often called the
“chloride” group because they form sparingly soluble to insoluble precipitates
with chloride ions. All three solids are white. The first step in isolating
these ions from a solution is to add HCl to form the chloride precipitates.
Silver and mercury(I) chlorides are much less soluble (Ksp
values of 1.8 x 10-10 and 1.3 x 10-18, respectively) than
lead(II) chloride (Ksp of 1.6 x 10-5). If solid
PbCl2 is heated in water to 100 °C for a few minutes, it will
dissolve. The other two chlorides will not. Lead(II) in solution will form an
insoluble white precipitate when allowed to react with sulfuric acid (Ksp
= 6.3x10-7). The addition of ammonia to solid silver chloride causes
the formation of a colorless silver-ammonia complex ([Ag(NH3)2]+,
Kf = 1.7 x 107). The addition of nitric acid will
cause the equilibrium to shift to free the silver which can then react with the
chloride again. Mercury(I) chloride reacts with ammonia to form Hg (metallic
liquid), HgNH2Cl (s, white), and Hg2O (s, black). The
solid mixture will have an overall grayish color.
Before examining an unknown mixture it is helpful to
observe the behavior of known ions in a mixture. You will separately
analyze two known mixtures of cations using the procedures outlined in Fig. 1.
Mixture A contains silver(I) (Ag+), mercury(I) (Hg22+),
aluminum (Al3+), barium (Ba2+) and potassium (K+)
ions. Mixture B contains lead(II) (Pb2+), iron(III) (Fe3+),
nickel(II) (Ni2+), magnesium (Mg2+) and zinc (Zn2+)
ions. There will also be solutions available that contain the individual ions
that you will be analyzing. You can use these solutions to confirm the behavior
of the ions in your mixture.
For each sample, you should record every step of the
analysis and your observations as you proceed in a table similar to the one
shown as Table 1. You should be as specific as possible when describing your
observations of the known mixtures so that you can use those observations to
identify your unknowns.
Table 1. Sample data table for recording the
results of each experimental step.
The first steps in the cation procedure are to identify
and remove the chloride group as described above. After these have been
isolated, the remaining cations will be separated based on their reactions with
hydroxide. The reactions of hydroxide ions with cations are very interesting. By
carefully controlling the pH of the solution, only certain metal hydroxides can
be caused to precipitate from solution or form soluble complexes. After the
chloride group ions are precipitated with hydrochloric acid, the solution will
be acidic. An ammonia/ammonium buffer is then created in order to make it just
neutral, which will shift the hydroxide concentration of the solution causing
the precipitation of only the most highly insoluble hydroxides, Fe(OH)3
(Ksp = 1.6 x 10-39) and Al(OH)3 (Ksp
= 3 x 10-34). It is important that the solution is not made overly
basic as the additional hydroxide will cause the [Al(OH)4]-
complex to form too soon. The remaining ions will either form soluble complex
ions with the added ammonia or remain dissolved in solution. After separation
from the supernatant, the aluminum hydroxide can be re-dissolved by increasing
the concentration of hydroxide ions with the addition of sodium hydroxide. This
addition will favor the formation of the complex, [Al(OH)4]-
(Kf = 2.0 x 1033). However, the iron(III) hydroxide
will not re-dissolve. Aluminum can be confirmed by adding aluminon, a dye, and
then making the solution alkaline with concentrated ammonia. The presence of a
pink lake (dyed precipitate) suspended in solution confirms the presence
of aluminum. Make sure that it is not the solution itself that is pink by
centrifuging. The presence of Fe3+ can be confirmed in two ways. Iron
forms a red complex with SCN- and a blue solid, KFe[Fe(CN)6],
upon the addition of K4Fe(CN)6.
After the iron and aluminum are removed from the solution,
the hydroxide concentration can be manipulated, again, to selectively
precipitate two cations. An increase in the concentration of hydroxide ions with
the direct addition of aqueous sodium hydroxide will cause the precipitation of
nickel hydroxide (Ksp = 6 x 10-16) and magnesium
hydroxide (Ksp = 6 x 10-10). The other cations will
remain in solution; zinc as the hydroxide complex [Zn(OH)4]-
and Ba2+ and K+ as the solvated ions. Like most
hydroxides, magnesium and nickel hydroxide can be dissolved in a solution that
is acidified and warmed. After re-establishing the ammonia buffer, the addition
of Na2HPO4 will cause the magnesium to slowly precipitate
as MgNH4PO4. Nickel will remain in solution in the form of
a nickel ammonia complex. The magnesium can be redissolved in hydrochloric acid.
The magnesium will form a blue lake in an alkaline solution containing the
organic compound 4-(p-nitrophenylazo)-resorcinol. Dissolved nickel ions will
form a deep pink precipitate upon the addition of another organic compound,
The only ions remaining in solution after the hydroxide
concentration is raised are zinc, barium and potassium. Barium forms an
insoluble precipitate with sulfate ions (Ksp = 1.1 x 10-10).
The barium can be further confirmed by the presence of a persistent green flame
test. Zinc can be precipitated by the addition of phosphoric acid, H3PO4
(Ksp for Zn3(PO4)2 is 5 x 10-36).
At this point, the only unknown ion remaining in solution
will be potassium. Potassium forms very few insoluble precipitates. The simplest
way to identify it is by a flame test after other ions are removed. The flame
will turn a fleeting violet color when exposed to potassium ions. Because this
color may be masked by the orange flame of sodium ions, the flame should be
viewed through a thickness of cobalt blue glass.
The strategy for the analysis of anions is similar to that
for cations; known reagents are added to a solution to selectively precipitate
dissolved anions. In this case, the precipitating reagents will be cations that
form insoluble salts with the dissolved anions. You will perform an analysis to
identify chloride (Cl-), iodide (I-), carbonate (CO32-),
sulfate (SO42-) and phosphate (PO43-).
The flow chart for the separation and identification of these anions is shown in
Fig. 2. You will be given only one known solution to analyze in this section.
This mixture will contain all five anions.
The first step in this procedure is the addition of barium
nitrate to cause the precipitation of BaCO3 (Ksp =
5.0 x 10-9), BaSO4 (Ksp = 1.1 x 10-10)
and Ba3(PO4)2 (Ksp = 6 x 10-39).
The addition of a strong acid (nitric, HNO3) to these precipitates
will adjust the solubility of the ions by taking advantage of their basic
nature. The reaction of CO32- and the acid will cause the
evolution of carbon dioxide gas. Because of the limited number of anions
possible in your procedure, the presence of this gas is a confirmation for the
carbonate ion. The nitric acid will also cause the dissolution of barium
phosphate. When the supernatant containing only barium phosphate is decanted and
made basic with Ba(OH)2 the precipitate will reappear. Unlike the
other two precipitates, barium sulfate will not redissolve when nitric acid is
added. The presence of a white solid after the acidification of the barium
precipitates is a confirmation of the sulfate ion.
The supernatant found after the addition of barium in the
previous step will contain the other anions in this procedure, I- and
Cl-, neither of which form insoluble salts with barium. (This
procedure could be performed on a fresh sample of the analyte, as well.) They
do, however, form insoluble salts with silver ions (Ksp(AgCl)
= 1.6 x 10-10, Ksp(AgI) = 1.5 x 10-16).
The solubility of these ions can be further decreased by acidifying the solution
with nitric acid. The precipitate should be washed to remove any other ions and
then stirred in clean distilled water. When aqueous ammonia and additional
silver nitrate are added to the mixture, the silver chloride will redissolve to
form the silver ammonia complex used in the detection of silver ions. The yellow
silver iodide will not redissolve. If the supernatant containing the complex is
acidified with nitric acid, the silver chloride will reprecipitate, confirming
the presence of chloride ions.
You will be given three unknowns to analyze. The first
will be a solution containing three cations. The second will be a solution
containing three anions. The third will be a solid binary salt. As soon as you
receive your unknowns, record your unknown numbers. You must include your
unknown numbers in your reports to receive credit. Use the procedures you
learned in the first two weeks to determine the compositions of your three
When analyzing unknown mixtures, you should keep a few
things in mind. Remember that you must have a confirmatory test for each cation
you believe is present. Even if you are told the number of ions present in your
mixture, you should not stop after finding that number of ions. It is possible
that you have made a mistake or have a false positive. Complete the entire flow
chart to make sure that no other species appear. If another is identified, you
should repeat the procedure on a fresh sample of analyte. You should have plenty
of your solution to repeat the entire analysis several times. There may be
penalties if you ask for extra, however, so be careful when using your unknown.
If you are ever uncertain as to whether a test is positive
for a given ion, you can repeat the test on the standard solutions provided to
confirm the behavior of that ion. After you have identified your unknown
mixtures, you may want to create a mixture containing the ions you believe are
present in your solution. If you have time, you can test this solution and
compare the results to your unknown mixture.
When you are analyzing for both cations and anions in a
single unknown, it is important to recognize that the analyses must be performed
separately. For example, the first step for a cation analysis is the addition of
hydrochloric acid. A test on a solution after HCl is added would, obviously, be
positive for chloride ions!
The various analytical procedures you will do must be
performed on ions dissolved in a solution. The first step in your solid unknown
analysis will be to dissolve the sample. Your sample may be insoluble or
sparingly soluble in distilled water! The fact that your unknown is insoluble in
water may give you an idea as to its identity. Try dissolving a small sample of
the solid first in water, then nitric acid or another aqueous solvent until you
find a solvent in which your unknown is completely soluble. You should then use
this solvent to prepare a mixture that is 1-3% of your unknown by weight. Try to
avoid using a solvent that contains any of the possible unknown ions!
Results and Analysis
Report to your instructor the identity of the ions in each
of your unknowns. Your instructor may allow you to repeat the
identification of an incorrect result, so you should submit your decision as
soon as possible after you have confirmed your unknowns’ identities. In the
Results section of your lab notebook you should include the balanced net
ionic equations of each reaction in the flow charts. You should attempt to
determine these reactions and may be able to find some of them in your textbook
or a reference text. Use equilibrium constants for the reactions to explain why
one ion precipitated or formed a complex while others remained solvated. Some of
the equilibrium constants you need are included in the procedures.
For this experiment you do not need to write any
conclusions until you have finished the analysis of your unknowns. Use the
reporting on physical phenomena from the Laboratory Notebook web page as a
starting point for your discussion.
Summary of Results
Use Table 2 to report the analysis of your unknowns.
Table 2. Summary table for reporting the identity of the unknowns.
- 1. Click here to download this file in PDF format
(link not yet active).
- 2. Wismer, Robert K. Qualitative Analysis with
Ionic Equilibrium; Macmillan Publishing Company: New York, 1991.
- 3. Zumdahl, S. S. Chemical Principles, 4th
Ed.; Houghton-Mifflin: New York, 2002, chapter 8.