Objectives Name: Separation and Identification of Group D Cations (Cu 2, Ni 2, Mg 2 and Zn 2 ) To understand the chemical reactions involved in the separation and identification of the Group D cations. To complete a flow diagram summarizing the qualitative analysis scheme for the Group D ions. To successfully identify the Group D cation(s) in an unknown. Background Chemistry and Discussion The Group D cations are characterized as having hydroxides, oxides and oxalates that are soluble in an ammonia/ammonium ion buffer solution. Three of the Group D cations, Cu 2, Ni 2 and Zn 2, form complex ions with ammonia. These complexes ions are very stable and prevent the cations from precipitating when oxalate is added even though their oxalates are relatively insoluble. Many complex ions are highly colored and the color of your solution at this stage may help in identification of the cations in your unknown. Testing known solutions to determine the color of the complex ions formed is recommended. Magnesium hydroxide will precipitate in aqueous ammonia, however the precipitate can be prevented from forming if the solution contains a significant concentration of NH 4. This can be understood by considering the common ion effect and the following equilibrium: NH 3 (aq) H 2 O(l) NH 4 (aq) OH (aq) In an NH 3 /NH 4 buffer solution, this equilibrium is shifted left, reducing the hydroxide ion concentration so that Mg 2 does not precipitate. Magnesium oxalate has a K sp value that is relatively high and under the conditions of the experiment should not have precipitated if a limited amount of oxalate was added in a previous separation step. The presence of ammonia and possibly oxalate ions in the supernatant liquid containing Group D can interfere with characterization tests for the cations in this group. Therefore, prior to testing the ammonia and oxalate, if present, must be removed from the solution. This is accomplished by heating the solution in a crucible over a low flame to dryness, adding concentrated HNO 3 and then reheating to dryness. THIS HEATING MUST BE DONE IN THE HOOD! The following reactions occur: NH 4 (aq) Cl (aq) NH 4 Cl(s) NH 3 (g) HCl(g) NH 4 (aq) NO 3 (aq) NH 4 NO 3 (s) N 2 O(g) 2H 2 O(g) And if oxalate is present: 2NO 3 (aq) C 2 O 4 2 4H (aq) 2NO 2 (g) 2CO 2 (g) 2H 2 O(g) Note that if oxalate is not present in your sample, then the HNO 3 need not be added prior to heating; NH 3 will still be removed as shown in the reaction that does not involve the nitrate ion given above. After heating to dryness, the residue contains the oxides of the Group D cations: Copper (II) oxide is black, nickel oxide is green and magnesium and zinc oxide are both white. The residue is dissolved in 6 M HCl(aq) and the resulting solution can be tested directly for Ni 2 and Cu 2. (In HCl, copper (II) and nickel ions form complex ions: CuCl 4 2 and NiCl 6 4.) After testing for Ni 2 and Cu 2, the solution is treated to remove Ni 2 and Cu 2 ions, if present, prior to testing for Mg 2 and Zn 2. Ni 2 and Cu 2 Confirmation Tests: Identification of Ni 2 : A portion of the unknown solution is made basic with aqueous ammonia and then a drop of dimethylglyoxime ((CH 3 ) 2 C 2 (NOH) 2, abbreviated as DMG) solution is added. The immediate formation of nickel dimethylglyoxime complex, a cherry red precipitate, confirms the presence of Ni 2. Nickel ion forms the only highly colored precipitate with dimethylglyoxiniate ions. The solution must be made basic with ammonia in order for the precipitate to form since the DMG must be deprotonated before it will react with any Ni 2 that is present. (CH 3 ) 2 C 2 (NOH) 2 (aq) NH 3 (aq) (CH 3 ) 2 C 2 (NOH)(NO) (aq) NH 4 (aq) 2(CH 3 ) 2 C 2 (NOH)(NO) (aq) Ni(NH 3 ) 6 2 (aq) Ni[(CH 3 ) 2 C 2 (NOH)(NO)] 2 (s) 6NH 3 (aq) Identification of Cu 2 : The confirmation test for Cu 2 is the formation of a maroon precipitate of copper hexacyanoferrate, Cu 2 [Fe(CN) 6 ], from slightly acidic solution: Dr. L.J. Larson 1 Revised/Printed 6/28/08
2Cu 2 (aq) Fe(CN) 6 4 (aq) Cu 2 [Fe(CN) 6 ](s) In the absence of Cu 2 a white or pale green precipitate may be observed since hexacyanoferrate, [Fe(CN) 6 ] 2, forms precipitates with other Group D cations. Before adding hexacyanoferrate, the strongly acidic solution is first treated with aqueous ammonia until it is only slightly acidic. The purpose of this is to avoid the addition of hexacyanoferrate to a strongly acidic solution, since this would result in the formation of hydrogen cyanide (HCN), a very toxic gas. By reducing the acidity of the solution, this possibility is avoided. HOWEVER, THIS TEST MUST STILL BE PERFORMED IN THE HOOD. After adjusting the ph of a portion of the test solution, potassium hexacyanoferrate, K 4 [Fe(CN) 6 ], is added. The immediate formation of a maroon precipitate confirms the presence of copper(ii) ion. Separation of Mg 2 and Zn 2 from Cu 2 and Ni 2 : Before testing for magnesium and zinc ions, the copper (II) and nickel ions are removed. This is accomplished by precipitating the copper (II) and nickel ions as their sulfides. The source of sulfide ion will be a hot, weakly acidic solution of sodium thiosulfate, Na 2 S 2 O 3. Under these conditions sodium thiosulfate disproportionates to form sulfate and sulfide ions. The unbalanced net-ionic chemical equation for this reaction is: S 2 O 3 2 (aq) S 2 (aq) SO 4 2 (aq) Copper (II) sulfide is insoluble, even in acidic conditions, and nickel sulfide has only slight solubility in the weakly acidic conditions imposed. Zinc sulfide is soluble under the conditions used; it does not precipitate. Magnesium does not form a sulfide. Thus, upon heating in weakly acidic sodium thiosulfate solution, Cu 2 and Ni 2 ions will precipitate, forming CuS(s) and NiS(s), both of which are black. Magnesium and zinc ions are left in the supernatant that is used to test for their presence. Due to the slight solubility of NiS in the weakly acidic solution, any Ni 2 ions present may not completely precipitate leaving some Ni 2 in solution along with the Zn 2 and Mg 2 ions. Mg 2 and Zn 2 Confirmation Tests: The identification of magnesium is accomplished using the supernatant remaining after the precipitation of Cu 2 and Ni 2 as their sulfides. This supernatant may contain Zn 2 and Mg 2, and possibly some Ni 2. Magnesium and zinc ions are typically colorless in solution; the presence of Ni 2 may give a characteristic color to the solution. The solution is made slightly basic with aqueous ammonia, and disodium hydrogen phosphate, Na 2 HPO 4, is added. Under these conditions, Zn 2 will precipitate as white zinc phosphate, Zn 3 (PO 4 ) 2, and Mg 2 will precipitate as white magnesium ammonium phosphate, MgNH 4 PO 4. Any other Group D cation present at this point remains dissolved since their phosphates are soluble in ammonia. An acid-base chemical equilibrium is involved is these precipitation reactions as shown below: Acid Base Equilibria: Precipitation Reactions: NH 3 (aq) HPO 4 2 (aq) NH 4 (aq) PO 4 3 (aq) Mg 2 (aq) NH 3 (aq) HPO 4 2 (aq) MgNH 4 PO 4 (s) 3Zn 2 (aq) 2NH 3 (aq) 2HPO 4 2 (aq) Zn 3 (PO 4 ) 2 (s) 2NH 4 If the solution is made too basic with ammonia, a precipitate may not form. In this case, adding a small amount of HCl(aq) may work to produce a precipitate. After being centrifuged, the supernatant is discarded and the precipitate is treated with NaOH. Any zinc phosphate will dissolve to form the colorless tetrahydroxozincate(ii) complex ion, [Zn(OH) 4 ] 2. The unbalanced net ionic chemical equation for this reaction is: Zn 3 (PO 4 ) 2 (s) OH (aq) [Zn(OH) 4 ] 2 (aq) PO 4 3 (aq) Magnesium ammonium phosphate is insoluble in NaOH solution. A white precipitate that remains after addition of the NaOH confirms the presence of Mg 2. The basic supernatant is tested for the presence of zinc by making the solution slightly acidic with aqueous acetic acid and then adding K 4 [Fe(CN) 6 ]. THIS TEST MUST BE PERFORMED IN THE HOOD. The immediate formation of a white precipitate of Zn 2 [Fe(CN 6 )] confirms the presence of zinc(ii) ion. 2Zn 2 (aq) Fe(CN) 6 4 (aq) Zn 2 [Fe(CN) 6 ](s) For both the Mg 2 and Zn 2 tests, the precipitates formed are white. The existence of traces of colored ions such as Ni 2 and Cu 2 can add some color to the white precipitates. Modifications to the procedure may be needed if results are ambiguous. Dr. L.J. Larson 2 June 28, 2008
Reagents Available 6-M NH 3 1% dimethylgloxime 6-M HCl 0.1-M K 4 Fe(CN) 6 6-M NaOH Solid Na 2 S 2 O 3 6-M acetic acid (CH 3 COOH) Saturated Na 2 HPO 4 (about 0.5 M) Concentrated HNO 3 Dilute, Known Solutions of Cu 2, Ni 2, Mg 2 and Zn 2 Safety and Waste Disposal Some of Group D cations and reagents used are toxic. The HCl, HNO 3, CH 3 COOH, NH 3 and NaOH are irritants. Avoid contact and wash immediately if any is spilled or splashed on you. Wear eye protection at all times. As you perform the experiment, collect all waste solutions in a waste beaker. This mixture should then be discarded in the appropriate waste container. DO NOT POUR ANY OF THE SOLUTIONS DOWN THE DRAIN. Unknowns and Knowns Separate known samples containing Cu 2, Ni 2, Mg 2 and Zn 2 are provided for your use. Testing known samples is helpful in this analysis since doing so will allow you to observe what a positive test looks like. It is usually convenient to test a known sample simultaneously with your unknown. To test a known sample, you can either prepare a known mixture of the Group D cations or the known Group D cation solutions can be tested separately. In the case of a known mixture, steps in the procedure that are required for separation of the cations must be followed before the confirmation tests can be conducted. If a known is prepared that contains only one of the cations, then the confirmation test can be conducted directly on the known. To prepare a known sample for testing, add 2 to 3 drops of the solution(s) containing the cation(s) you wish to test to about 0.5 ml of water. Note that the experimental conditions such as ph, oxidation state, etc. for the known test must be the same as that for the unknown. Outline of Procedure: Use a flow diagram in your notebook to record all observations for each step in the procedure. Chemistry and Relevant Background Information: 1. Removal of ammonia and oxalate: The presence of ammonia and possibly oxalate ions may interfere with later tests. These are removed from the sample solution by heating to dryness in the presence of concentrated nitric acid. The net ionic chemical equations for the reactions that take place were given in the Background Chemistry and Discussion section. You should rewrite them here for reference. 2. Identification of Nickel Ion: To test for Ni 2 a portion of the solution from step 1 is made Experimental Procedure: 1. CAUTION!!! The following must be done in the hood! (a) Pour the supernatant liquid containing the Group D cations into a clean crucible. Place the lid on the crucible. slightlv ajar, and heat the solution to dryness over a low flame. If your sample does not contain oxalate, skip to step 1(d). Otherwise continue with step 1(b). (b) Cool the crucible for 5 minutes, and then add 6 drops of concentrated nitric acid (Caution: HNO 3 is corrosive and an oxidizer), washing the inside of the crucible. Replace the lid and heat to drvness once again. (c) Cool the crucible for 5 minutes, repeat the addition of concentrated nitric acid and heat to dryness. (d) Cool the crucible for 5 minutes, and dissolve the residue by adding 5-10 drops of 6 M HCI. (It is alright if all of the solid does not dissolve.) Using a clean Pasteur pipet, transfer the solution to a clean, labeled test tube. (e) Rinse the crucible with 5 drops of deionized water. Add this rinse to the same test tube. 2. (a) Transfer one drop of the solution from step 1 to a clean Dr. L.J. Larson 3 June 28, 2008
Chemistry and Relevant Background Information: basic by adding NH 3. This is followed by addition of dimethylglyoxime solution. Formation of a cherry-red precipitate confirms the presence of Ni 2. In the space below, write the balanced net ionic chemical equation, including phase labels, for the formation of the cherry red precipitate (Refer to the Background Chemistry and Discussion section). Experimental Procedure: test tube. (b) Add 6 M NH 3, with stirring, until the solution tests basic to litmus paper. (c) Add 1 drop of 1% dimethylglyoxime. 3. Identification of Copper(II) Ion: To test for Cu 2 a portion of the solution from step 1 is made weakly acidic by adding NH 3. This is followed by addition of K 4 [Fe(CN) 6 ] solution. Formation of a maroon precipitate confirms the presence of Cu 2. CAUTION!!! If the solution is too acidic, toxic HCN(g) will be produced. In the space below, write the balanced net ionic chemical equation, including phase labels, for the formation of the maroon precipitate (Refer to the Background Chemistry and Discussion section). 3. CAUTION!!! The following must be done in the hood! (a) Transfer one drop of the solution from step 1 to a clean test tube. (b) Add 6 M NH 3 until the solution tests only weakly acidic. If the solution becomes basic, use 6 M acetic acid to make it slightly acidic. (c) Add 3-4 drops of 0.1 M K 4 [Fe(CN) 6 ], solution. 4. Separation of Cu 2 and Ni 2 from Mg 2 and Zn 2 Copper (II) and nickel ions are separated from magnesium and zinc by precipitating the Cu 2 and Ni 2 as sulfides. Magnesium ions do not form a sulfide and the Zn 2 ions remain in solution under the conditions imposed. The source of sulfide ion is the disproportionation of thiosulfate ion in hot, weakly acidic solution. Refer to the Background Chemistry and Discussion section for the unbalanced net ionic chemical equation for the disproportionation of thiosulfate ion. Balance this equation and write the balanced net ionic chemical equation, including phase label, in the space below: 4. (a) Add 6 M NH 3 to the remainder of the solution from step 1 until it tests slightly basic (ph 8-9). Then add 6 M acetic acid until the solution is weaklv acidic (ph 4-5). (b) Add about 0.2 g of solid sodium thiosulfate, Na 2 S 2 O 3, and heat for 5 minutes in a boiling water bath. Cool for 1 minute by swirling the test tube in cold tap water. (c) Centrifuge and decant the supernatant into a clean, labeled test tube. The precipitate may be discarded. 5. Identification of Magnesium Ion: The solution from step 4 is made basic by adding NH 3 and then sodium hydrogen phosphate is added. Under these consitions, Mg 2 will precipitate as MgNH 4 PO 4 and Zn 2 will precipitate as Zn 3 (PO 4 ) 2. Both of these are white solids. The net ionic chemical equations for the reactions that take place were given in the Background Chemistry and Discussion section. You should rewrite them here for reference. 5. (a) To the supernatant from step 4(c) add 6 M NH 3 until the solution tests slightly basic to litmus paper (ph 8). (b) Add 6-8 drops of saturated (about 0.5 M) Na 2 HPO 4 solution. Stir and then cool the solution in ice water for several minutes. (c) If a precipitate does not form, add 6 M HCl dropwise to lower the ph slightly. Watch carefully for precipitate formation; stop adding the HCl upon formation of a precipitate. Dr. L.J. Larson 4 June 28, 2008
Chemistry and Relevant Background Information: Experimental Procedure: (d) Centrifuge and discard the supernatant. To the precipitate formed, 6 M NaOH is added. Any Zn 3 (PO 4 ) 2 present will dissolve due to the formation of the [Zn(OH) 4 ] 2 complex ion. Any MgNH 4 PO 4, if present, will not dissolve, indicating the presence of Mg 2. In the space below write the balance net ionic chemical equation, including phase labels, for the reaction where Zn 3 (PO 4 ) 2 is dissolved in NaOH(aq) solution: (e) Wash the precipitate with 10 drops of deionized water. Discard the wash. (f) To the washed precipitate, add 6 drops of 6 M NaOH and stir thoroughly. Centrifuge and decant the supernatant into a clean, labeled test tube. 6. Identification of Zinc Ion: The supernatant from step 5(e) is tested for the presence of zinc ion. To test for Zn 2 the solution from is made weakly acidic by adding NH 3. This is followed by addition of K 4 [Fe(CN) 6 ] solution. Formation of a white precipitate confirms the presence of Zn 2. CAUTION!!! If the solution is too acidic, toxic HCN(g) will be produced. 6. CAUTION!!! The following must be done in the hood! (a) To the supernatant from step 5(f), add 6 M acetic acid until the solution tests only weakly acidic (ph 4-5). If the solution becomes too acidic, use 6 M NH 3 to make it slightly acidic. (b) Add 4 drops of 0.1 M K 4 [Fe(CN) 6 ], solution. In the space below, write the balanced net ionic chemical equation, including phase labels, for the formation of the precipitate (Refer to the Background Chemistry and Discussion section). Dr. L.J. Larson 5 June 28, 2008
Prelab Exercise: Name: 1. After adding concentrated HNO 3, heating to dryness and then dissolving the residue in 6 M HCl a Group D unknown is treated according to the procedure for the D analysis. For each step below, answer the questions asked. a) One drop of the solution is made basic with NH 3 and then 1% dimethoxylgloxime is added. A cherry red preciptate forms. What can you conclude? b) One drop of the solution is made weakly acidic by adding NH 3. When 0.1 M K 4 [Fe(CN) 6 ] is added a light green precipitate results. What can you conclude? c) The remainder of the solution is made weakly acidic by adding NH 3. When sodium thiosulfate is added and the solution is warmed, a black precipitate and a colorless supernatant results. What can you conclude? d) The supernatant from step c is made basic by adding ammonia. Adding 1 M Na 2 HPO 4 yields a white precipitate. What can you conclude? e) The precipitate from step d is treated with 6 M NaOH resulting in a white precipitate and a colorless supernatant. What can you conclude? f) The supernatant from step e is made weakly acidic by adding acetic acid. When 0.1 M K 4 [Fe(CN) 6 ] is added no precipitate forms. What can you conclude? 2. Describe a simple, ONE step test that would allow you to distinguish between Cu 2 and Zn 2. Dr. L.J. Larson 6 June 28, 2008
3. Complete the flow diagram below for Group D analysis. For each step, indicate the chemical form of each Group D cation that is present at the given step (for example CuO, Cu(NH 3 ) 4 2, CuS, etc. Add colors of solutions and precipitates where known. snt supernatant ppt precipitate heat Cu(NH 3 ) 4 2, Ni(NH 3 ) 6 2, Zn(NH 3 ) 4 2 and Mg 2 (1) to dryness (2) If oxalate present add Conc. HNO 3, to dryness 6M HCl 1 drop 6 M NH 3 until basic 1% DMG 1 drop 6 M NH 3 until slightly acidic 0.1 M K 4 [Fe(CN) 6 ] 6 M NH 3 until slightly basic then 6 M CH 3 COOH until slightly acidic (ph 4-5) 0.2 g Na 2 S 2 O 3, snt ppt 6 M NH 3 until ph 8 1 M Na 2 HPO 4 Discard snt 6 M NaOH ppt 6 M CH 3 COOH until slightly acidic 0.1 M K 4 [Fe(CN) 6 ] Dr. L.J. Larson 7 June 28, 2008
Follow-up Questions: For numerical problems, you must show all work for credit! Name: 1. The test for Cu 2 and Zn 2 ions involves adding K 4 [Fe(CN) 6 ] to a weakly acidic test solution. If the solution is too acidic, toxic HCN(g) will form. Write the balanced net ionic chemical equation, including phase labels, showing the production of HCN(g) when excess strong acid is added to [Fe(CN) 6 ] 4 (aq). 2. One of the reactions used in the Group D cation analysis is a disproportionation reaction. a) What is a disproportionation reaction? b) What was the purpose of using this reaction? In other words, what did it accomplish? c) Balance this disproportionation reaction and then write the balanced net ionic chemical equation, including phase labels, in the space below. 3. Cu(OH) 2, Zn(OH) 2 and Ni(OH) 2 are all insoluble in water. Explain why they do not precipitate with the Group B cations when the solution is made basic with ammonia. 4. In the Group D analysis scheme, Zn 3 (PO 4 ) 2 (s) is dissolved by adding 6 M NaOH. a) Write the balanced net ionic equation, including phase labels, for this reaction. b) Given that K sp for Zn 3 (PO 4 ) 2 (s) is 9.0x10 33 and that K f for [Zn(OH) 4 ] 2 is 4.6 x 10 17, calculated K net for the reaction written in part (a). c) Is this K net consistent with the fact that Zn 3 (PO 4 ) 2 (s) dissolves when 6 M NaOH is added? Why or why not? Dr. L.J. Larson 8 June 28, 2008