Cation Separation 6dc46
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Chemistry 102L Manual - Page 153
Separation and Identification of Cations Introduction Metals are a very important part of our society, and the majority of these are not found in the metallic form in nature. The more useful metals (including copper, nickel, zinc, etc.) occur as positively charged ions (cations) incorporated into minerals. The cations of the more useful metals are often mixed with an array of less valuable cations. Separation of the cations of the desired metal from the others is a critical step in metal production. This experiment will explore some of the strategies used to separate (and identify) metal cations. The flowchart shown below is a classic method for separating complex mixtures of metal cations. Its central strategy is to first separate the cations into smaller “classes” that can be then chemically sorted into the individual metals. This method was used to obtain specific metals from low-grade “messy” ores. Though purer higher-grade ores are used in modern practice, some of these old chemical separation “tricks” are still used to remove unavoidable contaminants from the desired metals. The major steps of this flow chart are explained below. Cation Mixture HCl (aq)
(s) Class I Cations
H2S pH=0.5
(s)
(aq)
Class II Cations (s) Class IIa Cations
H2S pH=8
OH(aq) Class IIb Cations
(aq)
(s) Class III Cations
3 M (NH4)2CO3 (s)
Class IV Cations
(aq) Class V Cations
Separation of Class I cations (Ag+,Pb2+,Hg22+) In the first step of this flow chart, the Class I cations (Ag +, Hg22+, Pb2+) are separated from the others by virtue of their unique ability to form insoluble chlorides. M2+(aq) + 2Cl-(aq)
MCl2(s)
Chemistry 102L Manual - Page 154
Upon the addition of HCl(aq) [which is really H+(aq) + Cl-(aq)] Ag +(aq), Hg 22+(aq), Pb 2+(aq) ions will precipitate as AgCl(s), Hg2Cl2(s), PbCl2(s) respectively while all the other ions will remain in solution (since their chlorides are soluble). that the appearance of such a solid merely indicates that one or more of the “Class I” cations Ag+(aq), Hg22+(aq), Pb2+(aq) were present in the original solution. After the solid is separated from the solution, the specific cations in the solid can now be easily identified through simple tests. For example, only solid AgCl(s) will dissolve in the presence of NH3 to form Ag(NH3)2+ and Cl-. AgCl(s) + 2NH3(aq) ––––> Ag(NH3)2+(aq) + Cl-(aq) In contrast, Hg2Cl2(s) would react with NH3 to give a black solid. This black solid would be positive evidence that the original solution contained Hg2+(aq). Hg2Cl2(s) + 2NH3 ––––> HgNH2Cl(s) + Hg(l) + NH4+ + ClThe addition of NH3 to PbCl2(s) would have no effect. Separation of Class II (Hg2+,Pb2+,Bi3+,Cu2+,Cd2+,Sb3+,Sn4+) and Class III cations (Co2+,Mn2+,Zn2+,Ni2+,Fe2+,Al3+,Cr3+) The next two steps in the flow chart separate Class II ions and Class III ions from the remaining solution by their ability to form insoluble sulfides upon reaction with H2S. M2+ + H2S
MS(s) + 2H+
This equilibrium is really a combination of three standard equilibria: H2S
HS- + H+
HS-
S2- + H+
M2+ + S2-
Ka1 = 8.9·10-8 Ka2 = 1.2·10-13
MS(s)
1/Ksp
When these are combined (added together), the resulting equilibrium constant is strongly dependent on the Ksp of the particular sulfide (MS). M2+ + H2S
MS(s) + 2H+
K = Ka1Ka2 / Ksp K = (1.07·10-20) / Ksp
The Ksp’s of Class II and Class III cations are all relatively small. This makes the overall equilibrium constant (K) large which favors the formation of solid MS(s). Nevertheless, the Ksp’s for the different cations vary from small to extremely small. We can take advantage of these differences to separate the two classes of cations.
Chemistry 102L Manual - Page 155
Class II cations have very small Ksp’s and very large overall equilibrium constants for the reaction with H2S. This favors the formation of solid MS(s) even when the concentration of H+ is high (pH = 0.5). Class III cations have small Ksp’s but not as small as those of Class II Cations. The resulting overall equilibrium constants are not as large and favor the formation of solid MS(s) only when the concentration of H+ is small (pH = 8). This solubility behavior allows Class II cations to be precipitated first with H2S at pH=0.5 and Class III cations to be precipitated later with H2S at pH=8. that a precipitate with H2S at pH=0.5 only tells you that the original sample solution contained one or more of the Class II ions. Further tests are needed to determine what particular ions were present. Similarly, a precipitate with H2S at pH=8 only tells you that the original sample solution contained one or more of the Class III ions. Further tests are needed. Use of H2S to precipitate these cations can sometimes accidentally produce solid sulfur. Don’t mistake this solid sulfur for a metal sulfide! Sulfur can accidentally form two ways: At pH=0.5, sulfur can form in the presence of nitrates (NO3-). Because of this, nitrates and nitric acid should be avoided (if possible) when doing cation analysis. 2H+ + 2NO3- + 3H2S ––––> 3S(s) + 2NO + 4H2O Sulfur can also form in the presence of Fe3+. Fe3+ forms from the reaction of Fe2+ (a Class III ion) with air. Though no one will receive Fe2+ as a cation in this laboratory, you should but be aware of the problem. 2Fe3+ + H2S ––––> 2Fe2+ + S(s) + 2H+ Sulfur appears as a white to yellow precipitate than can make the solution look smoky. If a small amount of such a solid appears, test the solid by adding Na2S. This will make the solid sulfur disappear by the following reaction. S(s) + S2-(aq)
S22-(aq)
Chemistry 102L Manual - Page 156
Separation of Class IIa (Hg2+,Pb2+,Bi3+,Cu2+,Cd2+) and IIb (Sn4+,Sb3+) cations Class II is large enough of a group of cations that another separation is necessary. The sulfides of the Class IIb cations (Sb2S3 and SnS2) can be separated from those of the Class IIa ions by their tendency to react with hydroxide to form soluble hydroxides Sb(OH)4- and Sn(OH)62-. Since the other Class II sulfides (the “Class IIa” sulfides) do not dissolve in base, the Class IIb hydroxides are easily separated for further testing. Sb2S3(s) + 8OH-(aq) SnS2(s) + 6OH-(aq)
2Sb(OH)4-(aq) + 3S2-(aq) Sn(OH)62-(aq) + S2-(aq)
Separation of Class IV Cations (Ca2+,Ba2+) Among the cations not yet precipitated with HCl or H2S, the Class IV cations (Ca2+ and Ba 2+) precipitate with the addition of 3M (NH4)2CO3. Ca2+ + CO32-
CaCO3(s)
Ba2+ + CO32-
BaCO3(s)
Since the remaining Class V ions (Mg2+, Na +, K +) form soluble salts with CO32-, they remain in solution. This makes them easy to separate from the solid Class IV carbonates: CaCO3 and BaCO3. Identification of Class V Cations (Mg2+,Na+,K+) These ions are the leftovers from the HCl, H2S, and (NH4)2CO3 precipitations. Of the Class V cations, only Mg2+ precipitates with the addition of hydroxide to form Mg(OH)2. Mg2+ + 2OH-
Mg(OH)2(s)
Na+ and K+ do not precipitate with any of the common anions. However, they are easily distinguishable by the distinctive colors they give to a flame when their solutions are burned: Na+ gives a bright yellow flame, and K+ gives a violet flame. Useful clues: Cation separation can be an arduous task. Simple clues can offer helpful insight into the presence or absence of specific cations, and this information could be used to simplify the analysis. Identification of Cations with Colors Most cations form colorless ions in water, but a few (Cr3+, Fe3+, Co2+, Ni2+, Cu2+) form colored ions in water. The appearance of one or more of these colors can be a very useful clue to the presence of these ions. We will see a few of these in this experiment.
Chemistry 102L Manual - Page 157
Identification of Cations by Flame Tests Some cations glow a distinctive color when they are caught in a flame. These colors offer yet another easy clue to the presence of particular ions.
Experimental Objective. The purpose of this experiment is to give you that opportunity to use the techniques discussed above to separate cations by their Classes and to use specific tests to distinguish specific ions. In the next couple experiments, you will be using these techniques to determine what ions are present in a given water solution. Some of the chemical tests require you to recognize the color and texture of solids, so pay particular attention to how these reactions look to you.
Important Notes • This is a two session experiment, so tightly stopper solutions and solids you want to save. • You may work in groups of 2 (but no larger without explicit permission from your instructor). • The report form at the end of the procedure and the questions must be completed and returned to your instructor one week after completing the experiment. • Though you are not expected to copy the detailed procedure in to your notebook, you are expected to summarize the procedure in your own words. Include important details such as about reagents and indicators you used. This will be important for notebook quizzes. No photocopies are allowed. Include spaces for you to write results. • You are also expected to record all results in your notebook as you collect them. Your instructor will be spot checking your work.
Experimental Procedure. Part A. Separation of Class I Cations: • Into test tube labeled “Ag+” add: • 5 drops of 0.1 M AgNO3 • 1 mL of water • drop-wise: 10 to 15 drops 6M H Cl (until no more white curd-like solid forms). • Into test tube labeled “Pb2+” add: • 5 drops of 0.1 M Pb(NO3)2 • 1 mL of water • drop-wise: 10 to 15 drops 6M H Cl (until no more granular white solid forms).
Chemistry 102L Manual - Page 158
• Into another test tube: • add a small pea sized portion of powdered Hg2(NO3)2. • add a 1 mL of water. • Heat the Hg22+ solution to a near boil. • Cool the Hg22+ solution and transfer it to a centrifuge tube. • Centrifuge the Hg22+ suspension and decant the solution into a test tube labeled “Hg22+”. • Into the “Hg22+” tube, add 10 to 15 drops of 6M HCl (until no more white precipitate forms). • Heat the “Ag+”, “Pb2+”, and “Hg22+” tubes in boiling water, and occasionally stir the contents of each test tube as it heats. •) Which of the solids dissolve? • Cool the tubes in ice water. •) Does the solid that dissolved above reform? > In principle, we can separate this ion from the other Class I cations by heating the chloride precipitate in hot water, centrifuging while hot, and then cooling the solution to get the chloride of that ion. If the solid indeed reforms, we have confirmed that that ion was in the original mixture. • With the chlorides that were insoluble in hot water: • Centrifuge the suspensions and discard the solutions. • Wash the solids with water. Reminder: “washing the solid” means: • Add water to the solid, and resuspend it with a stir rod. • Centrifuge the suspension, and discard the solution. • Add 2 mL of 6M NH3, and mix the suspension well. •) What happened in the two tubes: (Identify the cation and any changes in the solid and solution.)
> The solids in one tube should dissolve. For a mixture, we can now centrifuge the suspension and transfer the solution containing that ion into another tube leaving the other ion behind.
Chemistry 102L Manual - Page 159
> The solids in the other tube should change color. This is the test that confirms the presence of that ion. • If the solid dissolves in the NH3: • Add 1 M KI drop by drop until something happens. •) What happens? > The reaction you see is the test that confirms the presence of that ion in your solution. • Summarize your results above as a flow chart that allows you to separate the Class I cations and confirm their presence. Part B. Precipitation of Class II and Class III cations. • Into a test tube labeled “Cu2+, Zn2+” add: • 5 drops of 0.5 M Cu2+ • 5 drops of 0.2 M Zn2+ • 1 mL of water • 1 mL of thioacetamide (in the hood). • drop-wise 6 M HCl with stirring until the solution makes methyl violet paper turn blue-green. (Put a drop of 0.3 M HCl onto methyl violet paper for comparison.) • If the paper is purple or too blue, add more acid. • If the paper is yellow, add 6M NH3(aq). • Into a test tube labeled “Cu2+, Al3+” add: • 5 drops of 0.5 M Cu2+ • 5 drops of 0.2 M Al3+ • 1 mL of water • 1 mL of thioacetamide (in the hood). • drop-wise 6 M HCl with stirring until the solution makes methyl violet paper turn blue-green. (Put a drop of 0.3 M HCl onto methyl violet paper for comparison.) • If the paper is purple or too blue, add more acid. • If the paper is yellow, add 6 M NH3.
Chemistry 102L Manual - Page 160
• Heat these two tubes in a 70°C water bath (in the hood) for 10 minutes. •) Describe the precipitate in the “Cu2+, Zn2+” tube:
•) Describe the precipitate in the “Cu2+, Al3+” tube:
> The precipitates in both tubes should be CuS(s). • Centrifuge the precipitate in the two tubes and transfer the solutions into new test tubes with new “Cu2+, Zn2+” and “Cu2+, Al3+” labels. > Here, we have largely separated the Cu2+ from these solutions as CuS(s). • Add 6 M HCl drop-wise to both test tubes to adjust [H+] to 0.3 M as before. • Add 6 M NH3 drop-wise to both tubes until they are pH = 8. (Test this by touching a drop of the solution on a stir rod onto universal indicator paper.) • Add 1 mL of thioacetamide to both tubes and then heat them in the 70°C water bath for 10 minutes. • Separately, centrifuge the contents of both tubes and discard the solutions. • Wash the solids in both tubes twice with 2 or 3 mL of water. > The solid from the “Cu2+, Zn2+” tube is ZnS(s). > The solid from the “Cu2+, Al3+” tube is Al(OH)3. •) Is either solid colored?
> This color is from residual Cu2+ that was left over from the first separation. Since CuS is so darkly colored, only a little of it gives the color you see. This residual Cu2+ is nothing to worry about in a separation. Just be aware that it is there. • Add 5 drops of 6 M H Cl to these solids. •) Describe the odor from the “Cu2+, Zn2+” tube.
•) Describe the odor from the “Cu2+, Al3+” tube.
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•) Can you tell these tube apart by their odors. If so, how?
> Any difference you notice could be used to distinguish Zn2+ and Al3+ during an analysis. Part C. Sulfide Amphoterism: Separations of Class IIa and Class IIb. • Into a test tube labeled as “Bi3+”: • Add 1 mL water • 5 drops of 0.1 M Bi3+. • Into a test tube labeled as “Sn4+”: • Add 1 mL water • 5 drops of 0.1 M Sn4+. • Add 6 M HCl to both tubes until their solutions turn methyl violet paper blue-green. • Add 1 mL of thioacetamide to both tubes, and then heat them in the 70°C water bath (in the hood) for 5 minutes. • Separately centrifuge the suspensions from Bi3+ and Sn4+, and wash the solids with water. •) Describe the precipitate in the “Bi3+” tube.
> This solid should be Bi2S3(s). •) Describe the precipitate in the “Sn4+” tube.
> This solid should be SnS2(s). • With the solid from the “Sn4+” tube: • Add 6 M NaOH to the solid until it completely dissolves. > This is the reaction: SnS2(s) + 4OH-
Sn(OH)4- + 2S2-
• Add 6 M HCl to this Sn(OH)4- solution until SnS2(s) reprecipitates.
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• With the solid from the “Bi3+” tube: • Add about the same amount of 6 M NaOH to the solid, and warm the tube. •) Does much of the solid dissolve? • Centrifuge the contents of the “Bi3+” tube, and decant the solution into another tube. • Add 6 M H Cl to this solution until it is acidic. •) Does any solid form? •) If so, how does the amount compare to that from the “Sn4+” case above?
> Overall, though both Sn4+ and Bi3+ are Class II ions, they can be further separated by their solubilities in NaOH: the SnS2 dissolves while Bi2S3 does not. This allows the Sn(OH)4to be separated by centrifugation (as the solution), and then reprecipitated as the SnS2 by reacidification. The end result: Cation IIb ions (such as SnS2) can be separated from Cation IIa ions (such as Bi2S3). Part D. Separation within Class IV. • Into the first tube add: • 1 mL of water. • 5 drops of 0.1 M Mg2+. • 1 mL of 6 M NaOH •) Centrifuge the solution/suspension. Did a precipitate form? If so, describe it. Hint: look for a gel formed on the test tube wall. If it is not there, cool the tube in ice water for 5 minutes and then centrifuge a second time. Then look again.
• Save the tube for comparisons. • Into the second tube add: • 1 mL of water. • 5 drops of 0.1 M Mg2+. • 1 mL of 6 M NH3 •) Did a precipitate form? If so, describe it. (Note: a precipitate may not form.)
• If a solid formed, centrifuge and save it for comparisons.
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• Into the third tube add: • 1 mL of water. • 5 drops of 0.1 M Mg2+ • 0.5 mL of 3 M (NH4)2CO3 • 1 mL of 6 M NH3. •) Did a precipitate form? If so, describe it. (Note: a precipitate may not form.)
• If a solid formed, centrifuge and save it for comparisons. •) Compare the amount of solid in the three test tubes.
> The first, second, and third test tubes are in a progression of decreasing basicity. All three are capable of precipitating Ca2+ and Ba2+. We would like to precipitate these two ions but not Mg2+. Which procedure of the three above would I use to do this?
> After the Mg2+ is removed from the solid CaCO3 and BaCO3, we can then precipitate it as Mg(OH)2 with NaOH. Part E. Precipitation of Sulfur: How to identify it and what to do about it. • To two test tubes add: • 1 mL 6 M HNO3. • 1 mL thioacetamide. • Heat these tubes in a 70°C water bath (in the hood) for 10 to 15 minutes. • Centrifuge the sulfur precipitate, and wash it with water. Note: only a very small amount of solid may be there. • To one centrifuge tube, add 1 mL concentrated HCl. • To the other centrifuge tube, add 1 mL of 1.0 M Na2S.
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• Heat the tubes in a 70°C water bath for 5 minutes. •) Does the solid in the HCl dissolve?
•) Does the solid in the Na2S dissolve?
> We will be using the successful procedure to remove precipitate sulfur from sulfides during our analyses. This will prevent us from getting tricked into thinking we have precipitated sulfides when we really don’t. Part F. Using Solution Color for Cation Identification. The following solutions should be available in bottles on a tray in the laboratory. Just examine them to answer the following questions. •) What color is 0.1 M Ni(NO3)2? •) What color is 0.1 M Co(NO3)2? •) What color is 0.1 M Fe(NO3)3? •) What color is 0.1 M Cr(NO3)3? •) What color is 0.5 M Cu(NO3)2? Part G. Using Flame Tests for Cation Identification. • Find a clean ceramic spot plate and a clean Nichrome wire loop. • Light a Bunsen burner in a hood. As your instructor for help if needed. • Add 5 drops of each the following solutions into their own wells of the spot plate. (One cation per well.) 0.10 M Ca(NO3)2 0.10 M Ba(NO3)2 0.10 M Cu(NO3)2 0.10 M NaNO3 0.10 M KNO3 • Add 1 drop of 12 M HCl into each of these five wells. • Add 5 drops of 12 M HCl into a sixth well.
Chemistry 102L Manual - Page 165
Flame clean the Nichrome wire loop. Repeat the following three times. • Dip the Nichrome wire loop into the 12 M HCl (in the sixth well), and lift the loop out of the HCl such that a drop of the acid is hanging inside the loop. • Hold the loop and its drop of HCl in the middle of the Bunsen burner flame. You may see the flame quickly change color as the drop of HCl evaporates. This quick color change is due to cation residues on the wire (probably from the previous ). The flame will also likely begin to glow orange as the flame makes the Nichrome wire glow. This late color change (if it appears) is due to unavoidable contamination of the Nichrome and should be ignored during the following flame tests. For each of the cation/HCl solutions you prepared above, follow the following steps. (These are the flame tests). • Dip the Nichrome wire loop into a cation / HCl solution, and lift the loop out of the solution such that a drop of the solution is hanging inside the loop. • Hold the loop and its drop in the middle of the Bunsen burner flame. Watch for a quick color glow in the flame. This color should appear almost immediately. Don’t confuse this with the later orange color that appears when the wire begins glowing. Describe the flame color that quickly appears from each of the cation solutions: •) What color appears from Ca2+? •) What color appears from Ba2+? •) What color appears from Cu2+? •) What color appears from Na+? •) What color appears from K+? • Flame clean the Nichrome wire loop once as described above. In mixtures of Na+ and other cations, the strong color of Na+ could obstruct the color of others. To see these hidden colors, we can simply filter out the glow from Na+ with cobalt glass. • Find the dark purple glass squares in your equipment drawer. (You only need one.) • Add 1 drop of 0.10 M NaNO3 to the cation / HCl solutions above.
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Repeat the following for each of the Na+ / cation / HCl solutions. • Dip the Nichrome wire loop into a cation / HCl solution, and lift the loop out of the solution such that a drop of the solution is hanging inside the loop. • Hold the loop and its drop in the middle of the Bunsen burner flame as you watch through the cobalt glass. (Keep your goggles on!) Watch for a nearly immediate color glow in the flame. Again, don’t confuse this with the later orange color that appears when the wire begins glowing. Describe the flame color that quickly appears from each of the cation solutions: •) What color appears from the solution with only Na+? •) What color appears from Ca2+ and Na+? •) What color appears from Ba2+ and Na+? •) What color appears from Cu2+ and Na+? •) What color appears from K+ and Na+? • Flame clean the Nichrome wire loop once as described above.
Chemistry 102L Manual - Page 167
Report Form: SEPARATION AND IDENTIFICATION OF SELECTED CATIONS Name
Date
Laboratory Instructor
Section
A. Separation and Identification of Ag+, Hg22+, and Pb2+ 1. a. Which chloride (AgCl, Hg2Cl2 or PbCl2) dissolved in hot water? __________ b. Did precipitate form when the hot tube was cooled? __________ 2. a. Which chloride gave a black precipitate with 6 M NH3? __________ b. Net Reaction:
3. a. Which chloride dissolved in 6 M NH3? __________ b. Net Reaction:
c. What compound precipitated when KI was added? __________ d. Net Reaction:
4. Using the reactions described above, draw a flow chart that summarizes how to separate and identify AgCl, PbCl2 and Hg2Cl2 in a mixture.
Chemistry 102L Manual - Page 168
B. Precipitation of Sulfides - Separation of Class II and III 1. a. What was the color of the precipitate produced when a solution of Zn2+, Cu2+ reacted with H2S at 0.3 M H+? b. What was the formula for the precipitate?
2. a. What was the color of the ZnS precipitated at pH 8?
b. What was the color of the Al(OH)3 precipitated at pH 8? c. As the H+ decreases, does the concentration of S2- increase or decrease? As this happens, would you expect more CuS to form if some Cu2+ were still in solution?
Based on this, would you expect more CuS to form?
d. What happened when 6M HCl was added to Al(OH)3(s) and ZnS(s)?
e. Could you distinguish ZnS(s) and Al(OH)3(s) by odor when you added acid? How?
3. a. Given that Bi2S3 is black and SnS2 is yellow, which cation (Bi3+ or Sn4+) precipitated as a sulfide with thioacetamide at pH 8?
b. Which cation (Bi 3+ or Sn4+) would not form a sulfide precipitate at pH 8 but would when the H+ concentration is increased to 0.3 M H+.
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c. Given that the equilibrium constant K for hydrolysis of Sn4+ to Sn(OH)62- is large: Sn4+(aq) + 6H2O(l)
Sn(OH)62-(aq) + 6H+(aq)
that this is the reaction that couples to the reaction below to make SnS2(s) dissolve: SnS2(s)
Sn4+(aq) + 2S2-(aq)
i.
Considering the first equilibrium, what would happen if the pH decreased? Specifically, would the Sn4+ concentration increase or decrease? Explain.
ii.
How would the change in Sn4+ concentration effect the second equilibrium? Specifically, what would happen to the amount of SnS2(s)? Explain.
iii.
What did you see happen when KOH was added to SnS2? Does this agree with the behavior you predicted above?
C. Sulfide Amphoterism - Separations within Class II 1. Was Bi2S3 insoluble in the KOH solution? 2. Based on your tests, is KOH adequate (as it is used in the procedure) for separating Bi2S3 and SnS2? If not, explain.
D. Separations within Class IV 1. Did you get a precipitate when you added 1 mL of 6 M NH3 to the Mg2+ solution which didn't contain (NH4)2CO3? 2. Was the amount of precipitate similar to what you got adding 1 mL of 6 M NaOH?
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3. Did you get a precipitate when you added 1 mL of 6 M NH3 to the solution containing (NH4)2CO3 and Mg2+? If so, how did the amount compare with that you got without the (NH4)2CO3?
4. Based on your results, would it be wise to change the flow chart procedure and use 6 M NH3 in place of 6 M NaOH for precipitating Mg2+ as Mg(OH)2? Why?
E. Precipitation of Sulfur 1. Did sulfur precipitate rapidly or slowly?
2. a. Was the sulfur soluble in concentrated HCl?
b. Was sulfur soluble in 1 M Na2S?
F. Identification of Cations by Color Complete the table below with the colors you observed. Cation Ni2+ Co2+ Fe3+ Cr3+ Cu2+
Cation Color
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G. Identifications of Cations by Flame Color Complete the table below with the colors you observed (without the cobalt glass). Cation
Flame Color
Ca2+ Ba2+ Cu2+ Na+ K+ Complete the table below with the colors you observed with the cobalt glass. Cation + Na+ Na+ Ca2+ + Na+ Ba2+ + Na+ Cu2+ + Na+ K+ + Na+
Flame Color (through cobalt glass)
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Questions
1. The following observations were made for an unknown solution that possibly contained the ions Hg22+, Ag+, Cu2+, Al3+, Ba2+, Na+, Ca2+, K+, Mg2+. For each observation, briefly interpret its meaning (what ions are present or absent). • The solution began colorless.
• Addition of HCl gave a white precipitate.
• This white precipitate was not soluble in hot water.
• With the addition of NH3, the white precipitate turned black and none dissolved.
• When the solution, from which the white precipitate was separated, was heated with thioacetamide (to generate H2S), no precipitate formed at H+ concentrations of 0.3 M. • When 1 mL of 6M NH3 was added to a portion of this solution no precipitate formed.
at pH 8,
• When NH3 was added to another portion of the pH 8 solution (until its pH was 9) and the resulting solution was mixed with 3M (NH4)2CO3, a white precipitate formed.
• The white precipitate dissolved with fizzing in concentrated HCl, and this HCl solution gave a green flame test.
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2. Suppose we were adding concentrated HCl to 1.0·10-3 M Pb2+(aq). What minimum concentration of HCl would be needed to make PbCl2 begin precipitating? [Ksp(PbCl2) = 1.6·10-5]. Show your work. [Hint: this would be the concentration of Cl- in equilibrium with PbCl2(s) and 1.0·10-3 M Pb2+(aq).]
3. a. If a 0.05 M Mn2+ solution were saturated with H2S (0.10 M H2S) and the pH was adjusted to 0.50, could any MnS precipitate? [Ksp(MnS) = 2.5·10-10]. Show your work. Hint: In the introduction, we found that: M2+ + H2S MS(s) + 2H+
K = (1.07·10-20) / Ksp
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b. If a 0.05 M Mn2+ solution were saturated with H2S (0.10 M H2S) and the pH was adjusted to 8.0, would any MnS precipitate? Show your work..
c.Given the behavior in parts (a) and (b), what class of cations would Mn 2+ belong to and why?
Separation and Identification of Cations Introduction Metals are a very important part of our society, and the majority of these are not found in the metallic form in nature. The more useful metals (including copper, nickel, zinc, etc.) occur as positively charged ions (cations) incorporated into minerals. The cations of the more useful metals are often mixed with an array of less valuable cations. Separation of the cations of the desired metal from the others is a critical step in metal production. This experiment will explore some of the strategies used to separate (and identify) metal cations. The flowchart shown below is a classic method for separating complex mixtures of metal cations. Its central strategy is to first separate the cations into smaller “classes” that can be then chemically sorted into the individual metals. This method was used to obtain specific metals from low-grade “messy” ores. Though purer higher-grade ores are used in modern practice, some of these old chemical separation “tricks” are still used to remove unavoidable contaminants from the desired metals. The major steps of this flow chart are explained below. Cation Mixture HCl (aq)
(s) Class I Cations
H2S pH=0.5
(s)
(aq)
Class II Cations (s) Class IIa Cations
H2S pH=8
OH(aq) Class IIb Cations
(aq)
(s) Class III Cations
3 M (NH4)2CO3 (s)
Class IV Cations
(aq) Class V Cations
Separation of Class I cations (Ag+,Pb2+,Hg22+) In the first step of this flow chart, the Class I cations (Ag +, Hg22+, Pb2+) are separated from the others by virtue of their unique ability to form insoluble chlorides. M2+(aq) + 2Cl-(aq)
MCl2(s)
Chemistry 102L Manual - Page 154
Upon the addition of HCl(aq) [which is really H+(aq) + Cl-(aq)] Ag +(aq), Hg 22+(aq), Pb 2+(aq) ions will precipitate as AgCl(s), Hg2Cl2(s), PbCl2(s) respectively while all the other ions will remain in solution (since their chlorides are soluble). that the appearance of such a solid merely indicates that one or more of the “Class I” cations Ag+(aq), Hg22+(aq), Pb2+(aq) were present in the original solution. After the solid is separated from the solution, the specific cations in the solid can now be easily identified through simple tests. For example, only solid AgCl(s) will dissolve in the presence of NH3 to form Ag(NH3)2+ and Cl-. AgCl(s) + 2NH3(aq) ––––> Ag(NH3)2+(aq) + Cl-(aq) In contrast, Hg2Cl2(s) would react with NH3 to give a black solid. This black solid would be positive evidence that the original solution contained Hg2+(aq). Hg2Cl2(s) + 2NH3 ––––> HgNH2Cl(s) + Hg(l) + NH4+ + ClThe addition of NH3 to PbCl2(s) would have no effect. Separation of Class II (Hg2+,Pb2+,Bi3+,Cu2+,Cd2+,Sb3+,Sn4+) and Class III cations (Co2+,Mn2+,Zn2+,Ni2+,Fe2+,Al3+,Cr3+) The next two steps in the flow chart separate Class II ions and Class III ions from the remaining solution by their ability to form insoluble sulfides upon reaction with H2S. M2+ + H2S
MS(s) + 2H+
This equilibrium is really a combination of three standard equilibria: H2S
HS- + H+
HS-
S2- + H+
M2+ + S2-
Ka1 = 8.9·10-8 Ka2 = 1.2·10-13
MS(s)
1/Ksp
When these are combined (added together), the resulting equilibrium constant is strongly dependent on the Ksp of the particular sulfide (MS). M2+ + H2S
MS(s) + 2H+
K = Ka1Ka2 / Ksp K = (1.07·10-20) / Ksp
The Ksp’s of Class II and Class III cations are all relatively small. This makes the overall equilibrium constant (K) large which favors the formation of solid MS(s). Nevertheless, the Ksp’s for the different cations vary from small to extremely small. We can take advantage of these differences to separate the two classes of cations.
Chemistry 102L Manual - Page 155
Class II cations have very small Ksp’s and very large overall equilibrium constants for the reaction with H2S. This favors the formation of solid MS(s) even when the concentration of H+ is high (pH = 0.5). Class III cations have small Ksp’s but not as small as those of Class II Cations. The resulting overall equilibrium constants are not as large and favor the formation of solid MS(s) only when the concentration of H+ is small (pH = 8). This solubility behavior allows Class II cations to be precipitated first with H2S at pH=0.5 and Class III cations to be precipitated later with H2S at pH=8. that a precipitate with H2S at pH=0.5 only tells you that the original sample solution contained one or more of the Class II ions. Further tests are needed to determine what particular ions were present. Similarly, a precipitate with H2S at pH=8 only tells you that the original sample solution contained one or more of the Class III ions. Further tests are needed. Use of H2S to precipitate these cations can sometimes accidentally produce solid sulfur. Don’t mistake this solid sulfur for a metal sulfide! Sulfur can accidentally form two ways: At pH=0.5, sulfur can form in the presence of nitrates (NO3-). Because of this, nitrates and nitric acid should be avoided (if possible) when doing cation analysis. 2H+ + 2NO3- + 3H2S ––––> 3S(s) + 2NO + 4H2O Sulfur can also form in the presence of Fe3+. Fe3+ forms from the reaction of Fe2+ (a Class III ion) with air. Though no one will receive Fe2+ as a cation in this laboratory, you should but be aware of the problem. 2Fe3+ + H2S ––––> 2Fe2+ + S(s) + 2H+ Sulfur appears as a white to yellow precipitate than can make the solution look smoky. If a small amount of such a solid appears, test the solid by adding Na2S. This will make the solid sulfur disappear by the following reaction. S(s) + S2-(aq)
S22-(aq)
Chemistry 102L Manual - Page 156
Separation of Class IIa (Hg2+,Pb2+,Bi3+,Cu2+,Cd2+) and IIb (Sn4+,Sb3+) cations Class II is large enough of a group of cations that another separation is necessary. The sulfides of the Class IIb cations (Sb2S3 and SnS2) can be separated from those of the Class IIa ions by their tendency to react with hydroxide to form soluble hydroxides Sb(OH)4- and Sn(OH)62-. Since the other Class II sulfides (the “Class IIa” sulfides) do not dissolve in base, the Class IIb hydroxides are easily separated for further testing. Sb2S3(s) + 8OH-(aq) SnS2(s) + 6OH-(aq)
2Sb(OH)4-(aq) + 3S2-(aq) Sn(OH)62-(aq) + S2-(aq)
Separation of Class IV Cations (Ca2+,Ba2+) Among the cations not yet precipitated with HCl or H2S, the Class IV cations (Ca2+ and Ba 2+) precipitate with the addition of 3M (NH4)2CO3. Ca2+ + CO32-
CaCO3(s)
Ba2+ + CO32-
BaCO3(s)
Since the remaining Class V ions (Mg2+, Na +, K +) form soluble salts with CO32-, they remain in solution. This makes them easy to separate from the solid Class IV carbonates: CaCO3 and BaCO3. Identification of Class V Cations (Mg2+,Na+,K+) These ions are the leftovers from the HCl, H2S, and (NH4)2CO3 precipitations. Of the Class V cations, only Mg2+ precipitates with the addition of hydroxide to form Mg(OH)2. Mg2+ + 2OH-
Mg(OH)2(s)
Na+ and K+ do not precipitate with any of the common anions. However, they are easily distinguishable by the distinctive colors they give to a flame when their solutions are burned: Na+ gives a bright yellow flame, and K+ gives a violet flame. Useful clues: Cation separation can be an arduous task. Simple clues can offer helpful insight into the presence or absence of specific cations, and this information could be used to simplify the analysis. Identification of Cations with Colors Most cations form colorless ions in water, but a few (Cr3+, Fe3+, Co2+, Ni2+, Cu2+) form colored ions in water. The appearance of one or more of these colors can be a very useful clue to the presence of these ions. We will see a few of these in this experiment.
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Identification of Cations by Flame Tests Some cations glow a distinctive color when they are caught in a flame. These colors offer yet another easy clue to the presence of particular ions.
Experimental Objective. The purpose of this experiment is to give you that opportunity to use the techniques discussed above to separate cations by their Classes and to use specific tests to distinguish specific ions. In the next couple experiments, you will be using these techniques to determine what ions are present in a given water solution. Some of the chemical tests require you to recognize the color and texture of solids, so pay particular attention to how these reactions look to you.
Important Notes • This is a two session experiment, so tightly stopper solutions and solids you want to save. • You may work in groups of 2 (but no larger without explicit permission from your instructor). • The report form at the end of the procedure and the questions must be completed and returned to your instructor one week after completing the experiment. • Though you are not expected to copy the detailed procedure in to your notebook, you are expected to summarize the procedure in your own words. Include important details such as about reagents and indicators you used. This will be important for notebook quizzes. No photocopies are allowed. Include spaces for you to write results. • You are also expected to record all results in your notebook as you collect them. Your instructor will be spot checking your work.
Experimental Procedure. Part A. Separation of Class I Cations: • Into test tube labeled “Ag+” add: • 5 drops of 0.1 M AgNO3 • 1 mL of water • drop-wise: 10 to 15 drops 6M H Cl (until no more white curd-like solid forms). • Into test tube labeled “Pb2+” add: • 5 drops of 0.1 M Pb(NO3)2 • 1 mL of water • drop-wise: 10 to 15 drops 6M H Cl (until no more granular white solid forms).
Chemistry 102L Manual - Page 158
• Into another test tube: • add a small pea sized portion of powdered Hg2(NO3)2. • add a 1 mL of water. • Heat the Hg22+ solution to a near boil. • Cool the Hg22+ solution and transfer it to a centrifuge tube. • Centrifuge the Hg22+ suspension and decant the solution into a test tube labeled “Hg22+”. • Into the “Hg22+” tube, add 10 to 15 drops of 6M HCl (until no more white precipitate forms). • Heat the “Ag+”, “Pb2+”, and “Hg22+” tubes in boiling water, and occasionally stir the contents of each test tube as it heats. •) Which of the solids dissolve? • Cool the tubes in ice water. •) Does the solid that dissolved above reform? > In principle, we can separate this ion from the other Class I cations by heating the chloride precipitate in hot water, centrifuging while hot, and then cooling the solution to get the chloride of that ion. If the solid indeed reforms, we have confirmed that that ion was in the original mixture. • With the chlorides that were insoluble in hot water: • Centrifuge the suspensions and discard the solutions. • Wash the solids with water. Reminder: “washing the solid” means: • Add water to the solid, and resuspend it with a stir rod. • Centrifuge the suspension, and discard the solution. • Add 2 mL of 6M NH3, and mix the suspension well. •) What happened in the two tubes: (Identify the cation and any changes in the solid and solution.)
> The solids in one tube should dissolve. For a mixture, we can now centrifuge the suspension and transfer the solution containing that ion into another tube leaving the other ion behind.
Chemistry 102L Manual - Page 159
> The solids in the other tube should change color. This is the test that confirms the presence of that ion. • If the solid dissolves in the NH3: • Add 1 M KI drop by drop until something happens. •) What happens? > The reaction you see is the test that confirms the presence of that ion in your solution. • Summarize your results above as a flow chart that allows you to separate the Class I cations and confirm their presence. Part B. Precipitation of Class II and Class III cations. • Into a test tube labeled “Cu2+, Zn2+” add: • 5 drops of 0.5 M Cu2+ • 5 drops of 0.2 M Zn2+ • 1 mL of water • 1 mL of thioacetamide (in the hood). • drop-wise 6 M HCl with stirring until the solution makes methyl violet paper turn blue-green. (Put a drop of 0.3 M HCl onto methyl violet paper for comparison.) • If the paper is purple or too blue, add more acid. • If the paper is yellow, add 6M NH3(aq). • Into a test tube labeled “Cu2+, Al3+” add: • 5 drops of 0.5 M Cu2+ • 5 drops of 0.2 M Al3+ • 1 mL of water • 1 mL of thioacetamide (in the hood). • drop-wise 6 M HCl with stirring until the solution makes methyl violet paper turn blue-green. (Put a drop of 0.3 M HCl onto methyl violet paper for comparison.) • If the paper is purple or too blue, add more acid. • If the paper is yellow, add 6 M NH3.
Chemistry 102L Manual - Page 160
• Heat these two tubes in a 70°C water bath (in the hood) for 10 minutes. •) Describe the precipitate in the “Cu2+, Zn2+” tube:
•) Describe the precipitate in the “Cu2+, Al3+” tube:
> The precipitates in both tubes should be CuS(s). • Centrifuge the precipitate in the two tubes and transfer the solutions into new test tubes with new “Cu2+, Zn2+” and “Cu2+, Al3+” labels. > Here, we have largely separated the Cu2+ from these solutions as CuS(s). • Add 6 M HCl drop-wise to both test tubes to adjust [H+] to 0.3 M as before. • Add 6 M NH3 drop-wise to both tubes until they are pH = 8. (Test this by touching a drop of the solution on a stir rod onto universal indicator paper.) • Add 1 mL of thioacetamide to both tubes and then heat them in the 70°C water bath for 10 minutes. • Separately, centrifuge the contents of both tubes and discard the solutions. • Wash the solids in both tubes twice with 2 or 3 mL of water. > The solid from the “Cu2+, Zn2+” tube is ZnS(s). > The solid from the “Cu2+, Al3+” tube is Al(OH)3. •) Is either solid colored?
> This color is from residual Cu2+ that was left over from the first separation. Since CuS is so darkly colored, only a little of it gives the color you see. This residual Cu2+ is nothing to worry about in a separation. Just be aware that it is there. • Add 5 drops of 6 M H Cl to these solids. •) Describe the odor from the “Cu2+, Zn2+” tube.
•) Describe the odor from the “Cu2+, Al3+” tube.
Chemistry 102L Manual - Page 161
•) Can you tell these tube apart by their odors. If so, how?
> Any difference you notice could be used to distinguish Zn2+ and Al3+ during an analysis. Part C. Sulfide Amphoterism: Separations of Class IIa and Class IIb. • Into a test tube labeled as “Bi3+”: • Add 1 mL water • 5 drops of 0.1 M Bi3+. • Into a test tube labeled as “Sn4+”: • Add 1 mL water • 5 drops of 0.1 M Sn4+. • Add 6 M HCl to both tubes until their solutions turn methyl violet paper blue-green. • Add 1 mL of thioacetamide to both tubes, and then heat them in the 70°C water bath (in the hood) for 5 minutes. • Separately centrifuge the suspensions from Bi3+ and Sn4+, and wash the solids with water. •) Describe the precipitate in the “Bi3+” tube.
> This solid should be Bi2S3(s). •) Describe the precipitate in the “Sn4+” tube.
> This solid should be SnS2(s). • With the solid from the “Sn4+” tube: • Add 6 M NaOH to the solid until it completely dissolves. > This is the reaction: SnS2(s) + 4OH-
Sn(OH)4- + 2S2-
• Add 6 M HCl to this Sn(OH)4- solution until SnS2(s) reprecipitates.
Chemistry 102L Manual - Page 162
• With the solid from the “Bi3+” tube: • Add about the same amount of 6 M NaOH to the solid, and warm the tube. •) Does much of the solid dissolve? • Centrifuge the contents of the “Bi3+” tube, and decant the solution into another tube. • Add 6 M H Cl to this solution until it is acidic. •) Does any solid form? •) If so, how does the amount compare to that from the “Sn4+” case above?
> Overall, though both Sn4+ and Bi3+ are Class II ions, they can be further separated by their solubilities in NaOH: the SnS2 dissolves while Bi2S3 does not. This allows the Sn(OH)4to be separated by centrifugation (as the solution), and then reprecipitated as the SnS2 by reacidification. The end result: Cation IIb ions (such as SnS2) can be separated from Cation IIa ions (such as Bi2S3). Part D. Separation within Class IV. • Into the first tube add: • 1 mL of water. • 5 drops of 0.1 M Mg2+. • 1 mL of 6 M NaOH •) Centrifuge the solution/suspension. Did a precipitate form? If so, describe it. Hint: look for a gel formed on the test tube wall. If it is not there, cool the tube in ice water for 5 minutes and then centrifuge a second time. Then look again.
• Save the tube for comparisons. • Into the second tube add: • 1 mL of water. • 5 drops of 0.1 M Mg2+. • 1 mL of 6 M NH3 •) Did a precipitate form? If so, describe it. (Note: a precipitate may not form.)
• If a solid formed, centrifuge and save it for comparisons.
Chemistry 102L Manual - Page 163
• Into the third tube add: • 1 mL of water. • 5 drops of 0.1 M Mg2+ • 0.5 mL of 3 M (NH4)2CO3 • 1 mL of 6 M NH3. •) Did a precipitate form? If so, describe it. (Note: a precipitate may not form.)
• If a solid formed, centrifuge and save it for comparisons. •) Compare the amount of solid in the three test tubes.
> The first, second, and third test tubes are in a progression of decreasing basicity. All three are capable of precipitating Ca2+ and Ba2+. We would like to precipitate these two ions but not Mg2+. Which procedure of the three above would I use to do this?
> After the Mg2+ is removed from the solid CaCO3 and BaCO3, we can then precipitate it as Mg(OH)2 with NaOH. Part E. Precipitation of Sulfur: How to identify it and what to do about it. • To two test tubes add: • 1 mL 6 M HNO3. • 1 mL thioacetamide. • Heat these tubes in a 70°C water bath (in the hood) for 10 to 15 minutes. • Centrifuge the sulfur precipitate, and wash it with water. Note: only a very small amount of solid may be there. • To one centrifuge tube, add 1 mL concentrated HCl. • To the other centrifuge tube, add 1 mL of 1.0 M Na2S.
Chemistry 102L Manual - Page 164
• Heat the tubes in a 70°C water bath for 5 minutes. •) Does the solid in the HCl dissolve?
•) Does the solid in the Na2S dissolve?
> We will be using the successful procedure to remove precipitate sulfur from sulfides during our analyses. This will prevent us from getting tricked into thinking we have precipitated sulfides when we really don’t. Part F. Using Solution Color for Cation Identification. The following solutions should be available in bottles on a tray in the laboratory. Just examine them to answer the following questions. •) What color is 0.1 M Ni(NO3)2? •) What color is 0.1 M Co(NO3)2? •) What color is 0.1 M Fe(NO3)3? •) What color is 0.1 M Cr(NO3)3? •) What color is 0.5 M Cu(NO3)2? Part G. Using Flame Tests for Cation Identification. • Find a clean ceramic spot plate and a clean Nichrome wire loop. • Light a Bunsen burner in a hood. As your instructor for help if needed. • Add 5 drops of each the following solutions into their own wells of the spot plate. (One cation per well.) 0.10 M Ca(NO3)2 0.10 M Ba(NO3)2 0.10 M Cu(NO3)2 0.10 M NaNO3 0.10 M KNO3 • Add 1 drop of 12 M HCl into each of these five wells. • Add 5 drops of 12 M HCl into a sixth well.
Chemistry 102L Manual - Page 165
Flame clean the Nichrome wire loop. Repeat the following three times. • Dip the Nichrome wire loop into the 12 M HCl (in the sixth well), and lift the loop out of the HCl such that a drop of the acid is hanging inside the loop. • Hold the loop and its drop of HCl in the middle of the Bunsen burner flame. You may see the flame quickly change color as the drop of HCl evaporates. This quick color change is due to cation residues on the wire (probably from the previous ). The flame will also likely begin to glow orange as the flame makes the Nichrome wire glow. This late color change (if it appears) is due to unavoidable contamination of the Nichrome and should be ignored during the following flame tests. For each of the cation/HCl solutions you prepared above, follow the following steps. (These are the flame tests). • Dip the Nichrome wire loop into a cation / HCl solution, and lift the loop out of the solution such that a drop of the solution is hanging inside the loop. • Hold the loop and its drop in the middle of the Bunsen burner flame. Watch for a quick color glow in the flame. This color should appear almost immediately. Don’t confuse this with the later orange color that appears when the wire begins glowing. Describe the flame color that quickly appears from each of the cation solutions: •) What color appears from Ca2+? •) What color appears from Ba2+? •) What color appears from Cu2+? •) What color appears from Na+? •) What color appears from K+? • Flame clean the Nichrome wire loop once as described above. In mixtures of Na+ and other cations, the strong color of Na+ could obstruct the color of others. To see these hidden colors, we can simply filter out the glow from Na+ with cobalt glass. • Find the dark purple glass squares in your equipment drawer. (You only need one.) • Add 1 drop of 0.10 M NaNO3 to the cation / HCl solutions above.
Chemistry 102L Manual - Page 166
Repeat the following for each of the Na+ / cation / HCl solutions. • Dip the Nichrome wire loop into a cation / HCl solution, and lift the loop out of the solution such that a drop of the solution is hanging inside the loop. • Hold the loop and its drop in the middle of the Bunsen burner flame as you watch through the cobalt glass. (Keep your goggles on!) Watch for a nearly immediate color glow in the flame. Again, don’t confuse this with the later orange color that appears when the wire begins glowing. Describe the flame color that quickly appears from each of the cation solutions: •) What color appears from the solution with only Na+? •) What color appears from Ca2+ and Na+? •) What color appears from Ba2+ and Na+? •) What color appears from Cu2+ and Na+? •) What color appears from K+ and Na+? • Flame clean the Nichrome wire loop once as described above.
Chemistry 102L Manual - Page 167
Report Form: SEPARATION AND IDENTIFICATION OF SELECTED CATIONS Name
Date
Laboratory Instructor
Section
A. Separation and Identification of Ag+, Hg22+, and Pb2+ 1. a. Which chloride (AgCl, Hg2Cl2 or PbCl2) dissolved in hot water? __________ b. Did precipitate form when the hot tube was cooled? __________ 2. a. Which chloride gave a black precipitate with 6 M NH3? __________ b. Net Reaction:
3. a. Which chloride dissolved in 6 M NH3? __________ b. Net Reaction:
c. What compound precipitated when KI was added? __________ d. Net Reaction:
4. Using the reactions described above, draw a flow chart that summarizes how to separate and identify AgCl, PbCl2 and Hg2Cl2 in a mixture.
Chemistry 102L Manual - Page 168
B. Precipitation of Sulfides - Separation of Class II and III 1. a. What was the color of the precipitate produced when a solution of Zn2+, Cu2+ reacted with H2S at 0.3 M H+? b. What was the formula for the precipitate?
2. a. What was the color of the ZnS precipitated at pH 8?
b. What was the color of the Al(OH)3 precipitated at pH 8? c. As the H+ decreases, does the concentration of S2- increase or decrease? As this happens, would you expect more CuS to form if some Cu2+ were still in solution?
Based on this, would you expect more CuS to form?
d. What happened when 6M HCl was added to Al(OH)3(s) and ZnS(s)?
e. Could you distinguish ZnS(s) and Al(OH)3(s) by odor when you added acid? How?
3. a. Given that Bi2S3 is black and SnS2 is yellow, which cation (Bi3+ or Sn4+) precipitated as a sulfide with thioacetamide at pH 8?
b. Which cation (Bi 3+ or Sn4+) would not form a sulfide precipitate at pH 8 but would when the H+ concentration is increased to 0.3 M H+.
Chemistry 102L Manual - Page 169
c. Given that the equilibrium constant K for hydrolysis of Sn4+ to Sn(OH)62- is large: Sn4+(aq) + 6H2O(l)
Sn(OH)62-(aq) + 6H+(aq)
that this is the reaction that couples to the reaction below to make SnS2(s) dissolve: SnS2(s)
Sn4+(aq) + 2S2-(aq)
i.
Considering the first equilibrium, what would happen if the pH decreased? Specifically, would the Sn4+ concentration increase or decrease? Explain.
ii.
How would the change in Sn4+ concentration effect the second equilibrium? Specifically, what would happen to the amount of SnS2(s)? Explain.
iii.
What did you see happen when KOH was added to SnS2? Does this agree with the behavior you predicted above?
C. Sulfide Amphoterism - Separations within Class II 1. Was Bi2S3 insoluble in the KOH solution? 2. Based on your tests, is KOH adequate (as it is used in the procedure) for separating Bi2S3 and SnS2? If not, explain.
D. Separations within Class IV 1. Did you get a precipitate when you added 1 mL of 6 M NH3 to the Mg2+ solution which didn't contain (NH4)2CO3? 2. Was the amount of precipitate similar to what you got adding 1 mL of 6 M NaOH?
Chemistry 102L Manual - Page 170
3. Did you get a precipitate when you added 1 mL of 6 M NH3 to the solution containing (NH4)2CO3 and Mg2+? If so, how did the amount compare with that you got without the (NH4)2CO3?
4. Based on your results, would it be wise to change the flow chart procedure and use 6 M NH3 in place of 6 M NaOH for precipitating Mg2+ as Mg(OH)2? Why?
E. Precipitation of Sulfur 1. Did sulfur precipitate rapidly or slowly?
2. a. Was the sulfur soluble in concentrated HCl?
b. Was sulfur soluble in 1 M Na2S?
F. Identification of Cations by Color Complete the table below with the colors you observed. Cation Ni2+ Co2+ Fe3+ Cr3+ Cu2+
Cation Color
Chemistry 102L Manual - Page 171
G. Identifications of Cations by Flame Color Complete the table below with the colors you observed (without the cobalt glass). Cation
Flame Color
Ca2+ Ba2+ Cu2+ Na+ K+ Complete the table below with the colors you observed with the cobalt glass. Cation + Na+ Na+ Ca2+ + Na+ Ba2+ + Na+ Cu2+ + Na+ K+ + Na+
Flame Color (through cobalt glass)
Chemistry 102L Manual - Page 172
Questions
1. The following observations were made for an unknown solution that possibly contained the ions Hg22+, Ag+, Cu2+, Al3+, Ba2+, Na+, Ca2+, K+, Mg2+. For each observation, briefly interpret its meaning (what ions are present or absent). • The solution began colorless.
• Addition of HCl gave a white precipitate.
• This white precipitate was not soluble in hot water.
• With the addition of NH3, the white precipitate turned black and none dissolved.
• When the solution, from which the white precipitate was separated, was heated with thioacetamide (to generate H2S), no precipitate formed at H+ concentrations of 0.3 M. • When 1 mL of 6M NH3 was added to a portion of this solution no precipitate formed.
at pH 8,
• When NH3 was added to another portion of the pH 8 solution (until its pH was 9) and the resulting solution was mixed with 3M (NH4)2CO3, a white precipitate formed.
• The white precipitate dissolved with fizzing in concentrated HCl, and this HCl solution gave a green flame test.
Chemistry 102L Manual - Page 173
2. Suppose we were adding concentrated HCl to 1.0·10-3 M Pb2+(aq). What minimum concentration of HCl would be needed to make PbCl2 begin precipitating? [Ksp(PbCl2) = 1.6·10-5]. Show your work. [Hint: this would be the concentration of Cl- in equilibrium with PbCl2(s) and 1.0·10-3 M Pb2+(aq).]
3. a. If a 0.05 M Mn2+ solution were saturated with H2S (0.10 M H2S) and the pH was adjusted to 0.50, could any MnS precipitate? [Ksp(MnS) = 2.5·10-10]. Show your work. Hint: In the introduction, we found that: M2+ + H2S MS(s) + 2H+
K = (1.07·10-20) / Ksp
Chemistry 102L Manual - Page 174
b. If a 0.05 M Mn2+ solution were saturated with H2S (0.10 M H2S) and the pH was adjusted to 8.0, would any MnS precipitate? Show your work..
c.Given the behavior in parts (a) and (b), what class of cations would Mn 2+ belong to and why?
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