7# qpXXXXXf... :@z44x.Z z*X  OChemistry 142 QUALITATIVE ANALYSIS FOR CERTAIN CATIONS (Adapted from an experiment by Rod O'Connor and Warren C. Woelfel) Object This experiment illustrates the procedures used in developing a systematic scheme of separation and analysis for a selected group of cations. Introduction Chemists use inorganic qualitative analysis to detect the elements or groups of elements that are present in a sample of material. This can be done for electrolytes by putting the sample into aqueous solution and then testing the solution to prove the presence or absence of specific ions. In this experiment you will learn how this method can be used for the cations Ag+, Cu2+, Fe3+, Cr3+, Zn2+, and Ba2+. For the separation and detection of these cations, you will use reactions that involve different tendencies of these ions to form precipitates, to form complex ions, or to show amphoteric behavior. To familiarize yourself with the laboratory techniques involved and with the differences in chemical properties of the cations that allow their separation and detection, you will find it necessary to do some preliminary tests with each cation to discover its characteristic behavior. From the data obtained, you will devise a scheme for separating and identifying this group of cations. You will then apply this scheme to the analysis of an unknown solution containing one or more of these cations. Before coming to the laboratory, you should study this experiment and the references cited or others suggested by your instructor. In particular, you should balance the 24 equations given near the end of this write-up. These equations represent the reactions you may want to use in developing a scheme of analysis, and will be submitted as part of your report at the end of this experiment. References 1. J. Olmsted III, G. M. Williams, Chemistry, The Molecular Science, Mosby, St. Louis, MO, 1994, pp. 761-798. 2. W.E. Caldwell, G.B. King, Brief Course in Semimicro Qualitative Analysis, 3rd Ed., New York: D. Van Nostrand. 3. R.H. Petrucci, General Chemistry, New York: Macmillan, 1989, pp. 146-148 and 700-705. 4. J.C. Kotz, K.F. Purcell, Chemistry and Chemical Reactivity, Second Edition, Philadelphia: Saunders, 1991, pp. 797-806. Theory Sparingly Soluble Salts. The solubility of Fe3+ ions in basic solution is governed by the reaction: Fe(OH)3(s)  Fe3+ + 3OH-. (1) The equilibrium expression (law of mass action) is: Ksp = [Fe3+] [OH-]3. (2) The Ksp of 6.3x10-38 M4 shows that Fe(OH)3 is a very sparingly soluble salt under most conditions. Equation (2) can be solved for the solubility of Fe3+ ions in solution: solubility = [Fe3+] = f(Ksp, [OH-]3) (3) A plot of the solubility of Fe3+ versus [H+] is shown in Figure 1. For hydroxide concentrations greater than 10-11 M (pH> 3), Fe3+ is sparingly soluble. Therefore, in a qual scheme Fe3+ may be removed from solution by making the solution basic, precipitating Fe(OH)3. The precipitate can be removed by centrifugation. If a pure Fe(OH)3 precipitate is desired, we must also consider any other cations in solution. For example, if Zn2+ is also present, the following reaction and relationships also hold: Zn(OH)2(s)  Zn2+ + 2 OH-, (4) Ksp = [Zn2+] [OH-]2, (5) and solubility = [Zn2+] = F(Ksp,[OH-]2) . (6) A plot of the solubility of Zn2+ is also shown in Figure 1. At hydroxide concentrations greater than 10-6 M (pH>8), both Fe3+ and Zn2+ are insoluble and would coprecipitate as their hydroxides.  Figure 1. Solubility of Fe(OH)3 and Zn(OH)2 at various concentrations of H+. Using a weak base, it might be possible to adjust the pH to the range where Fe3+ will precipitate but Zn2+ will not. A buffer solution is a better approach to pH control. Acetic acid-sodium acetate buffers fall in this range (Figure 1). But in practice, this approach is difficult to use. An alternative approach to separating Fe3+ and Zn2+ is to use the amphoteric nature of Zn(OH)2 or the ability of Zn2+ to form complexes. Amphoteric Hydroxides. Some hydroxide precipitates dissolve in excess hydroxide solutions because of the formation of soluble hydroxo-complexes: Zn(OH)2(s) + 2 OH-  [Zn(OH)4]2-. (7) Such hydroxides are called amphoteric hydroxides because they will dissolve in both acid and base. The equilibrium constant for this reaction is very favorable, K = 2x1020M-1. Therefore, while Zn(OH)2 will form on addition of small amounts of a strong base to a Zn2+ solution, continued addition of strong base will cause the precipitate to dissolve. This amphoteric behavior is useful in qualitative analysis. If excess base is added to a mixture of Fe3+ and Zn2+ ions, Fe(OH)3 will precipitate and [Zn(OH)4]2- will remain in solution; thus a separation has been effected. Complex Formation. The formation of hydroxo-complexes of amphoteric hydroxides is useful for dissolving precipitates or for the prevention of precipitate formation. The formation of complexes other than hydroxo-complexes can also be useful. Ammonia is useful for forming complexes in qual schemes. For example, the reaction: Zn2+ + 4 NH3  [Zn(NH3)4]2+ (8) has an equilibrium constant of 2.9x109M-4. This shows that the formation of this complex is favorable. The usefulness of ammonia in qualitative analysis schemes can be shown by the behavior of a mixture of Fe3+ and Zn2+ ions. Remember that an ammonia solution is basic, since ammonia is a weak base: NH3 + H2O  NH4+ + OH- . (9) If concentrated ammonia is added to a solution containing Fe3+ and Zn2+ ions, Fe(OH)3 will precipitate (Equation 1) and [Zn(NH3)4]2+ will remain in solution; thus a separation has been effected. Oxidation-Reduction Reactions. Oxidation reactions can also play a role in qualitative analysis. In the list of ions for this experiment (Ag+, Cu2+, Fe3+, Cr3+, Zn2+, Ba2+), only Cr3+ is not in its maximum oxidation state in aqueous solution. Hydrogen peroxide is a good oxidizing agent, especially in basic solution: H2O2 + OH-  OOH- + H2O (10) The half reaction for its reduction is: OOH- + H2O+ 2 e- > 3 OH-. (11) Cr3+ is oxidized to soluble chromate ions, CrOA(2-,4) , in a basic solution containing hydrogen peroxide. Chromate ions are brightly colored, and therefore the oxidation of Cr3+ is a good test for chromium. Experimental Procedure Chemicals (in small dropper bottles): Knowns* Test Reagents 0.1 M AgNO3 3 M HCl 0.2 M Cu(NO3)2 3 M HNO3 0.2 M Fe(NO3)3 6 M NH3 0.2 M Cr(NO3)3 6 M NaOH 0.2 M Zn(NO3)2 3 M H2SO4 0.1 M Ba(NO3)2 *15 M NH3 *16 M HNO3 *3% H2O2 for confirmatory tests: 6 M acetic acid, CH3COOH *0.5 M K2CrO4 *0.2 M K4Fe(CN)6 or Na4Fe(CN)6 *0.2 M KSCN *available in front of room Apparatus and Techniques. Centrifuge. The centrifuge is used to speed up the separation of a precipitate from a liquid. When a mixture of solid and liquid is placed in a tube and rotated at high speed in a centrifuge, the more dense precipitate is forced to the bottom of the tube by a centrifugal force that is many times greater than the force of gravity. After centrifuging, the supernatant, or clear liquid above the precipitate, can easily be poured off or withdrawn with a capillary (Pasteur) pipet. The centrifuge may be damaged if allowed to run unbalanced. Therefore, before centrifugation, prepare another tube to balance the sample tube by filling it with water until the liquid levels in both tubes are the same. Insert the tubes in opposite positions in the centrifuge, then set the machine in motion. Usually, 30 seconds of centrifugation will effectively separate the solid from the liquid. Capillary Pipet. This device is convenient for removing and transferring liquids from precipitates in small test tubes or centrifuge tubes. To use the pipet, squeeze the air from the bulb before immersing the capillary end in the liquid so as not to disturb the precipitate. Stirring Rods. Stirring rods for qualitative analysis are much smaller than normally used in chemistry labs. For thorough washing of precipitates you must have good mixing of precipitate and wash liquid. Mixing is helped by tilting the test tube and moving the rod up and down without striking the bottom of the test tube; this method tends to spread the precipitate over a larger area. Mixing of Reagents. Always use a medicine dropper or capillary pipet when adding a small quantity of a liquid to a test tube or other vessel. One ml is about 20 drops. If the reagent bottle has a dropper, replace it promptly after use. Do not allow the dropper to touch the container or solution to which you are adding the dropper's contents. Do not set the dropper down on the desk top or another surface, and be sure to return it to the right bottle. If the bottle does not have a dropper, do not insert your own into it; instead, pour a small amount of reagent into a clean beaker, insert your droppper into that portion and discard the rest into a waste container. Be sure to clean your dropper before and after use. Always stir the mixture with a clean stirring rod after adding reagent. Precipitation. To detect the formation of a precipitate on mixing two solutions, it is essential that both solutions be initially clear; if necessary, centrifuge to clarify. A "clear" solution is transparent but not necessarily colorless. After adding a reagent to bring about precipitation, one should always test for complete precipitation if the purpose is to separate one substance from another. Suppose, for example, that a solution contains 0.10 millimole of Ba2+. We add to this solution a few drops of dilute H2SO4, and the amount happens to contain 0.08 millimole of SOA(2-,4). We have thus produced 0.08 millimole of solid BaSO4, but 0.02 millimole of Ba2+ remains in the supernatant liquid, because we did not add enough H2SO4 to precipitate all the Ba2+ as BaSO4. This fact can be discovered by centrifuging and adding another drop of reagent (dilute H2SO4 in this case) to the clear supernatant. If precipitation is complete, no additional precipitate will form. However, if insufficient reagent was added the first time, the additional drop will cause formation of more precipitate. If more precipitate is observed, add several more drops of reagent, centrifuge, and again test for completeness of precipitation. Repeat until no precipitate is formed on adding reagent. After the precipitate and supernatant are separated, the precipitate is washed by adding a few drops of the washing reagent (usually water), mixing thoroughly with a stirring rod, centrifuging, and removing the washings with a capillary pipet. Two or more washings are generally necessary to prevent contamination of the precipitate. Failure to wash precipitates is one of the most common sources of error in qualitative analysis. Transfer of Precipitates. The easiest way to transfer a residue from one container to another is to mix the residue with small amounts of washing liquid and pour the suspension quickly into the new container. Repeat this two or three times to get complete transfer. For small volumes, draw the suspension into a pipet for transfer. Heating of Solutions. To avoid excessive evaporation on prolonged heating, use a water bath. Always stir when heating test tubes. Water Bath. Fill a 250 or 400 ml beaker three-quarters full of water and set it on a hot plate. Fit the aluminum test tube holder into the beaker. Heat to a gentle boil. Adjusting Acidity. Always stir well when adding acid or base. Use a stirring rod and test a drop of the solution on litmus paper placed on a watch glass. Never dip the paper into the solution. Week One In this part, you will study the properties of the individual cations Ag+, Cu2+, Fe3+, Cr3+, Zn2+, and Ba2+. Record all results (precipitates, colors, etc.) in your notebook. A typical layout for your notebook is given in appendix 1. (1) Precipitation of Chlorides with Dilute HCl. To 10 drops of cation solution, add 1 ml of 3 M HCl. Caution: Never treat chlorides with strong oxidizing agents. Chlorine gas may be produced. (2) Precipitation of Hydroxides with Dilute NaOH or NH3 Solution. The purpose of this step is to test the properties of each cation in dilute base. You will produce hydroxide precipitates of most of the cations; save these precipitates for the tests in Step 3. To 10 drops of cation solution, add 15 drops of water and 2 drops of 6 M NaOH solution. If no precipitate forms, test the solution with litmus paper. If the solution is still acidic, add 6 M NaOH until the solution is just basic. (Addition of too much NaOH may cause the amphoteric hydroxides to redissolve.) To 10 drops of cation solution, add 15 drops of water and 2 drops of 6 M NH3 solution, and mix. If no precipitate forms, test the solution with litmus paper. If the solution is still acidic, add more 6 M NH3 until the solution is just basic, then add two drops in excess. (Addition of too much NH3 may cause the cations that form NH3 complexes to redissolve.) At this point you should have hydroxide precipitates for most of the cations. Keep only one of the precipitates if you obtained precipitates upon addition of both NaOH and NH3. Centrifuge the precipitates, discard the supernatant solutions and proceed to the next step. (3) Amphoteric Nature of the Hydroxides: Excess NaOH. Amphoteric hydroxides will dissolve in excess base. To the hydroxide precipitates from Step 2, add 6 drops of 6 M NaOH solution. Stir carefully for 1 minute. Observe to see if the precipitate dissolves. (4) Ability to Form Coordination Complexes with NH3: Excess NH3. Test 10 drops of the cation solution with litmus paper. If the solution is acidic, add enough 15 M NH3 to make the solution just basic. Then add an additional 20 drops of 15 M NH3, with stirring. If the cation solution is initially neutral to litmus, add 20 drops of 15 M NH3, with stirring. Note any changes. (5) Precipitation of Sulfates with Dilute Sulfuric Acid. To 10 drops of the cation solution, add 5 drops of 3 M H2SO4. (6) Oxidation with Alkaline Hydrogen Peroxide. To 10 drops of the cation solution, add, drop-by-drop, 6 M NaOH until the solution is basic. Add another 6 drops of 6 M NaOH, then 10 drops of 3% H2O2. Heat for 3 minutes in a beaker of boiling water, with vigorous stirring. Note the results that differ from Steps 2 and 3. (In this step some of the precipitates change color but do not dissolve. This color change is probably due to the thermal decomposition of the hydroxide to an oxide.) NOTE: Before leaving the lab, submit a clean, corked, 13x100mm test tube clearly labelled with your name and lab day. Week Two Based on your results in Week One, devise a scheme for separating and identifying the individual cations studied. The scheme should take advantage of the solubility differences among the salts, hydroxides, and complexes, and redox properties of the several metal cations. You should also consider confirmatory tests as possible separation steps. If possible, use a reagent that causes only one cation to precipitate, then remove the insoluble compound by centrifugation. It may be necessary to precipitate two or more cations together. After separation, the mixed precipitate can be redissolved and the cations then separated by some other reagents, or it may be possible to dissolve only one component of a mixed precipitate by the addition of a suitable reagent. Most students feel more confident of their results if confirmatory tests are carried out after a cation has been separated from the solution. The procedures for several confirmatory tests are given in appendix 2. You should feel free to include these tests in your qual scheme. Your task is to select a scheme of separation that is straightforward, that entails the fewest steps, and that achieves the cleanest separations. Make a flow diagram of your scheme (in your lab notebook), including test reagent amounts and concentrations. A copy of your scheme is due at the beginning of the discussion section during the second week of this experiment. This scheme will be graded and incorporated into the report grade for this experiment. A sample separation scheme for a different set of cations and test reagents is shown in Figure 2. During the laboratory you should test your scheme on known mixtures of the cations to make sure it works. You should also try some of the tests on a spot plate, instead of in test tubes. Initial testing of your unknown with spot plate tests can quickly rule out some cations, so that the steps that are designed to separate these cations can be droppped from your scheme. You should also try the confirmatory tests. Make sure to include an up-to-date separation scheme in your lab notebook. Include amounts and concentrations of test reagents in your final scheme.  Week Three As an unknown for you to analyze according to your proposed scheme, your instructor will give you a mixture of one or more cations from the group you have studied. In working on your unknown, START WITH 10 DROPS of your unknown and use about 5-10 drops of most reagents. Be sure you do not need the supernatant before discarding it. Also, make certain that the total volumes do not exceed about 3 ml. Always test for complete precipitation; this precaution is especially important here, since no one has previously worked on your scheme to determine what quantities of reagents are needed. Simplify your scheme as you go along. If you know that a certain cation is absent, either from spot plate tests or early steps in your scheme, then the steps involving that cation can be eliminated from your scheme. For example, if a spot plate test shows that Ag+ is absent, don't bother to add HCl to your unknown. In your lab notebook, keep track of all your observations, including the amounts and concentrations of the test reagents you used. Sometimes the unknown solution contains acid. When a basic reagent, such as NH3 or NaOH, is added to an acidic solution, no result is to be expected until enough reagent has been added to neutralize the acid. Test with litmus paper to be sure that you have made the solution basic before you draw any conclusions. Report Hand in a list of cations that you know to be present in your unknown and a list of the cations for which you tested but were not present. Also, in your notebook, balance all of the equations on page 13 of this write-up; hand in a copy of these balanced equations as part of your report. After you have completed your analysis and submitted your report, you are strongly encouraged to verify your results by analyzing your unknown using our inductively coupled plasma emission spectrometer (ICP). This is a rapid and extremely sensitive way to detect the metal ions in your sample. Please read the following discussion of emission spectroscopy and sample preparation; an instructor or student assistant will be available to help you run your sample. Emission Spectroscopy Emission spectroscopy has long been a sensitive and selective technique for analysis of metals. Flame emission spectroscopy of Na+ and K+ is regularly used in clinical and environmental laboratories. Geochemistry laboratories have long used high current arc sources, which are similar to searchlight and theatrical spotlight systems, for emission spectroscopy of mineral samples. A simple example of emission spectroscopy is lighting a match. The yellow glow is emission from excited Na ions in the flame. In emission spectroscopy a sample is heated in a flame or electrical discharge to cause it to emit light. The emission frequencies are characteristic of the elements present in a sample and the intensity of the emission at each frequency is proportional to concentration, over a wide range of concentrations. Conventional flame or arc sources suffer from a number of interferences and instabilities that limit the usefulness of emission determinations. The interferences, such as oxide formation, can largely be eliminated by choosing higher temperature sources. The instability problem is solved by choosing a plasma that is heated by radio frequency or microwave energy. A plasma is an electrically conducting gas composed of free electrons and positive ions. Examples of plasmas include conventional flames, neon signs, fluorescent light fixtures, and the ionosphere of the earth. Plasmas are often considered a separate state of matter distinct from solids, liquids, and gases. Plasmas require a constant input of energy for existence. The source of the energy can be chemical reactions, as in a flame, an electrical discharge, as in a neon sign, or radiant energy, as in a radio frequency or microwave supported plasma. The gas chosen for the plasma is the most important determining factor for choosing the energy source. The most generally used gas for analytical plasmas is argon. Argon plasmas are very hot, stable, and provide a minimum of background emission. Argon plasmas are started with an electrical discharge, but in emission spectroscopy the energy for heating the plasma is provided by a 1 kWatt radiofrequency (R.F.) source, generally at a frequency of 12 MHz. This frequency lies between the AM and FM bands on your radio dial. The R.F. energy is coupled into the plasma by wrapping a coil of wire (essentially a transmitter antenna) around the plasma. A coil of wire acts as an inductor, hence the name for this emission source, inductively coupled plasma, or ICP. ICP sources are easy to use, very hot (around 10,000oK), stable, and provide relatively low background emission. The sample is introduced into the flowing argon stream as an aerosol from a nebulizer. The light emitted from the plasma is analyzed by a very high resolution monochrometer and a sensitive photomultiplier detector. A diagram of the instrument is shown in Figure 3.  Figure 3. ICP optical layout The analytical wavelength to use is determined by introducing a known sample of the desired element and scanning the monochrometer to determine the emission spectrum (In our instrument the exit slit is scanned to determine spectra). The wavelength of maximum emission is then used in subsequent measurements. The computer system on the instrument records the monochrometer settings for this emission wavelength in a "Peak Table" to which it can return whenever determinations of this particular element are desired. ICP emission spectrometers are calibrated by introducing a series of solutions of known concentration and fitting the emission intensities to a linear or polynomial calibration function. This is exactly the same process as is used for a visible absorption spectrophotometer (such as Spectronic 20's) except that emission intensity is used rather than absorbance. We will be able to use a linear calibration function in this laboratory. This is shown schematically in Figure 4. The solution of unknown concentration is then introduced into the plasma and the resulting emission intensity is then used to calculate the concentration using the calibration curve.  Figure 4. ICP calibration curve for three standard solutions. Sample Preparation Place three drops of your unknown in a clean test tube. Then fill the test tube half-way with deionized water. Stir well. Take your sample to the Instrument Lab where the student assistant will help you analyze your unknown. Make sure to note the position of your test tube in the auto sampler rack. The results are printed in millimolar units, where 1mM = 1x10-3M. Qualitative Analysis for Certain Cations Appendix 1. Record all your observations in your notebook. Include the formation (and dissolution) of precipitates, along with the colors of precipitates and solutions. Colors are very helpful. An example of the type of table you might use is given below. ReagentsReactions of CationsAg+Cu2+Fe3+Cr3+Zn2+Ba2+HCldilute NaOHdilute NH3excess NaOHexcess NH3H2SO4H2O2 + excess NaOH Appendix 2. CONFIRMATORY TESTS FOR QUALITATIVE ANALYSIS FOR CERTAIN CATIONS After the ions have been separated in your qual scheme, confirmatory tests should be carried out to assure their proper identification. Zinc Ferrocyanide, Fe(CN)64-, properly called hexacyanoferrate(II), when added to a slightly acidic solution containing zinc ions produces a whitish precipitate which is quite satisfactory as a test for zinc. Check your solution with litmus to make sure it is acidic. If it is not, add 6 M acetic acid (CH3COOH), dropwise, until it is barely acidic. Then add five drops of 0.2 M K4Fe(CN)6 and stir. If a white or off-white precipitate forms, the presence of zinc is confirmed. Copper Copper ions can be detected in very dilute solutions by forming Cu2Fe(CN)6, which is red. To the solution, which should contain excess NH3, add 6 M acetic acid (CH3COOH) a few drops at a time, testing with litmus, until the solution is just acidic. Then add two drops of 0.2 M potassium ferrocyanide solution, K4Fe(CN)6. Mix well. A brick red precipitate indicates that copper is present. Chromium Chromium is usually oxidized to chromate in qual schemes. The yellow color of chromate is very sensitive, but not completely conclusive. To the yellow solution add 6 M acetic acid until the solution is barely acid. Then add ten drops of the Ba2+ test ion solution. Stir the mixture. The precipitate should be yellow if CrO42- is present. However, if you added excess sulfate earlier, some BaSO4 will also form. BaSO4 is white. To avoid this problem, if you do not have Ba2+ in your unknown, it is not necessary to perform the separation step using sulfuric acid. (This test can also be used in your qual separation scheme, if you need to precipitate CrO42- from solution. Just remember this test doesn't work with Cr3+; you must oxidize Cr3+ to CrO42- first.) Iron After the addition of excess base, the formation of a reddish or brownish precipitate indicates Fe(OH)3. Centrifuge and remove the supernatant. Wash the precipitate with ten drops of distilled water and discard the washing after centrifuging again. To the precipitate of Fe(OH)3 add 1 ml (twenty drops) of 3 M H2SO4 solution. If the precipitate does not dissolve at once, heat it in the water bath until it does. Cool and add four drops of 0.2 M potassium thiocyanate (KSCN) solution. If a deep red solution is produced, the presence of iron(III) is confirmed. A pink color is not significant, since there is usually a trace of iron in the reagents used, sufficient to produce a pink color. QUALITATIVE ANALYSIS FOR CERTAIN CATIONS POSSIBLE SEPARATION REACTIONS (1) Ag+ + Cl- > AgCl (2) Ag+ + NH3 + H2O > Ag2O + NH4+ (3) Ag2O + NH3 + H2O > Ag(NH3)2+ + OH- (4) Ag+ + OH- > Ag2O + H2O (5) Cu2+ + NH3 + H2O > Cu(OH)2 + NH4+ (6) Cu(OH)2 + NH3 > Cu(NH3)42+ + OH- (7) Cu2+ + OH- > Cu(OH)2 (8) Fe3+ + NH3 + H2O > Fe(OH)3 + NH4+ (9) Fe3+ + OH- > Fe(OH)3 (10) Cr3+ + NH3 + H2O > Cr(OH)3 + NH4+ (11) Cr3+ + OH- > Cr(OH)3 (12) Cr(OH)3 + OH- > Cr(OH)4- (13) Cr(OH)4- +H2O2 + OH- > CrO42- +H2O (14) Zn2+ + NH3 + H2O > Zn(OH)2 + NH4+ (15) Zn(OH)2 + NH3 > Zn(NH3)42++ OH- (16) Zn2+ + OH- > Zn(OH)2 (17) Zn(OH)2 + OH- > Zn(OH)42- (18) Ba2+ + SO42- > BaSO4 POSSIBLE CONFIRMATORY TESTS (19) Ag(NH3)2+ + H+ + Cl- > AgCl + NH4+ (20) Cu2+ + Fe(CN)64- > Cu2Fe(CN)6 (21) Fe3+ + SCN- > FeSCN2+ (22) CrO42- + Ba2+ > BaCrO4 (23) Zn2+ + K+ + Fe(CN)64- > Zn3K2[Fe(CN)6]2 (24) Ba2+ + SO42- > BaSO4 Chemistry 142 QUALITATIVE ANALYSIS FOR CERTAIN CATIONS PRELABORATORY ASSIGNMENT WEEK ONE NAME (1) How do you test for complete precipitation? (2) A test solution was made from 10 drops of 0.1 M AgNO3 test ion solution and 30 drops of water. Calculate the number of drops of 3 M HCl that would be needed to supply the same number of moles of Cl- as there are moles of Ag+ in this solution. (In practice, an excess of HCl would be used to ensure complete precipitation). Assume 1 ml=20 drops. (3) One drop of 6 M NaOH is added to a neutral test solution of total volume 3 ml. (a) Calculate the hydroxide ion concentration and the pH of the solution. (Remember that 20 drops = 1ml). 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()S ddPro( +:(5 centrifuge ddPro( D5Sg(N8 Precipitate ddPro( P5_l* PbS (black) ddPro( \5kr* CdS (yellow) ddPro( t5R*HNO ddPro  yJV +3 ddPro( tNo (~Q, boil, ddPro( HbW(Re Supernatant ddPro( Wbft*Ni ddPro  Ul`} (]o2+ ddPro( Wtf +, Ca ddPro  U` (]2+ ddPro( hUwk (rXNH ddPro  mcxo +4 ddPro( hiw (rlCl, NH ddPro  mx +3 ddPro( hw (r, ddPro( t_x(~b(NH ddPro  yp| +4 ddPro( tt (~w) ddPro  yw +2 ddPro( t| (~S, ddPro( b(e centrifuge ddPro( :i(= centrifuge ddPro( #U+ Precipitate ddPro( #Y* NiS (black) ddPro( !U($HCl, HNO ddPro  MY +,3 ddPro( Q\ (T, ddPro( 2J(5boil ddPro( 2D*Ni ddPro  <M (?2+ ddPro( 0F (3NH ddPro  >L +3, ddPro( 1 ([(CH ddPro  )5 +3 ddPro( -8 (0) ddPro  0< +2 ddPro( 4C (7C ddPro  ;G +2 ddPro( ?b (B(NOH) ddPro  Zf +2 ddPro( ^i (a] ddPro(   9(#NiC ddPro  1 = +8 ddPro( 5 D (8H ddPro  < L +14 ddPro( D S (GN ddPro  K W +4 ddPro( O ^ (RO ddPro  V b +4 ddPro(  d (!nickel dimethyl ddPro( ( X+  gloximate ddPro( (a(#"(strawberry red ddPro( #(2\+ precipitate) ddPro( -!<Y(7$ confirms Ni ddPro  +Q6b (3T2+rrr @2===     q Q*@ (*@ )Q) ɫ @گ             躪           * .j       j2=z=z=z @@      *   @j@ @ @      * @  j @ ! 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