Experiment 24. Qualitative Analysis I

Transcription

1 Experiment 24 Qualitative Analysis I GOAL: In this experiment you will explore characteristic reactions of a number of ions. The reactions you observe will be used to identify unknown solutions next week. We will use this study to develop logical schemes for analyzing unknowns, and also to learn and review some descriptive chemistry of the elements. INTRODUCTION: You will be studying seven cations (Na 1+, Mg 2+, Ni 2+, Cr 3+, Zn 2+, Ag 1+, Pb 2+ ) and four anions (NO 3 1-,Cl 1-, I 1-, SO 4 2- ). These ions may be separated from each other and identified using simple chemical and physical properties. Characteristic colors of aqueous ions as well as solid compounds can be very useful in identification. In general, we expect main group cations to be colorless in solution and to form white solids. Transition metal cations often have characteristic colors both in the aqueous solution and in solid compounds. Some ions that generally form colorless compounds can give brilliant colors if mixed with the right partner. For example, both Pb(NO 3 ) 2 and KI dissolve to give colorless solutions, but when mixed, a bright yellow precipitate of PbI 2 is produced. I. Reactions of cations with NaOH and NH 3 Most metal ions react with aqueous OH 1-, hydroxide ion, to form gelatinous precipitates of the metal hydroxides. For example, Al 3+ (aq) reacts to form Al(OH) 3 (s). Al 3+ (aq) + 3 OH 1- (aq) Al(OH) 3 (s) Eqn 1 Ag 1+ is a rare exception, forming Ag 2 O(s) rather than AgOH. All other hydroxide precipitates you encounter in this experiment are neutral hydroxides. Adding more OH 1- causes some insoluble hydroxides to re-dissolve. For example, Al(OH) 3 (s) reacts to form the complex ion Al(OH) 4 - (aq). In lab, if you observe a precipitate re-dissolving upon addition of more OH 1-, it is doing a reaction like Eqn 2. Al(OH) 3 (s) + OH 1- (aq) Al(OH) 4 1- (aq) Eqn 2 In this experiment metal hydroxide precipitates that re-dissolve in excess OH 1- form complex ions with the general formula M(OH) 4 n- (aq). Metals ions may react with NH 3 to form either insoluble precipitates or soluble complex ions. If a precipitate forms, it is the neutral hydroxide. For example, Al 3+ (aq) reacts with NH 3 (aq) to form Al(OH) 3 (s). In lab, if you observe an ion reacting with NH 3 to form a precipitate, it is doing a reaction like that in Eqn 3. Al 3+ (aq) + 3 NH 3 (aq) + 3 H 2 O(l) Al(OH) 3 (s) + 3 NH 4 1+ (aq) Eqn 3 If no precipitate forms or is only briefly present, a complex ion with NH 3 has formed, for example, Zn(NH 3 ) 4 2+ (aq), Ni(NH 3 ) 4 2+ (aq), and Ag(NH 3 ) 2 1+ (aq). See Eqn 4. Zn 2+ (aq) + 4 NH 3 (aq) Zn(NH 3 ) 4 2+ (aq) Eqn 4 1

2 II. Anion Precipitation Reactions Of the cations we are studying, only Ag 1+ and Pb 2+ react with any of our four study anions to form precipitates. Ag 1+ should form precipitates with two of the anions; while Pb 2+ should form precipitates with three of them. The colors and solubilities of these precipitates will allow us to determine the presence of these anions in unknown mixtures. Any precipitates that form in these reactions are simple insoluble salts of the ions. Occasionally when a large excess of one ion, especially the anion, is present, anomalous results may be noted. Ag 2 SO 4, which is moderately soluble, normally remains dissolved under our conditions, but may form a fine white precipitate when concentrated. PbCl 2 and PbI 2, which are normally insoluble under our conditions, may re-dissolve by forming complex ions with excess anion. The solubility of PbCl 2 is also highly temperature dependent. A PbCl 2 precipitate will re-dissolve if it is heated in water. This fact may be used to identify the precipitate since the other Pb 2+ precipitates will not re-dissolve when heated. III. Stability Sequences Occasionally individual precipitation reactions are insufficient to allow identification of an unknown mixture. A stability series such as those shown below may be helpful. A stability series is a ranking of salts and ions for a given cation with the most stable species on the right. Addition of an appropriate reagent to a compound can convert that compound into a compound further to the right in the series. However, the equilibria will not allow you to go from right to left. For Ag + : Ag 2 O(s) < AgCl(s) < Ag(NH 3 ) 2 + (aq) < AgI(s) For Pb 2+ : PbCl 2 (s) < PbSO 4 (s) < PbI 2 (s) < Pb(OH) 2 (s) < Pb(OH) 4 2- (aq) in xs OH - For Zn 2+ : Zn(OH) 2 (s) < Zn(NH 3 ) 4 2+ (aq) < Zn(OH) 4 2- (aq) in xs OH - For Ni 2+ : Ni(OH) 2 (s) < Ni(NH 3 ) 4 2+ (aq) < Ni(OH) 2 (s) in xs OH - For Cr 3+ : Cr(OH) 3 (s) < Cr(OH) 4 1- (aq) in xs OH 1- From the lead stability series, for example, we can see that addition of SO 4 2- (aq) to solid PbCl 2 (or aqueous Pb 2+ ) would cause PbSO 4 (s) to form. You could not, however, go the other direction. Adding Cl - (aq) to PbSO 4 (s) will not form PbCl 2 (s). We could, however, add I - (aq) to Pb 2+ (aq), PbCl 2 (s) or PbSO 4 (s) and get PbI 2 (s). Addition of excess OH 1- to any of species in the lead series would cause the formation of the complex ion Pb(OH) Remember that any species shown as an ion in the stability series above will be soluble in water. The neutral compounds will precipitate. IV. Nitrate Ions Nitrate ions, NO 3 1-, will not precipitate, regardless of the ions mixed with them. The brown ring "spot test" has been developed to identify this ion. V. Sodium Ions Sodium ions, Na 1+, will not precipitate, regardless of the ions mixed with it. A flame test is the easiest way to identify Na 1+. 2

3 WASTES AND HAZARDS: All waste from qualitative analysis experiments should be placed in the combined qual waste beaker in the hood. Acid and base solutions are corrosive. Use care when handling these solutions. Although the ion solutions should be handled with respect, low concentrations minimize hazards that might be associated with solid or concentrated forms of the same compounds. One of the bases used in this experiment is aqueous ammonia. Look up its MSDS. Record in the Hazards section of your notebook entries, appearance and odor concentration (found under Composition) potential health effects and the appropriate first aid for o inhalation o skin contact o eye contact Given that our 6M NH 3 is about 10%, do you think the hazards of our solutions are quite similar to or quite different from what you found on the MSDS? Explain. LABORATORY OBSERVATIONS AND DATA: In the work that follows, keep clear, complete laboratory notes in your notebook. Be sure to fully describe each reagent before mixing, and each final solution after mixing. Include colors, precipitates, etc. Do not simply record no change upon mixing. Give the color and state of the system, e.g. green precipitate remains. Some of the precipitates will be gelatinous. Any cloudiness is a precipitate, even if the solid doesn t settle to the bottom. Be sure to mix all reactions thoroughly. PROCEDURE: Stock solutions of the seven cations (Na 1+, Mg 2+, Ni 2+, Cr 3+, Zn 2+, Ag 1+, Pb 2+ ) are supplied as 0.1 M solutions of the nitrate salts. Stock solutions of NO 3 1-, I 1-, and SO 4 2- are 0.1 M sodium salt solutions. The Cl 1- solution is 0.5 M NaCl and is slightly acidified. Other reagents available include 6M solutions of NaOH, NH 3, HCl, and HNO 3. Record your observations in the table provided. I. Reactions of cations with NaOH and NH 3 All Na 1+ salts are soluble, including NaOH. Consequently, Na 1+ will not form a precipitate with NaOH or NH 3. We will save a little time, and not perform these tests. You should still, however, record in your table of results that Na 1+ gives no precipitate with any of the reagents. Obtain a well plate that has six columns of four wells each. Add 10 drops of Mg 2+ (aq) to each of two wells in the first column. Similarly, prepare two wells each with 10 drops of the remaining five cations: Ni 2+, Cr 3+, Zn 2+, Ag 1+, and Pb 2+. Add one drop of 6M NaOH to one of the wells for each cation. Tap the well plate to mix. Note any reactions that occur. Record your observations. Any precipitates (even just cloudy ones) are insoluble hydroxide salts. Now add 10 more drops of 6M NaOH to the same wells. Tap the well plate to mix. Note any reactions that occur. Any precipitates that re-dissolve are reacting with aqueous OH 1- to form complex ions. Formulas for these complex ions can be found in the introduction. If a precipitate does not re-dissolve, no reaction has occurred. Record colors and the presence of any precipitates (even cloudiness). 3

4 Finally, add 10 drops of NH 3 (aq) to each of the wells in the second row. Mix and observe any reactions. When precipitates form, the metal cations have reacted with NH 3 and H 2 O to form insoluble hydroxide salts. If no precipitate forms or forms briefly and immediately re-dissolves, the metal cation has reacted with NH 3 to form a soluble complex ion with the ammonia. Formulas for these complex ions can be found in the introduction. Record colors and the presence of any precipitates (even cloudiness). Dump your used solutions in the designated waste container. Rinse the well plate thoroughly with distilled water and shake it briefly to remove any excess water. II. Anion Precipitation Reactions Put 10 drops of Ag 1+ into each of four wells, and 10 drops of Pb 2+ into another four wells. Add 1 drop of Cl 1- solution to the first Ag 1+ and Pb 2+ wells. Note the colors of any precipitates. Add an additional 10 drops of Cl 1- to each of these wells. Note any changes. Repeat this procedure using the remaining three anions (I 1-, SO 4 2-, and NO 3 1- ) and the remaining unused wells of Ag 1+ and Pb 2+. If you do not get the proper number of precipitates (two for Ag 1+, and three for Pb 2+), then repeat. III. A Stability Sequence The following four reactions illustrate the stability sequence for Ag 1+. Record observations for each of the reactions, noting precipitates, colors, and the re-dissolution of precipitates. Rxn 1: Combine 1 ml AgNO 3 (aq) and 1 ml NaOH(aq) in a test tube. Rxn 2: Centrifuge and discard the liquid from Rxn 1. Add several droppersful of NaCl(aq) to the precipitate from Rxn 1. Mix thoroughly. Rxn 3: Centrifuge and discard the liquid from Rxn 2. Add NH 3 (aq) to the solid until it dissolves. Rxn 4: Add NaI dropwise to the solution from Rxn 3 until a reaction is seen. IV. Sodium Ions For this experiment, test a solution of NaNO 3 to see the characteristic test for Na 1+. Sodium flame test procedure: Dip a cotton swab into the solution to be tested and then heat end of the swab in a Bunsen burner flame. A very intense yellow flame is observed when Na 1+ is present. Douse the swab in water to extinguish the flame and dispose of it in the container provided. When using a flame test to determine the presence of Na 1+ in an unknown, you must always use a new swab to avoid contamination. Despite even your best efforts, you may still see a slight yellow flame when no Na 1+ is present in the unknown. This happens because nearly everything is contaminated with small amounts of Na 1+. 4

5 V. Nitrate Ions Perform the brown ring test for nitrates on a sample of aqueous NaNO 3. Record your observations. Be careful! It is very easy to get a false negative result from the brown ring test. Nitrate brown ring test. Put 20 drops of the solution to be tested into a test tube, and then, carefully and slowly, add 20 drops of concentrated H 2 SO 4. (The test tube will get hot, and may splatter if the acid is added too quickly.) If necessary, cool this mixture by allowing cold tap water to run over the outside of the test tube. In a second test tube, dissolve about 0.1 g FeSO 4. 7H 2 O in 1 ml of water. Now, hold the test tube containing the H 2 SO 4 solution at a 45 o angle while you allow 5 drops of the FeSO 4 solution to run slowly down the inside of the tube. Do NOT mix. The aqueous FeSO 4 should form a layer above the acid. If nitrate ion is present, a smoky brown ring will form at the solution interface. This ring may take several minutes to form, and will disappear eventually. Look at the interface between the two layers from several angles. If you were doing this test on a real unknown sample, and there were a chance that the sample also contained I 1-, this would interfere with the test. To remove I 1-, add several drops of saturated Ag 2 SO 4 solution until no more AgI precipitates. RESULTS AND DISCUSSION: At the end of this handout, a blank table is provided for you to record your observations for when the cationic solutions are mixed with the following: OH 1-, xs OH 1-, NH 3, and anion precipitation. (You may not have mixed some of the combinations because we knew nothing would happen. Record this information in your table anyway.) In the last column record your initial observations, noting especially the color of each ion solution. Some additional notes are provided in this column as well. In each box, list solution colors, formulas for any complex ions formed, formulas of any precipitates formed, and precipitate colors. See the Introduction for help with formulas. DO NOT write something like no change. Record the color and state of the chemicals present. The anion precipitation column will have entries for only Ag 1+ and Pb 2+. The other cations will not precipitate with any of the anions. Make a copy of this table to keep for yourself after you turn in your report. You will need this table next week to help identify unknown solutions. AN INTRODUCTION TO FLOWCHARTS: The reactions you observed this week can be used to separate and identify unknowns. Flowcharts are a useful way to track these reactions. You will learn much more about flowcharts next week when you work with solutions of unknown ions. This week, you'll learn just enough to answer the question below and get you ready for next week. Let s imagine we were testing a solution that could contain any of our seven cations. Would adding I 1- to this mixture let us decide what was present? Let s answer this by drawing a flowchart to show the possibilities. Start by listing all the ions that could be present. The vertical line with I 1- next to it is our way of showing that we add some I 1-. Na 1+ (aq), Mg 2+ (aq), Zn 2+ (aq), Ag 1+ (aq), Pb 2+ (aq) all colorless Ni 2+ (aq) green, Cr 3+ (aq) purple/blue I 1-5

6 Look at your results from this week and consider what would happen as each ion reacts with I 1-. As you saw in lab, Ag 1+ and Pb 2+ will precipitate as AgI(s) and PbI 2 (s), while all other metal ions remain dissolved in solution. We represent this with the flow chart below. Na 1+ (aq), Mg 2+ (aq), Zn 2+ (aq), Ag 1+ (aq), Pb 2+ (aq) all colorless Ni 2+ (aq) green, Cr 3+ (aq) purple/blue I 1- AgI(s) pale yellow PbI 2 (s) bright yellow Na 1+ (aq), Mg 2+ (aq), Zn 2+ (aq) all colorless Cr 3+ (aq) purple/blue, Ni 2+ (aq) green Notice that the chart separates precipitates (to the left) from solutions (to the right) and that the formulas of precipitates and complex ions are used. In the lab we can accomplish a similar separation. By centrifuging and decanting, the precipitated AgI and PbI 2 can be physically separated from the other ions that are still in solution. Also notice that when precipitates form, we use a double drop line, rather than a single line, in the flowchart. Note that the flowchart above is being used to show all the possible things you might see. If you were actually identifying an unknown in lab, you would include only the things you actually see. How would we record our in-lab work? Let s say you were given an unknown and told that it contained only one of the seven cations. What is the first thing you would observe? The color! Your unknown is a clear and colorless solution. You would write Unknown #1, only 1 cation, colorless Could be Na 1+ (aq), Mg 2+ (aq), Zn 2+ (aq), Ag 1+ (aq) or Pb 2+ (aq) You have already eliminated Ni 2+ and Cr 3+ based upon color, so there is no reason to include them in your flowchart. Now which test do you run? Several good choices are available, but let s stick with the addition of I 1- since we used that above. In lab, you add some I 1- to a bit of your unknown and still have a clear, colorless solution with no precipitate. You should record this as Unknown #1, only 1 cation, colorless Could be Na 1+ (aq), Mg 2+ (aq), Zn 2+ (aq), Ag 1+ (aq) or Pb 2+ (aq) I 1- No ppt., so don t have Colorless, could be Pb 2+ or Ag 1+ Na 1+ or Mg 2+ or Zn 2+ Since you didn t get a precipitate, do you really need to record no ppt, or can you just write nothing under the double drop line? It is safer to write no ppt. When you come back to this later, your observations will be clear. If you had simply left it blank, you might think that you had forgotten to record a precipitate that formed. Was this addition of I 1- a useful way to identify your unknown? In this case, we narrowed the 6

7 possibilities but will need to do additional steps to choose from among the remaining possibilities. We will need to do a second test of the original solution, or we might add tests to the separated precipitate and solution. We would then add these steps to the flowchart. Next week you will do such multi-step flowcharts. Does that mean that adding I 1- isn t a good way to start our unknown identification? It s certainly not the only step we might have chosen, but a different unknown might make things seem quite different. Let s say you get a second unknown, it is also clear and colorless and contains only one of our seven cations. Again you choose to add I 1-, but this time you get a pale cream precipitate. You should record this as Unknown #2, only 1 cation, colorless Could be Na 1+ (aq), Mg 2+ (aq), Zn 2+ (aq), Ag 1+ (aq) or Pb 2+ (aq) I 1- pale cream ppt must be AgI(s) so unknown contained Ag 1+ colorless soln In one step we have identified this unknown. We don t need to do any additional tests or try to list possibilities that we didn t actually see. Since you now know that the unknown contained Ag 1+, you no longer list any other possibilities. QUESTIONS: 1. Give balanced chemical equations. Check your lab results and the Introduction to decide what happens in each case. Some may be N.R., no reaction. a. Cr 3+ (aq) + xs OH 1- (aq) b. Ni 2+ (aq) + NH 3 (aq) c. Na 1+ (aq) + NH 3 (aq) d. Cr 3+ (aq) + NH 3 (aq) e. Pb 2+ (aq) + xs OH 1- (aq) f. Pb 2+ (aq) + Cl 1- (aq) g. Ni 2+ (aq) + Cl 1- (aq) 2. Draw a flow chart, similar to the first one in the Introduction to Flowcharts section, to show what would happen if you did each of the tests below. Start each flowchart with the specified ions in the mixture listed at top. Include the color of each of those ions in aqueous solution. After the specified reactant has been added, you will have both a precipitate and a supernatant solution. Give correct formulas and colors for the precipitates. Give correct formulas and colors for the ions in the supernatant solution. Be careful; you may get some complex ion formation. a. excess OH 1- added to a mixture of all seven metal ions b. NH 3 added to a mixture of aqueous Na 1+, Mg 2+, Ni 2+, Cr 3+, and Zn 2+ c. I 1- added to a mixture of Ag 1+, Zn 2+, and Ni Carefully re-read the Introduction Section III Stability Sequences. Write balanced equations. a. PbCl 2 (s) + excess OH 1- (aq) [see the stability sequence for Pb 2+ ] b. AgCl(s) + excess OH 1- (aq) [see the stability sequence for Ag 1+ ] 7

8 4. While again watching those stability sequences, draw flow charts showing what would happen if you added a. excess OH 1- to a mixture of AgCl(s) and PbCl 2 (s) [see your answers to #3] b. NH 3 to a mixture of AgCl(s) and PbCl 2 (s) [see stability sequences!] 5. You are given a test tube containing an aqueous solution of a nitrate salt of one of our cations (Na 1+, Mg 2+, Ni 2+, Cr 3+, Zn 2+, Ag 1+, or Pb 2+ ) but you don t know which it is. You must use the reactions of this experiment to identify the single cation in the test tube. If you could do only one test, which would be most helpful in identifying the unknown or at least narrowing the possibilities? Consider these options: add 1 drop OH 1-, add excess OH 1-, add NH 3, add Cl 1-, add I 1-, do a flame test, or do a brown ring test. Write several well-organized paragraphs in which you consider these possibilities, discuss the merits of each option, and then finally draw a conclusion about the best choice or choices. 8

9 Characteristic Reactions for Identifying Unknown Ions Test Ion OH - XS OH - NH 3 Anions (Cl -, I -, SO 2-4 ) Na + Miscellaneous Yellow flame Mg 2+ Slow pption Ni 2+ Cr 3+ Zn 2+ Ag + Pb 2+ NO 3 - Stains black; Ag 2 SO 4 may ppt at high conc PbCl 2 will dissolve in HOT water All nitrates are soluble; Brown ring test Stability sequences Some precipitates are less soluble than others. You can only move left to right. For Ag + : Ag 2 O(s) < AgCl(s) < Ag(NH 3 ) + 2 (aq) < AgI(s) For Pb 2+ : PbCl 2 (s) < PbSO 4 (s) < PbI 2 (s) < Pb(OH)2(s) < Pb(OH) 2-4 (aq) in XS OH - For Zn 2+ : Zn(OH) 2 (s) < Zn(NH 3 ) 2+ 4 (aq) < Zn(OH) 2-4 (aq) in XS OH - For Ni 2+ : Ni(OH) 2 (s) < Ni(NH 3 ) 2+ 4 (aq) < Ni(OH) 2 (s) in XS OH - For Cr 3+ : Cr(OH) 3 (s) < Cr(OH) - 4 (aq) in XS OH - 9