Qualitative Analysis of Cations

Goals
In this experiment, you will use qualitative analysis to identify the cations in known and unknown samples. By observing the reactions of known cation samples, you will construct a logical flow chart for identifying the components of a mixture of unknown cations. This experiment also illustrates a practical application for the theory of solubility equilibria.

Key Questions
1. What ions are present in my test tube?

2. What ions are in solution?

3. What is the precipitate?

4. What do I expect to happen as I add the next reagent?


Chemistry & Background
Qualitative analysis schemes are often based on selective precipitation of different cations. By carefully selecting reagents and conditions for reactions, groups of ions can be separated on the basis of their reactivity and the solubility of their salts. This enables the experimenter to use simple techniques with standard laboratory equipment, as well as to learn some descriptive chemistry.

The Separation Scheme
The possible cations in your unknown sample are given in the following list, organized according to the separation scheme. These ions should all be familiar to you.

Group I: Pb2+, Ag+Group II: Cu2+Group III: Fe2+, Ni2+, Mn2+

The first step in any scheme separates the ions into several groups, each of which exhibits a common chemical property. The groups are then examined to identify the components. The unknown solution contains ions from three groups which may be separated as follows:

1. The Group I ions form insoluble chlorides upon addition of dilute HCl.
2. In the absence of Group I ions, the Group II ions are insoluble in acidified H2S solution.
3. After Group II ions have been removed, the Group III ions are insoluble in a basic solution saturated with H2S.

Once the ions are separated into these three groups, additional reactions are used to separate and analyze for individual cations. For example, Pb2+ and Ag+ ions are separated by dissolving the PbC12 precipitate in hot water and removing it from the less-soluble AgCl. Following this separation, a test is performed on the solution to confirm the presence of Pb2+. The individual separation reactions and confirmation tests for each ion are outlined in the procedure section.

These procedures will provide crucial information regarding the presence or absence of particular ions in your mixture. You will obtain an "unknown" solution, which contains a subset of the six ions, and you will use your scheme to determine its composition. Some of the experiments may be unnecessary for your particular combination of unknown cations, but that is for you to decide. In any event, you should convince yourself that you have devised a reliable scheme before you pick up your unknown sample.

In addition to your unknown cation mixture, you will also have in lab a set of 0.1 M solutions of each individual cation. You can carry out parallel experiments with these solutions and your unknown mixture of cations, to see what a positive test looks like. If your unknown does not contain a particular ion, you should do the confirmation test for that cation on the known sample and record your observations in your notebook. Note that it is not appropriate to ask your TA or instructor questions like, "Is this the color of the precipitate that copper gives with thioacetamide?" or "Is this the color of the solution that you get when ammonia is added to cupric ion?" When you have the urge to ask such questions, try the experiment yourself with a single-cation solution and see the answer!

Finally, when you think that you know which of the six cations are present in your unknown (and, just as importantly, which are not) you should carefully repeat your analysis to confirm your results.

A Note on Net Reactions
The balanced net reaction describes the chemical changes of any reaction. It is balanced in the usual sense: as many moles of each element in all the reactants as in all the products, and a net charge for all reactants equal to that for all products. It is a net reaction in the sense that only those species in solution that actually change or participate in new bonding situations are included. For example, in the confirming test for Fe3+,  solid NH4SCN is added to a solution containing Fe3+. The net reaction, however, does not mention NH4SCN, since it dissociates in solution to the ammonium ion, NH4+, and the thiocyanate ion, SCN-. Only the thiocyanate ion is involved in the chemistry; the ammonium ion is a spectator. Thus, we write the balanced net reaction as

Fe3+ + 6 SCN-  Fe(SCN)63-
Writing net reactions thus requires you to be able to identify the reacting species in solution as well as the resultant new product, be it solid precipitate or, as in this case, a complex ion that stays in solution.

Some of the net reactions are redox reactions that will include H3O+ or OH- reactants (according to the pH of the solution). These may require more time and effort to balance. Always start with oxidation and reduction half-reactions. Balance these for both mass and charge. Then add the half-reactions together for the overall reaction. See your textbook or TA for help with these reactions, if you have difficulty.


Key QuestionsThe key questions to ask, as you perform this and next week's experiments are:
1. What ions are present in my test tube?
2. What ions are in solution?
3. What is the precipitate?
4. What do I expect to happen as I add the next reagent?
Prelab Problems
1. Describe the process of washing a precipitate. Why must this be done carefully and completely? How will you test the pH of the sample in your test tube? What are the steps for using a centrifuge?

2. How could you separate each of these pairs?:

 (i) Ag+(aq), Cu+2(aq)

 (ii) AgCl(s), CuC12(s)

 (iii) AgCl(s), PbC12(s)

 (iv) Ag+(aq), Fe2+(aq)

3. Construct a logical flowchart for the analysis of an unknown mixture of cations. Include balanced chemical equations and confirmation tests for each ion. Show clearly what is contained in the precipitate and centrifugate for each step of the analysis. A partial flowchart is given below, to help you get started.
Logical flowchart for analysis of unknown mixture of cations
4. Use the web applet Qualitative Analysis of Cations to test a web-based unknown sample. Practice with known samples and unknowns that you can reveal before trying a "true unknown". Follow the procedure and use your flowchart from prelab problem 3. In your lab notebook, record the code number of your "true unknown" and its identity for your TA to grade. Like all prelab problems, this one should be done independently, not in collaboration with a lab partner or friend. You can access the applet by clicking on the link below.
Qualitative Analysis of Cations Applet


Techniques
In this lab, you will perform simple test tube experiments and make careful observations. To separate ions of each group, you will form an insoluble precipitate and centrifuge to remove solid from the ions that remain in solution. You will need to carefullywash the precipitate and decant the solution to remove all soluble ions. You will use pH paper when you acidify and basify your samples. Finally, you will heat in a water bath to perform some reactions. Be sure to use clean glassware to avoid contamination of your sample. Wash glassware with soap and rinse with distilled water.

Each of the techniques listed above are described and illustrated on the ChemLab website. Use the information on the website to learn each boldfaced technique before coming to lab. Thetechnique information on the website will be helpful in answeringPrelab Problem 1.




Procedure
First start heating your water bath, on a hotplate. It should contain distilled water. The bath should be boiling gently when used.

You are strongly advised to retain the test tubes of all confirmation tests until you complete the experiment. This way, you can reinterpret your results, if necessary.

Unknown mixture of cations
Obtain your unknown mixture, which contains a subset of the six possible ions. Note its number in your notebook and note as well any simple physical characteristics of the mixture such as color, pH (from indicator paper), etc. Your task is to determine the constituents of the mixture. Use your flowchart and observations to confirm the presence or absence of each cation. Check your flowchart with your TA before beginning your analysis.

To confirm your analysis of the unknown mixture, use the known solutions of single cations to observe positive confirmation tests. Record an observation for each cation's confirmation test in your notebook, either from your unknown or from a known solution.

When you have completed the analysis of your unknown, repeat it with a new sample from your unknown vial, if time permits. This will confirm your initial conclusions.

1. Separation and Analysis of Group I Ions
The solubility products (Ksp) of silver chloride and lead chloride are 1.8 x 10-10 and 1.7 x 10-5, respectively. This means that neither salt is soluble and that AgCl is even less soluble than PbCl2. These two Group I ions (Ag+ and Pb2+) are separated from the other cations in the sample by addition of HCl, precipitation of the Group I metal chlorides, and separation of the solid precipitate from the remainder of the solution, which contains the dissolved Group II and III ions. Be careful to avoid a large excess of HCl, since both cations form soluble complex cations with excess chloride ions.

Add 6 drops of 6 M HCl to no more than 1 mL of the solution to be analyzed. If a precipitate forms, centrifuge the sample, and save the liquid solution for further analysis in Section 2.

Image 4A sample after centrifugation is shown. Make sure the solution is completely clear. Add an additional drop or two of HCl and let the solution sit, to see if more AgCl or PbCl2 precipitate forms. If so, remove by centrifugation. Carefully wash the precipitate and test it for lead and/or silver.

To make the most efficient use of time, you should begin the separation of Group II ions, in section 2 at this point. While that solution is heating in the water bath, you can return to your Group I ions and perform the confirmation tests for lead and silver. Careful labeling and notebook records will help you to keep track of your test tubes.

Image 5A. Confirmation of Lead (Pb2+)
Lead chloride is almost three times more soluble in hot water than cold. One may use this as a basis for separating it from silver chloride. The presence of lead is then confirmed by precipitation of yellow lead chromate.

Add ~20 drops of hot distilled water to the solid precipitate from above. Centrifuge while hot, decant, and save both the solution and the solid. Add 2-3 drops of 1 M K2CrO4 to the solution. A bright yellow precipitate confirms the presence of Pb2+.

B. Confirmation of Silver (Ag+)
Silver forms a soluble complex ion with aqueous ammonia. The presence of silver is confirmed by dissolving any remaining solid residue in 6 M NH3 (aq) and then re-precipitating the chloride by freeing the silver ion from the complex ion using 6 M acid.

Image 6Add ~5 drops of 6 M NH3 (aq) to the solid from A, keeping your test tube near the inlet of the fume exhaust vent. The solid should dissolve, but if any precipitate remains, centrifuge and proceed using only the centrifugate. Add 6 M HNO3 to the solution until the solution is acidified, using pH indicator paper to test for acidification. A white precipitate (AgCl) confirms the presence of Ag+. The Cl- needed for precipitation will be present from the prior dissolution of AgCl.

To perform this test on a known solution containing silver ions, add HCl to the solution to observe the AgCl precipitation.

2. Separation and Analysis of Group II Ions
The Group II ions are separated from those in Group III by adding hydrogen sulfide to the mixture in acidic solution. We use an organic precursor, thioacetamide (CH3CSNH2), which decomposes to hydrogen sulfide in the presence of acid according to the net reaction
CH3CSNH2 + 2 H2 CH3CO2- + NH4+ + H2S
The Group II ions precipitate as sulfides on reaction with S2-from H2S.

Image 7Add 2 drops of 6 M HCl to the solution from Section 1 and dilute to ~2.5 mL. Confirm acidity with pH paper. With the sample solution placed near the fume exhaust vent, add 10 drops 5% thioacetamide solution, stir, and heat in a water bath for at least 10 minutes. The precipitate contains copper sulfide. Centrifuge, decant, and save the centrifugate for Group III identification. Wash the precipitate by stirring it with ~10 drops of 1 M HCl, centrifuge, and combine the centrifugate with that from the previous centrifugation. It is essential to remove all solid group II ions, so they will not interfere with the group III analysis later on.

A. Confirmation of Copper (Cu2+)
The precipitated sulfides are redissolved by addition of conc. HNO3, producing free ions and elemental sulfur. Once the ions are redissolved, one can confirm the presence of the Group II ion, Cu2+. Addition of conc. ammonia (NH3 (aq) ) initially results in the precipitation of copper hydroxide. Excess aqueous ammonia redissolves the copper hydroxide via the complex ion, Cu(NH3)42+. A characteristic blue color confirms the presence of Cu2+ in solution.

Image 8Add ~8 drops of 16M HNO3to dissolve the precipitate which contains Cu2+. If large black "clumps" remain, centrifuge and discard the solid. It is probably PbS(s) or S(s). If solid floats, remove the solution to a clean test tube with a dropper, leaving the solid behind. Now add conc. NH3 (aq) until the solution is strongly basic according to pH paper. This could take as much as 20 drops. A blue solution indicates the presence of Cu2+ ions.

3. Analysis of Group III Ions
The Group III ions form sulfides that are more soluble than those of Group II, but they may be precipitated as sulfides if the sulfide concentration is sufficiently large. Addition of conc. NH3 (aq) to the thioacetamide insures that the necessary sulfide concentration will be reached.

Your cation solution should contain only Group III ions after the procedure outlined in Section 2 above. Keeping your sample near the fume exhaust vent, add ~5 drops of 5% thioacetamide solution, stire, and heat for ~5 minutes. Add ~5 drops of conc. NH3 (aq), stir and heat for 5 additional minutes. The precipitate contains the sulfides of the Group III ions. Centrifuge and discard the liquid.

A. Separation of Nickel (Ni2+) from Iron (Fe2+) and Manganese (Mn2+)
The sulfides of Fe2+ and Mn2+ are soluble in 1 M HCl, but nickel sulfide is not. This is used as the basis for separating Ni2+ from the remaining two ions.

Add ~10 drops of 1 M HCl to the precipitate prepared above, stir and centrifuge. Decant and wash any solid residue (NiS) with ~5 drops of 1 M HCl, and add the wash to the centrifugate. If present, Fe2+ and Mn2+ should now be in the solution. If solid floats, remove the solution to a clean test tube with a dropper, leaving the solid behind.

B. Confirmation of Ni2+Nickel sulfide is soluble in a mixture of nitric and hydrochloric acids that converts sulfide to elemental sulfur. In this reaction sulfur is oxidized from the -2 oxidation state (S2-) to the zero oxidation state (S(s)). At the same time, the nitrogen is reduced from a +5 oxidation state (NO3-) to a +2 oxidation state (NO). The free Ni2+ that results is first complexed with ammonia and then detected as an insoluble, scarlet coordination compound of dimethylglyoxime (DMGH2). The structures of DMGH2 and the coordination compound are shown below. Note that DMGH2 is weakly acidic and loses one of its two acidic protons in order to form an electrically neutral complex. This DMGH2conjugate base anion can be abbreviated DMGH- in your net reactions.
Figure 1
Image 9Add ~6 drops of conc. HCl and 2 drops of conc. HNO3 to any precipitate from A and heat the mixture. Add 6 M NH3 (aq) a few drops at a time. Stir with a clean stirring rod after each addition of NH3 (aq) and test the pH with indicator paper. Add NH3 (aq) until a strongly basic pH is reached. A total of 10-15 drops are typically required. This produces the Ni(NH3)62+ complex ion. Dilute to 1 mL with H2O. Add ~3 drops of dimethylglyoxime to the solution. Ni2+ is confirmed by the formation of a scarlet to strawberry-red precipitate.

C. Confirmation of Mn2+
Manganese is detected by oxidation of Mn2+ to permanganate ion (MnO4-) by sodium bismuthate (NaBiO3) in nitric acid. In this reaction, manganese is oxidized from the +2 oxidation state (Mn+2) to the +7 oxidation state (MnO4-). At the same time, bismuth is reduced from the +5 to the +3 oxidation state (Bi+3). You can tell that reaction has occurred from the distinctive purple or pink color of permanganate.

Image 10Add 5 drops of 6 M HNO3 and ~2 drops of 1 M sodium nitrite (NaNO2) to the solution from part A, and dilute to 1 mL. Heat the solution and, after cooling, divide it into 2 parts, setting aside one part for section D below. Add a spatula tip of NaBiO3 and ~2-3 drops or more of 6 M HNO3. The formation of a transient pink to purple color of MnO4- confirms Mn2+. An additional spatula tip of solid and a gentle shake of the test tube may make this subtle color change more apparent. Bismuthate (BiO3-) in strong acid is predominantly molecular bismuthic acid (HBiO3) and is reduced to Bi3+.

Image 11D. Confirmation of Fe2+/Fe3+
The confirmation test for Fe3+ is simple. Treating the solution with nitric acid in step C, above, oxidizes Fe2+to Fe3+. The latter ion reacts with thiocyanate ion to produce a well-known "blood red" hexathiocyanato complex ion.

To the solution set aside in part C add 2-3 crystals of NH4SCN. Very little solid is required for the solution to turn a dark blood red, if Fe3+ is present.
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