Chemistry 203
Experiment 7 – Group I Cations (Adapted from DePaul University)
Objective: To qualitatively determine the presence or absence of specific cations from a
mixture. This will be done with a set of knowns, and a set of unknowns.
Physical and Chemical Principles
Silver (I), lead (II), and mercury (I) ions are separated from other cations by addition of
HCl to form their insoluble chlorides (Eq. 1 – 3). If they were not precipitated as
chlorides in group I, they would precipitate as insoluble sulfides in group II. If a large
excess of chloride ions is used to precipitate silver (I) or lead (II), they can form the
soluble complex ions AgCl2− or PbCl42− (Eq. 4 and 5, respectively). Although it is
advisable to test for complete precipitation of these chlorides, especially in a general
unknown scheme, no more than a drop of excess HCl should be added.
Ag+(aq) + Cl−(aq)
AgCl (s)
Eq. 1
Pb2+(aq) + 2Cl−(aq)
PbCl2 (s)
Eq. 2
Hg22+(aq) + 2Cl−(aq)
Hg2Cl2 (s)
Eq. 3
AgCl (s) + Cl−(aq)
AgCl2−(aq)
Eq. 4
PbCl2 (s) + 2Cl−(aq)
PbCl42−(aq)
Eq. 5
Lead
Lead (II) chloride is about three times more soluble in hot water than in cold water; this
fact is used to separate lead (II) chloride form the other two chlorides of group I. In fact,
lead (II) may not precipitate out in group (I) if the solution is too warm or the lead ion
concentration is not great enough to exceed the Ksp for lead chloride. Because of this, the
more insoluble sulfide could be how the lead precipitates out, making it is necessary to
account for lead (II) in cation group II as well as in group I.
It may be necessary to wash the precipitate more than once with hot water to dissolve all
the lead chloride so it does not interfere with subsequent tests, especially for an unknown.
The presence of lead (II) ions is confirmed by the formation of a more insoluble
precipitate of lead (II) chromate (yellow), lead (II) sulfate (white), or lead (II) iodide
(yellow).
Pb2+(aq) + CrO42−(aq)
PbCrO4 (s)
Eq. 6
Pb2+(aq) + SO42−(aq)
PbSO4 (s)
Eq. 7
Pb2+(aq) + 2I−(aq)
PbI2 (s)
Eq. 8
Mercury
Ammonia is added to separate the mercury ions from silver ions. Because silver forms an
ammine complex while mercury does not, silver dissolves instead of mercury. At the
same time, the ammonia solution promotes a disproprtionation reaction to produce two
mercury containing products.
Hg2Cl2 (s) + 2NH3 (aq)
HgNH2Cl (s) + NH4+(aq) + Cl−(aq) + Hg(l)
Eq. 9
Although the mercury (II) amidochloride, HgNH2Cl, is a white compound, the precipitate
normally appears gray or black because of the formation of elemental mercury. Mercury
is thus verified by the formation of a dark precipitate.
Silver
With the addition of ammonia to separate the mercury and silver ions, the reaction:
AgCl (s) + 2NH3(aq)
Ag(NH3)2+(aq) + Cl−(aq)
Eq. 10
which is a combination of:
AgCl (s)
Ag+(aq) + Cl−(aq)
Eq. 11
Ag(NH3)2+(aq)
Eq. 12
and
Ag+(aq) + 2NH3 (aq)
causes AgCl(s) to dissolve so that it can be separated from mercury. In order to
determine whether silver (I) is present, it is necessary to reprecipitate the AgCl(s). This
is done by adding acid, which reacts with the ammonia in equilibrium with the
diamminesilver(I) ion (Eq. 13). The complex ion then dissociates to replace the ammonia
that is neutralized (Eq. 14), and as the Ag+ concentration increases, the Ksp of the AgCl(s)
will be exceeded (Eq. 1). White AgCl (s) soon precipitates, thus confirming the presence
of silver (I) ion.
NH3 (aq) + H+(aq)
Ag(NH3)2+(aq)
NH4+(aq)
Eq. 13
Ag+(aq) + 2NH3 (aq)
Eq. 14
Experimental
Table 1: Equipment and Chemicals for Group I Cation Analysis
Equipment
Quantity
Beaker, 250 mL
1
Dropper bottles
4
Hot plate (or Bunsen burner with ring stand, ring, and wire gauze)
1
Pipets, disposable, (or eyedroppers)
as needed
Test tubes
as needed
Chemicals
Quantity
Ammonium hydroxide (NH4OH), 6 M (student droppers)
as needed
Group I cations, 0.1 M
as needed
Group I unknown
1 mL
Hydrochloric acid (HCl), 6 M (student droppers)
as needed
Nitric acid (HNO3), 6 M (student droppers)
as needed
Potassium chromate (K2CrO4), 1 M
as needed
Sulfuric acid (H2SO4), 6 M (student droppers)
as needed
Water, deionized
as needed
Chemicals marked as “as needed” are generally available on the benchtop.
Procedure
The procedure for this experiment is given in flowchart form as shown in Figure 1.
add 10 drops 6M HNO3
Figure 1: Flow-chart procedure for Group I qualitative analysis.
Waste Disposal
Dispose of all Group I materials in the container labeled Group I Waste.
Retain your 6 M acids and bases for the next experiment.
Report
The report for this experiment is due at the beginning of the next laboratory period. It
must be neatly typewritten and contain the following elements:
Abstract
Write a brief paragraph describing the experiment and summarizing the findings. It
should 4-5 sentences in length, and end with a statement that identifies the ions present in
the unknown solution.
Background
Describe the chemistry going on in the reactions found in the flow chart. Use the
chemical equations to describe the reactions taking place. Describe what you would be
looking for in each reaction to confirm the presence of a cation.
Results
Use observations from the experiment to confirm the presence or absence of each cation
in the solution. In one clear, complete sentence, state the unknown number and the ions
that are present in the unknown. Include the net ionic equations for all of the reactions for
the ions that are present in the unknown. Equations may be typed or handwritten. In
either case, the equations should be easy to read, with clear notation of the phase of each
participant (aqueous, liquid, solid…) and careful attention to superscript and subscript
notations. Sets of equations that are poorly organized or illegible will not receive any
credit.
Discussion Questions
1. Briefly explain why it is necessary to wash the precipitate with hot water after the
lead chloride is dissolved in hot water to separate it from the other chloride
precipitates.
2. Could HNO3 be substituted for the HCl in the original precipitation of Group I?
Why or why not?