Qualitative Analysis


Qualitative analysis is used in the determination of the identity of a
substance. It is different from quantitative analysis, which deals with the
determination of the amount of a substance. In this experiment, qualitative
analysis techniques are used to determine whether or not a sample contains a
certain ion. When using this method, an unknown and a reactant are mixed. The
result of the reaction leads to a conclusion about the presence or absence of
certain ions in the unknown. Many ions react in similar ways, and although the
addition of one reagent to an unknown may not identify the ion, it limits the
possibilities as to what the ion could be. A sequence of reactions used to
analyze a sample is called a scheme, and it usually requires a large number of
reagents and separation steps. For this experiment, the unknown may contain
anywhere from 2 to all of the following cations and anions: Cations Anions Ag+

Cl- Ba2+ SO42- Fe3+ PO43- Cu2+ Cr3+ The following reagents are used to identify
the ions: 1M H2SO4 2M HCl 2M NH4OH (labeled as NH4+) 2M NaOH .1M Ba(NO3)2
(labeled as .1M Ba2+) .1M AgNO3 (labeled as .1M Ag+) The first four are used to
identify the cations, and the last two, used in conjunction with the first four,
are used to identify the anions. The identification of the ions is mainly based
on solubilities. This means that something must be known about the solubility
characteristics of the different ions in the presence of the available reagents.

The point of the first part of the experiment is to learn which reagents cause
the ions to form precipitates, and which reagents dissolve the precipitates
formed by the ions. This information is used to make the flow charts for the
identification on the unknown ions. For example, it is important to know that a
certain reagent will dissolve the precipitate formed by one ion, while it will
not dissolve the precipitate formed by another ion. This can be used to
distinguish between two different precipitates present in a solution, or to
confirm which ion formed the precipitate and therefore was present in the
solution. When carrying out the reactions, avoid adding an excess of reagent to
the solution. This is because some precipitates redissolve in an excess of the
reagent. Therefore, in cases where one drop of reagent produces a precipitate, 3
or more drops could completely dissolve the precipitate without it ever being
visible to the eye. This would cause a large error in the scheme developed to
identify the unknown ions. Experimental: The first part of the experiment
consists of reacting the cations and anions with the reagents in order to see
what the reaction will result in (precipitate or no precipitate). The cations
were each reacted with the first four reagents listed in the introduction
(H2SO4, HCl, NH4+, and NaOH). Then, the anions were each reacted with Ba2+ and

Ag+. This was done by placing 2 drops of the ion in the test tube and then
adding 2 drops of reagent. Each cation was reacted with each of the 4 reagents
before moving on to the next cation to be tested. Prior to performing the
reactions, a chart was made like the one in the data and calculations section.

As each reaction was performed, the chart was filled in with the observation of
what happened. If there was no change, NR was written in the chart for "no
reaction." If a precipitate formed, the color of the precipitate was written
in the chart. If there was no precipitate but there was a color change in the
solution, that was also recorded. As each reaction was carried out, it was
sometimes difficult to determine whether a precipitate formed or not. If there
was uncertainty, the test tubes had to be placed into the centrifuge in order to
separate the precipitates from the solution. There are some very important
things to remember when using the centrifuge. First, when tubes are placed in
the centrifuge, a tube with an approximately equal volume of solution should be
placed exactly opposite each sample tube to counterbalance it (use a test tube
filled with an equivalent amount of water if necessary). Second, the centrifuge
should come to a stop before it is opened and the test tubes removed. This is to
avoid injury. Once the tubes were removed from the centrifuge, it was obvious
whether there was a precipitate present or not. If a solid has settled onto the
bottom or side of the test tube, there was a precipitate present. If the tube
appears to contain the same solution as before the test tube was placed in the
centrifuge, no reaction occurred. The next part of the experiment consists of
determining which reagents dissolve certain precipitates. This information can
be especially helpful when determining the ions present in the unknown. The
precipitates tested were AgCl, BaSO4, and Ag3(PO4). They were reacted with HCl,

H2SO4, NH4OH, and NaOH. This was done by making the precipitate using the
information from the first chart, and then adding 2 drops of reagent. For
example, the precipitate AgCl was made by reacting Ag+ with HCl. Four samples of
this were prepared, and each of the reagents was added to the samples to see if
the precipitate dissolved. A chart was filled in with the results of the
reactions. In the final part of the experiment, the unknown was tested to
determine which ions were present in it. This was done using flow charts created
with the information from part 1 of the experiment (see data and calculations
section). To test for the ions in unknown #2, it was first made into a solution
by adding 25 mL of distilled water to the sample in a 100 mL beaker. It was
mixed until all of the solid dissolved. To speed up the dissolving process, the
beaker was held in the palm of the hand in order to slightly heat the solution.

Once the solution was ready, it was tested for the ions by following the flow
charts. For each step, 2 drops of reagent was added to 2 drops of unknown
solution. To test for the cations, the cation flow chart was followed. First,

HCl was added to the solution. There was no reaction, so H2SO4 was added to
another sample of the unknown solution. This also resulted in no reaction, so

NaOH was added to another sample of the solution. A rusty-brown precipitate
appeared, which meant either PO43- or Cu2+ was present. To determine which one
of these ions was present, NH4+ was added to the solution, and a rust-colored
precipitate formed. This confirmed the presence of Fe3+. Next, the unknown had
to be tested for anions. The anion flow chart was followed. 2 drops of unknown
were reacted with Ba2+, and there was no reaction. 2 more drops of unknown were
reacted with Ag+, and white and tan precipitates formed. H2SO4 was added to the
test tube containing the precipitates, and a white precipitate was left. This
confirmed that PO43- was present but dissolved when the H2SO4 was added (as was
found in part 1 of the experiment), leaving the Cl-. Therefore, the anions
present were Cl- and PO43-. Data and Calculations: The data charts and the flow
charts are on the following pages. Unknown #2 contains Fe3+, Cl-, PO43-. Net
ionic equations for precipitates and reactions on flow charts: Ag+ + Cl- AgCl

Ba2+ + SO42- Ba SO4 Fe3+ + 3OH- Fe (OH)3 Cu2+ + 2OH- Cu (OH)2 Cr3+ +

3OH- Cr (OH)3 3Ag+ + PO43- Ag3(PO4) Ag3(PO4) + H2SO4 Ag2(SO4) + H3PO4

Thought process for flow charts: 2 separate flow charts had to be made, one for
the cations and one for the anions. Starting with the cations, the flow chart
must begin by listing all cations possibly present because the unknown can
contain any number of them. HCl was the first reagent added on the flow chart
because it only produced a precipitate with one of the cations, Ag+. This was
determined using the data chart from part 1 of the experiment, where the
precipitates formed with each reagent were clearly delineated. By beginning the
flow chart with reagents that produce fewer precipitates and ending it with the
ones that produce more, the chart was easier to follow during the testing for
ions. Therefore, the next reagent used on the chart was H2SO4. Since the Ag+
precipitated out as AgCl, the ions left to react with H2SO4 were Ba2+, Fe3+,

Cu2+, and Cr3+ (these were the ions that were left unreacted by the HCl). By
referring to the data chart from part 1 of the lab, it was found that H2SO4 only
formed a precipitate with Ba2+(BaSO4). This meant that the Fe3+, Cu2+, and Cr3+
ions were left unreacted. NaOH was added to these ions and it formed a
rust-colored precipitate with Fe3+, while it formed a blue precipitate with

Cu2+. To confirm whether the precipitate was from Fe3+ or Cu2+, NH4+ was added
to the unknown. If a rusty-brown precipitate appeared, the ion present was

Fe3+(Fe(OH)3). If no precipitate formed but the solution turned dark blue, Cu2+
was present. The only ion left unreacted after the NaOH was added was Cr3+. If a
bluish-white precipitate formed when NH4+ was added to the Cr3+, it confirmed
the presence of Cr3+(Cr(NH4)3). The anion flow chart began with all three of the
anions listed (SO42-, Cl-, PO43-). Ba2+ was the first reagent added because it
formed a precipitate with only one of the ions, SO42-(BaSO4). The ions left
unreacted were Cl- and PO43-. Ag+ was added to these. At this point, a white
precipitate could form with Cl-(AgCl), or a yellow precipitate could form with

PO43-(Ag3(PO4)). If the precipitate was purely white, the ion was Cl-. This was
confirmed by adding H2SO4 to the precipitate, which would result in no reaction.

This is because H2SO4 does not dissolve AgCl (from part 1 of the experiment).

Then, the addition of NH4+ would dissolve the precipitate and prove that only Cl-
was present. However, if the unknown contained both Cl- and PO43-, both
precipitates would form, but the two colors of the precipitates would be
indistinguishable from each other. A yellow precipitate would be seen, but it
would be impossible to tell if there was also a white precipitate present.

Therefore, H2SO4 would again be added to the precipitates. If a white
precipitate appeared, it would mean that PO43- had been present, but it was
dissolved by the H2SO4. It would leave only the AgCl precipitate visible (H2SO4
dissolves Ag3(PO4), but not AgCl), but because the precipitate originally had a
yellow color to it, it would be known that both the Cl- and the PO43- ions were
present in the unknown solution. If the solution had no precipitate left (turned
clear) when the H2SO4 was added to the yellowish precipitate, it would indicate
that only PO43- was present in the solution. This is because H2SO4 dissolves

Ag3(PO4), and there were no other precipitates left in the solution to be seen.

Results and Discussion: In conclusion, unknown #2 contained Fe3+, Cl-, PO43-.

This was determined using qualitative analysis, and the purpose of the
experiment was therefore fulfilled. One possible source of error for this lab
could occur in part 1, where the reactions of different reagents with different
ions are recorded in data charts. If there is incorrect information about
whether or not a precipitate formed, it will most likely result in an incorrect
flow chart and an incorrect identification of the ions in the unknown. That is
why it is important to use the centrifuge if there is uncertainty about a
precipitate, or the reaction should be performed again. Another source of error
would be to add too many drops of reagent to the ion or sample of unknown. This
is because a precipitate may form with the reagent, but dissolve in an excess of
the reagent. Therefore, the precipitate could form but then be dissolved without
ever being seen. This would also result in an incorrect flow chart and an
incorrect identification of the unknown ions. Another source of error could
occur in making or following the flow chart. Incorrect reasoning when designing
the flow chart will make it difficult to correctly identify the ions in the
unknown, and not following the flow chart correctly would obviously cause error
in the final results.