Foothill College-Chemistry 1C
Name:
Separation and Identification of Group D Cations
(Cu2+, Ni2+, Mg2+ and Zn2+)
Objectives
•
•
•
To understand the chemical reactions involved in the separation and identification of the Group D cations.
To complete a flow diagram summarizing the qualitative analysis scheme for the Group D ions.
To successfully identify the Group D cation(s) in an unknown.
Background Chemistry and Discussion
The Group D cations are characterized as having hydroxides, oxides and oxalates that are soluble in an ammonia/ammonium
ion buffer solution. Three of the Group D cations, Cu2+, Ni2+ and Zn2+, form complex ions with ammonia. These complexes
ions are very stable and prevent the cations from precipitating when oxalate is added even though their oxalates are relatively
insoluble. Many complex ions are highly colored and the color of your solution at this stage may help in identification of the
cations in your unknown. Testing known solutions to determine the color of the complex ions formed is recommended.
Magnesium hydroxide will precipitate in aqueous ammonia, however the precipitate can be prevented from forming if the
solution contains a significant concentration of NH4+. This can be understood by considering the common ion effect and the
following equilibrium:
NH3(aq) + H2O(l) ⇔ NH4+(aq) + OH–(aq)
In an NH3/NH4+ buffer solution, this equilibrium is shifted left, reducing the hydroxide ion concentration so that Mg2+ does
not precipitate. Magnesium oxalate has a Ksp value that is relatively high and under the conditions of the experiment should
not have precipitated if a limited amount of oxalate was added in a previous separation step.
The presence of ammonia and possibly oxalate ions in the supernatant liquid containing Group D can interfere with
characterization tests for the cations in this group. Therefore, prior to testing the ammonia and oxalate, if present, must be
removed from the solution. This is accomplished by heating the solution in a crucible over a low flame to dryness, adding
concentrated HNO3 and then reheating to dryness. THIS HEATING MUST BE DONE IN THE HOOD! The following
reactions occur:
NH4+(aq) + Cl–(aq) ➝ NH4Cl(s) ➝ NH3(g) + HCl(g)
NH4+(aq) + NO3–(aq) ➝ NH4NO3(s) ➝ N2O(g) + 2H2O(g)
And if oxalate is present:
2NO3–(aq) + C2O42– + 4H+(aq) ➝ 2NO2(g) + 2CO2(g) + 2H2O(g)
Note that if oxalate is not present in your sample, then the HNO3 need not be added prior to heating; NH3 will still be
removed as shown in the reaction that does not involve the nitrate ion given above.
After heating to dryness, the residue contains the oxides of the Group D cations: Copper (II) oxide is black, nickel oxide is
green and magnesium and zinc oxide are both white. The residue is dissolved in 6 M HCl(aq) and the resulting solution can
be tested directly for Ni2+ and Cu2+. (In HCl, copper (II) and nickel ions form complex ions: CuCl42– and NiCl64–.) After
testing for Ni2+ and Cu2+, the solution is treated to remove Ni2+ and Cu2+ ions, if present, prior to testing for Mg2+ and Zn2+.
Ni2+ and Cu2+ Confirmation Tests:
Identification of Ni2+: A portion of the unknown solution is made basic with aqueous ammonia and then a drop of
dimethylglyoxime ((CH3)2C2(NOH)2, abbreviated as DMG) solution is added. The immediate formation of nickel
dimethylglyoxime complex, a cherry red precipitate, confirms the presence of Ni2+. Nickel ion forms the only highly
colored precipitate with dimethylglyoxiniate ions. The solution must be made basic with ammonia in order for the
precipitate to form since the DMG must be deprotonated before it will react with any Ni2+ that is present.
(CH3)2C2(NOH)2(aq) + NH3(aq) ⇔ (CH3)2C2(NOH)(NO)–(aq) + NH4+(aq)
2(CH3)2C2(NOH)(NO)–(aq) + Ni(NH3)62+(aq) ⇔ Ni[(CH3)2C2(NOH)(NO)]2(s) + 6NH3(aq)
Identification of Cu2+: The confirmation test for Cu2+ is the formation of a maroon precipitate of copper
hexacyanoferrate, Cu2[Fe(CN)6], from slightly acidic solution:
Dr. L.J. Larson
QualGroupD.doc
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Revised/Printed 6/28/08
Foothill College-Chemistry 1C
Qualitative Analysis Group D
2Cu2+(aq) + Fe(CN)64–(aq) ⇔ Cu2[Fe(CN)6](s)
In the absence of Cu2+ a white or pale green precipitate may be observed since hexacyanoferrate, [Fe(CN)6]2–, forms
precipitates with other Group D cations. Before adding hexacyanoferrate, the strongly acidic solution is first treated with
aqueous ammonia until it is only slightly acidic. The purpose of this is to avoid the addition of hexacyanoferrate to a
strongly acidic solution, since this would result in the formation of hydrogen cyanide (HCN), a very toxic gas. By
reducing the acidity of the solution, this possibility is avoided. HOWEVER, THIS TEST MUST STILL BE
PERFORMED IN THE HOOD. After adjusting the pH of a portion of the test solution, potassium hexacyanoferrate,
K4[Fe(CN)6], is added. The immediate formation of a maroon precipitate confirms the presence of copper(II) ion.
Separation of Mg2+ and Zn2+ from Cu2+ and Ni2+:
Before testing for magnesium and zinc ions, the copper (II) and nickel ions are removed. This is accomplished by
precipitating the copper (II) and nickel ions as their sulfides. The source of sulfide ion will be a hot, weakly acidic
solution of sodium thiosulfate, Na2S2O3. Under these conditions sodium thiosulfate disproportionates to form sulfate and
sulfide ions. The unbalanced net-ionic chemical equation for this reaction is:
S2O32–(aq) ⇔ S2–(aq) + SO42–(aq)
Copper (II) sulfide is insoluble, even in acidic conditions, and nickel sulfide has only slight solubility in the weakly
acidic conditions imposed. Zinc sulfide is soluble under the conditions used; it does not precipitate. Magnesium does not
form a sulfide. Thus, upon heating in weakly acidic sodium thiosulfate solution, Cu2+ and Ni2+ ions will precipitate,
forming CuS(s) and NiS(s), both of which are black. Magnesium and zinc ions are left in the supernatant that is used to
test for their presence. Due to the slight solubility of NiS in the weakly acidic solution, any Ni2+ ions present may not
completely precipitate leaving some Ni2+ in solution along with the Zn2+ and Mg2+ ions.
Mg2+ and Zn2+ Confirmation Tests:
The identification of magnesium is accomplished using the supernatant remaining after the precipitation of Cu2+ and Ni2+
as their sulfides. This supernatant may contain Zn2+ and Mg2+, and possibly some Ni2+. Magnesium and zinc ions are
typically colorless in solution; the presence of Ni2+ may give a characteristic color to the solution. The solution is made
slightly basic with aqueous ammonia, and disodium hydrogen phosphate, Na2HPO4, is added. Under these conditions,
Zn2+ will precipitate as white zinc phosphate, Zn3(PO4)2, and Mg2+ will precipitate as white magnesium ammonium
phosphate, MgNH4PO4. Any other Group D cation present at this point remains dissolved since their phosphates are
soluble in ammonia. An acid-base chemical equilibrium is involved is these precipitation reactions as shown below:
Acid Base Equilibria:
NH3(aq) + HPO42–(aq) ⇔ NH4+(aq) + PO43–(aq)
Precipitation Reactions:
Mg2+(aq) + NH3(aq) + HPO42–(aq) ⇔ MgNH4PO4(s)
3Zn2+(aq) + 2NH3(aq) + 2HPO42–(aq) ⇔ Zn3(PO4)2(s) + 2NH4+
If the solution is made too basic with ammonia, a precipitate may not form. In this case, adding a small amount of
HCl(aq) may work to produce a precipitate. After being centrifuged, the supernatant is discarded and the precipitate is
treated with NaOH. Any zinc phosphate will dissolve to form the colorless tetrahydroxozincate(II) complex ion,
[Zn(OH)4]2–. The unbalanced net ionic chemical equation for this reaction is:
Zn3(PO4)2(s) + OH–(aq) ⇔ [Zn(OH)4]2–(aq) + PO43–(aq)
Magnesium ammonium phosphate is insoluble in NaOH solution. A white precipitate that remains after addition of the
NaOH confirms the presence of Mg2+. The basic supernatant is tested for the presence of zinc by making the solution
slightly acidic with aqueous acetic acid and then adding K4[Fe(CN)6]. THIS TEST MUST BE PERFORMED IN THE
HOOD. The immediate formation of a white precipitate of Zn2[Fe(CN6)] confirms the presence of zinc(II) ion.
2Zn2+(aq) + Fe(CN)64–(aq) ⇔ Zn2[Fe(CN)6](s)
For both the Mg2+ and Zn2+ tests, the precipitates formed are white. The existence of traces of colored ions such as Ni2+
and Cu2+ can add some color to the white precipitates. Modifications to the procedure may be needed if results are
ambiguous.
Dr. L.J. Larson
QualGroupD.doc
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June 28, 2008
Foothill College-Chemistry 1C
Reagents Available
6-M NH3
6-M HCl
6-M NaOH
6-M acetic acid (CH3 COOH)
Concentrated HNO3
Dilute, Known Solutions of Cu2+, Ni2+, Mg2+ and Zn2+
Qualitative Analysis Group D
1% dimethylgloxime
0.1-M K4Fe(CN)6
Solid Na2S2O3
Saturated Na2HPO4 (about 0.5 M)
Safety and Waste Disposal
Some of Group D cations and reagents used are toxic. The HCl, HNO3, CH3COOH, NH3 and NaOH are irritants.
Avoid contact and wash immediately if any is spilled or splashed on you. Wear eye protection at all times.
As you perform the experiment, collect all waste solutions in a waste beaker. This mixture should then be discarded in
the appropriate waste container. DO NOT POUR ANY OF THE SOLUTIONS DOWN THE DRAIN.
Unknowns and Knowns
Separate known samples containing Cu2+, Ni2+, Mg2+ and Zn2+ are provided for your use. Testing known samples is
helpful in this analysis since doing so will allow you to observe what a positive test looks like. It is usually convenient to
test a known sample simultaneously with your unknown. To test a known sample, you can either prepare a known mixture of
the Group D cations or the known Group D cation solutions can be tested separately. In the case of a known mixture, steps in
the procedure that are required for separation of the cations must be followed before the confirmation tests can be conducted.
If a known is prepared that contains only one of the cations, then the confirmation test can be conducted directly on the
known. To prepare a known sample for testing, add 2 to 3 drops of the solution(s) containing the cation(s) you wish to test to
about 0.5 mL of water. Note that the experimental conditions such as pH, oxidation state, etc. for the known test must
be the same as that for the unknown.
Outline of Procedure: Use a flow diagram in your notebook to record all
observations for each step in the procedure.
Chemistry and Relevant Background Information:
1. Removal of ammonia and oxalate:
The presence of ammonia and possibly oxalate ions may
interfere with later tests. These are removed from the sample
solution by heating to dryness in the presence of
concentrated nitric acid. The net ionic chemical equations for
the reactions that take place were given in the Background
Chemistry and Discussion section. You should rewrite them
here for reference.
Experimental Procedure:
1. CAUTION!!! The following must be done in the hood!
(a) Pour the supernatant liquid containing the Group D
cations into a clean crucible. Place the lid on the crucible.
slightlv ajar, and heat the solution to dryness over a low
flame. If your sample does not contain oxalate, skip to step
1(d). Otherwise continue with step 1(b).
(b) Cool the crucible for 5 minutes, and then add 6 drops of
concentrated nitric acid (Caution: HNO3 is corrosive and
an oxidizer), washing the inside of the crucible. Replace the
lid and heat to drvness once again.
(c) Cool the crucible for 5 minutes, repeat the addition of
concentrated nitric acid and heat to dryness.
(d) Cool the crucible for 5 minutes, and dissolve the residue
by adding 5-10 drops of 6 M HCI. (It is alright if all of the
solid does not dissolve.) Using a clean Pasteur pipet, transfer
the solution to a clean, labeled test tube.
(e) Rinse the crucible with 5 drops of deionized water. Add
this rinse to the same test tube.
2. Identification of Nickel Ion:
2.
To test for Ni2+ a portion of the solution from step 1 is made
(a) Transfer one drop of the solution from step 1 to a clean
Dr. L.J. Larson
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June 28, 2008
QualGroupD.doc
Foothill College-Chemistry 1C
Qualitative Analysis Group D
Chemistry and Relevant Background Information:
basic by adding NH3. This is followed by addition of
dimethylglyoxime solution. Formation of a cherry-red
precipitate confirms the presence of Ni2+.
Experimental Procedure:
test tube.
(b) Add 6 M NH3, with stirring, until the solution tests basic
to litmus paper.
In the space below, write the balanced net ionic chemical
equation, including phase labels, for the formation of the
cherry red precipitate (Refer to the Background
Chemistry and Discussion section).
(c) Add 1 drop of 1% dimethylglyoxime.
3. Identification of Copper(II) Ion:
To test for Cu2+ a portion of the solution from step 1 is made
weakly acidic by adding NH3. This is followed by addition
of K4[Fe(CN)6] solution. Formation of a maroon precipitate
confirms the presence of Cu2+. CAUTION!!! If the solution
is too acidic, toxic HCN(g) will be produced.
3. CAUTION!!! The following must be done in the hood!
(a) Transfer one drop of the solution from step 1 to a clean
test tube.
(b) Add 6 M NH3 until the solution tests only weakly acidic. If
the solution becomes basic, use 6 M acetic acid to make it
slightly acidic.
In the space below, write the balanced net ionic chemical
equation, including phase labels, for the formation of the
maroon precipitate (Refer to the Background Chemistry
and Discussion section).
(c) Add 3-4 drops of 0.1 M K4[Fe(CN)6], solution.
4. Separation of Cu2+ and Ni2+ from Mg2+ and Zn2+
Copper (II) and nickel ions are separated from magnesium
and zinc by precipitating the Cu2+ and Ni2+ as sulfides.
Magnesium ions do not form a sulfide and the Zn2+ ions
remain in solution under the conditions imposed. The source
of sulfide ion is the disproportionation of thiosulfate ion in
hot, weakly acidic solution.
Refer to the Background Chemistry and Discussion
section for the unbalanced net ionic chemical equation
for the disproportionation of thiosulfate ion. Balance this
equation and write the balanced net ionic chemical
equation, including phase label, in the space below:
4.
(a) Add 6 M NH3 to the remainder of the solution from step 1
until it tests slightly basic (pH 8-9). Then add 6 M acetic
acid until the solution is weaklv acidic (pH 4-5).
5. Identification of Magnesium Ion:
The solution from step 4 is made basic by adding NH3 and
then sodium hydrogen phosphate is added. Under these
consitions, Mg2+ will precipitate as MgNH4PO4 and Zn2+ will
precipitate as Zn3(PO4)2. Both of these are white solids.
5.
(a) To the supernatant from step 4(c) add 6 M NH3 until the
solution tests slightly basic to litmus paper (pH 8).
(b) Add about 0.2 g of solid sodium thiosulfate, Na2S2O3, and
heat for 5 minutes in a boiling water bath. Cool for 1 minute
by swirling the test tube in cold tap water.
(c) Centrifuge and decant the supernatant into a clean,
labeled test tube. The precipitate may be discarded.
(b) Add 6-8 drops of saturated (about 0.5 M) Na2HPO4
solution. Stir and then cool the solution in ice water for
several minutes.
The net ionic chemical equations for the reactions that
take place were given in the Background Chemistry and
Discussion section. You should rewrite them here for
reference.
Dr. L.J. Larson
QualGroupD.doc
(c) If a precipitate does not form, add 6 M HCl dropwise to
lower the pH slightly. Watch carefully for precipitate
formation; stop adding the HCl upon formation of a
precipitate.
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June 28, 2008
Foothill College-Chemistry 1C
Qualitative Analysis Group D
Chemistry and Relevant Background Information:
Experimental Procedure:
(d) Centrifuge and discard the supernatant.
(e) Wash the precipitate with 10 drops of deionized water.
Discard the wash.
To the precipitate formed, 6 M NaOH is added. Any
Zn3(PO4)2 present will dissolve due to the formation of the
[Zn(OH)4]2– complex ion. Any MgNH4PO4, if present, will
not dissolve, indicating the presence of Mg2+.
(f) To the washed precipitate, add 6 drops of 6 M NaOH and
stir thoroughly. Centrifuge and decant the supernatant into a
clean, labeled test tube.
In the space below write the balance net ionic chemical
equation, including phase labels, for the reaction where
Zn3(PO4)2 is dissolved in NaOH(aq) solution:
6. Identification of Zinc Ion:
The supernatant from step 5(e) is tested for the presence of
zinc ion. To test for Zn2+ the solution from is made weakly
acidic by adding NH3. This is followed by addition of
K4[Fe(CN)6] solution. Formation of a white precipitate
confirms the presence of Zn2+. CAUTION!!! If the solution
is too acidic, toxic HCN(g) will be produced.
6. CAUTION!!! The following must be done in the hood!
(a) To the supernatant from step 5(f), add 6 M acetic acid until
the solution tests only weakly acidic (pH 4-5). If the solution
becomes too acidic, use 6 M NH3 to make it slightly acidic.
(b) Add 4 drops of 0.1 M K4[Fe(CN)6], solution.
In the space below, write the balanced net ionic chemical
equation, including phase labels, for the formation of the
precipitate (Refer to the Background Chemistry and
Discussion section).
Dr. L.J. Larson
QualGroupD.doc
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June 28, 2008
Foothill College-Chemistry 1C
Qualitative Analysis Group D
Prelab Exercise:
1.
Name:
After adding concentrated HNO3, heating to dryness and then dissolving the residue in 6 M HCl a Group D unknown is
treated according to the procedure for the D analysis. For each step below, answer the questions asked.
a) One drop of the solution is made basic with NH3 and then 1% dimethoxylgloxime is added. A cherry red preciptate
forms. What can you conclude?
b) One drop of the solution is made weakly acidic by adding NH3. When 0.1 M K4[Fe(CN)6 ] is added a light green
precipitate results. What can you conclude?
c)
The remainder of the solution is made weakly acidic by adding NH3. When sodium thiosulfate is added and the
solution is warmed, a black precipitate and a colorless supernatant results. What can you conclude?
d) The supernatant from step c is made basic by adding ammonia. Adding 1 M Na2HPO4 yields a white precipitate.
What can you conclude?
2.
e)
The precipitate from step d is treated with 6 M NaOH resulting in a white precipitate and a colorless supernatant.
What can you conclude?
f)
The supernatant from step e is made weakly acidic by adding acetic acid. When 0.1 M K4[Fe(CN)6 ] is added no
precipitate forms. What can you conclude?
Describe a simple, ONE step test that would allow you to distinguish between Cu2+ and Zn2+.
Dr. L.J. Larson
QualGroupD.doc
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June 28, 2008
Foothill College-Chemistry 1C
3.
Qualitative Analysis Group D
Complete the flow diagram below for Group D analysis. For each step, indicate the chemical form of each Group D
cation that is present at the given step (for example CuO, Cu(NH3)42+, CuS, etc. Add colors of solutions and precipitates
where known.
snt ≡ supernatant
ppt ≡ precipitate
∆ ≡ heat
Cu(NH3)42+, Ni(NH3)62+, Zn(NH3)42+ and Mg2+
(1) ∆ to dryness
(2) If oxalate present add Conc. HNO 3, ∆ to dryness
6M HCl
1 drop
6 M NH3
until basic
+
1% DMG
1 drop
6 M NH3 until
slightly acidic
+
0.1 M K4[Fe(CN)6]
6 M NH3 until slightly
basic then 6 M
CH3COOH until
slightly acidic (pH 4-5)
+
0.2 g Na2S2O3, ∆
snt
ppt
Discard
6 M NH3 until
pH 8
+
1 M Na2HPO4
6 M NaOH
snt
ppt
6 M CH3COOH
until slightly acidic
+
0.1 M K4[Fe(CN)6]
Dr. L.J. Larson
QualGroupD.doc
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June 28, 2008
Foothill College-Chemistry 1C
Qualitative Analysis Group D
Follow-up Questions:
Name:
For numerical problems, you must show all work for credit!
1. The test for Cu2+ and Zn2+ ions involves adding K4[Fe(CN)6] to a weakly acidic test solution. If the solution is too acidic,
toxic HCN(g) will form. Write the balanced net ionic chemical equation, including phase labels, showing the production
of HCN(g) when excess strong acid is added to [Fe(CN)6]4–(aq).
2.
One of the reactions used in the Group D cation analysis is a disproportionation reaction.
a) What is a disproportionation reaction?
b) What was the purpose of using this reaction? In other words, what did it accomplish?
c)
Balance this disproportionation reaction and then write the balanced net ionic chemical equation, including phase
labels, in the space below.
3.
Cu(OH)2, Zn(OH)2 and Ni(OH)2 are all insoluble in water. Explain why they do not precipitate with the Group B cations
when the solution is made basic with ammonia.
4.
In the Group D analysis scheme, Zn3(PO4)2(s) is dissolved by adding 6 M NaOH.
a) Write the balanced net ionic equation, including phase labels, for this reaction.
b) Given that Ksp for Zn3(PO4)2(s) is 9.0x10–33 and that K f for [Zn(OH)4]2– is 4.6 x 1017, calculated Knet for the reaction
written in part (a).
c)
Is this Knet consistent with the fact that Zn3(PO4)2(s) dissolves when 6 M NaOH is added? Why or why not?
Dr. L.J. Larson
QualGroupD.doc
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June 28, 2008