Qualitative Analysis I

Before the advent of modern instrumentation, a plethora of 'wet chemical' techniques had been devised to determine if a given element was present in a sample.  One classic example was the grizzled old prospector in the wild west.  He would take his sample of 'gold' or 'silver' ore to the assay office were they would determine if any precious elements were present (qualitative analysis).  If precious elements were found, then further testing (quantitative analysis) would determine how much of the element was present in the ore.   Another classic example (if you're a murder mystery fan) is the analysis of various body parts to determine if those lonely, elderly gentlemen actually died of 'yellow fever', or had they been helped along with small doses of arsenic or some other poison.

In both of these case, the most difficult and time consuming part of the analysis is the separation of the element(s) of interest from the rest of the sample (the so-called matrix).  For these wet chemical techniques, this normally means 'digesting' the sample with strong acid to dissolve the matrix and make the element(s) of interest soluble.  Then based on the general solubility guidelines below (remember them from the first semester?), various reagents are added to produce compounds that are insoluble.  Remember that if either the cation or the anion is considered soluble, then the whole compound is soluble.  A compound is only considered to be insoluble when both the cation and the anion portions are insoluble.  

The trick, of course, is to pick a set of conditions so that only a single element or a small subset of elements precipitate while leaving all of the other element(s) in solution.  To accomplish this we need to consider several variables:  concentration, pH, temperature, charge density, and health risks.  Luckily, over the last several hundred years, these variables have been studied extensively and fairly standardized schemes now exist to analyze almost any element in the presence of almost any other element(s).  Together these procedures are know as Qualitative Analysis.

Purpose:

The purpose of this experiment is to use Qualitative Analysis to positively determine the presence or absence of a specific group of ions in an unknown mixture, namely:  Na⁺, K⁺, NH₄⁺, Ag⁺, Cu²⁺, and Bi³⁺.  You will first have to separate a given ion from the rest of the mixture, and then perform a series of tests to confirm the identity of that ion.  As discussed above, most of the separation and identification procedures are based on manipulating the solubility of these ions.  The actual procedures for conducting this experiment are well understood.  The actual purpose of this lab is more an exercise in organizational/laboratory skills and attention to detail:  scrupulous cleaning of glassware, careful observation, labeling all containers, separating liquids and solids, etc.

Lab Tips:

Procedure:

During this experiment, you will be performing specific ion separation and confirmation procedures with two samples.  One will be a reference solution labeled "Qual I Known' and contains all of the ions of interest: Na⁺, K⁺, NH₄⁺, Ag⁺, Cu²⁺, and Bi³⁺.  The second sample will be a solution that contains any or all of these ions.  The most effective method for determining the ions in your unknown involves performing all of the procedures simultaneously on the known and unknown solutions.  This practice provides you with invaluable feedback on proper ion separation as well as both positive and negative confirmation test results.  NOTE:  It is very important to understand the chemical reactions that are taking place as you analyze your samples.  If you try to 'cookbook' these procedures, or skip the known solution, you WILL fail to properly identify the contents of your unknown.

Testing for presence of Na⁺:

Since the sodium ion is always soluble, there is no reagent we can add to form a precipitate.  The only way to test for the presence of the sodium ion is to perform a flame test.  This involves repeatedly cleaning a loop of Nichrome wire by dipping it into a 6 M HCl solution and then heating it in the inner cone of a Bunsen burner flame until there is no color produced from the wire.  Then the wire is dipped in the solution and heated in the flame.  If the sodium ion is present, a bright yellow flame will be produced.

1. Add approximately 2 mL of your original solution to a clean evaporating dish.
2. Add solid calcium hydroxide until the solution is basic to litmus paper.
3. Add a little extra calcium hydroxide and a pinch of ammonium carbonate to the solution.
4. Heat the solution until a moist residue remains.  DO NOT HEAT TO DRYNESS.
5. Allow the residue to cool and then add approximately 1 mL of water (distilled of course) to the residue.
6. Pour the solution into a small test tube and centrifuge.
7. Clean the Nichrome wire loop with 6 M HCl and heat it in the inner cone of a Bunsen burner until no color is seen.
8. Dip the wire into the supernatant and heat in the flame.
9. A bright yellow flame confirms the presence of sodium.

Like the sodium ion, the potassium ion is also soluble and so can not be precipitated.  Its presence is also determined by a flame test.  However, the yellow flame from sodium is so intense that it obscures the less intense lavender flame of potassium.  To block the sodium's yellow flame, a piece of cobalt blue glass is used to view the flame test.  Unlike the sodium flame which is persistent, the potassium flame only lasts a second or two..

1. Clean the Nichrome wire loop with 6 M HCl and heat it in the inner cone of a Bunsen burner until no color is seen.
2. Dip the wire into the same solution used to determine the presence of sodium and heat in the flame.
3. View the flame through a piece of cobalt blue glass.  A fleeting lavender flame confirms the presence of potassium.

Like sodium and potassium, the ammonium ion is also soluble and so can not be precipitated.  To confirm the presence of ammonium, we will take advantage of the equilibrium between ammonia (a weak base) and water:

Since this is an equilibrium reaction, if an excess of hydroxide ion is added, the reaction will shift to the left and produce more ammonia.  If the solution is also heated at the same time, the ammonia produced will evaporate from the solution as a gas.  While this gas has a characteristic odor, a more sensitive test is to use a piece of wet litmus paper.

1. Add approximately 3 mL of your original solution to a clean 100 mL beaker.
2. Heat the solution until one or two drops of solution remain.  DO NOT HEAT TO DRYNESS.
3. Allow the residue to cool and then add approximately 1-2 mL of distilled water to the residue.
4. Moisten a piece of red litmus paper with distilled water and attach it on the underneath (convex) side of a watch glass.
5. Add 1-2 mL of 6 M NaOH to the solution in the beaker.
6. Place the watch glass on the beaker with the litmus paper facing the bottom of the beaker.
7. Gently heat the beaker.  The presence of ammonium is confirmed if the red litmus turns blue.  Note:  be very careful not to allow any NaOH to contact the litmus paper.

According to the solubility guidelines, AgCl is insoluble.  Therefore if hydrochloric acid is added to a solution contains Ag⁺ ions, AgCl will precipitate:

However, we must be careful, because addition of too much HCl will result in the equilibrium formation of AgCl₂^- ion which is soluble:

To confirm that the precipitate formed by the addition of HCl is actually AgCl, the solid is mixed with ammonia.  In the presence of an aqueous ammonia solution, AgCl will dissolve to form the Ag(NH₃)₂⁺ ion:

However, this silver ammonia complex is unstable in acidic solution.  The addition of nitric acid (actually ANY acid BUT HCl) will react with the ammonia and force this equilibrium back toward the left.  This causes the AgCl to reprecipitate.

1. Add approximately 2 mL of your original solution to a small test tube.
2. Add 2-3 drops of 6 M HCl and stir.
3. Centrifuge the test tube to separate the precipitate from the supernatant.  Note:  the supernatant must be saved because it may contain Cu²⁺ and/or Bi³⁺ ions.
4. Now add 2-3 more drops of 6 M HCl to see if any further AgCl precipitates.  If more precipitate forms, repeat Steps #3 & #4.  If no more precipitate forms, proceed to Step #5.
5. Wash the precipitate with 5-10 drops of distilled water, stir, and centrifuge.  Add this supernatant to the supernatant from previous Step #3.
6. Add 6 M NH₃ drop wise to the precipitate until it completely dissolves.
7. Add 6 M HNO₃ drop wise until the solution is acidic to litmus paper.  A white precipitate confirms the presence of Ag⁺.

According to the solubility guidelines, CuS is insoluble.  Therefore if H₂S(aq) (hydrogen sulfide is the compound responsible for the odor of rotten eggs) is added to a solution containing Cu²⁺ ions, CuS will precipitate:

Back in the dark prehistoric days of chemistry (during the LAST CENTURY, when I took this course), the hydrogen sulfide was generated as a gas and then bubbled through the solution to precipitate the copper ion.  Everyone in the entire building knew when the first year chemistry students were performing Qualitative Analysis!  The nose is very sensitive to H₂S, you can smell less than 1 ppm of the gas.  However, if the level exceeds 100 ppm, the nose become desensitized (all of the receptors in the nose are occupied), and at levels in excess of 1000 ppm, H₂S is toxic.  However, since the nose is already desensitized at this level, you will have no warning of potential danger.  The best solution is to minimize the amount of  H₂S generated.

In these enlightened times, the hydrogen sulfide is generated in situ by the thermal hydrolysis of thioacetamide (CH₃CSNH₂):

Hydrogen sulfide is a weak, diprotic acid that reacts with water to produce the sulfide ion, S^2-:

To confirm the presence of copper, the precipitate must first be dissolved with hot nitric acid:

The presence of Cu²⁺ can be confirmed by two tests.  A solution containing Cu²⁺ is light blue in color.  If concentrated ammonia is added a deep blue copper ammonia complex is formed:

The second confirmation test is the addition of potassium hexacyanoferrate(II), which forms a reddish brown precipitate:

1. Add 10-15 drops of 0.1 M CH₃CSNH₂ to the supernatant you saved from the analysis of the Ag⁺ (Steps 3 & 4 above) and heat it in a boiling water bath for 10-15 minutes.
2. If nothing happens after 5 minutes of boiling, add an additional 10 - 15 drops of 0.1 M CH₃CSNH₂ to the solution and continue heating.
   
3. Let the solution cool and then centrifuge to isolate any black sulfides.
4. Test the solution for incomplete precipitation by repeating Steps 1 & 2.
5. Centrifuge, and discard the supernatant. NOTE:  if this is a combined Qual I and Qual II unknown (Lab Final), save this supernatant, since it contains your Qual II ions!
6. Wash the combined precipitate with 0.1 M HCl and stir.
7. Centrifuge, and discard the supernatant.
8. Repeat Steps 5 & 6.
9. Add 15 drops of 6 M HNO₃ to the precipitate.
10. Heat the test tube in the hot water bath until the precipitate has dissolved.
11. Centrifuge, and discard any elemental sulfur that may have formed.
12. Add 5-10 drops of concentrated ammonia to the supernatant.  The formation of a deep blue color confirms the presence of copper.
13. Centrifuge, and save the precipitate for bismuth testing.
14. Acidify the supernatant from Step 12 with 6 M acetic acid.
15. Add 3 drops of 0.2 M K₄[Fe(CN)₆]^4-, a reddish brown precipitate reconfirms the presence of copper.

According to the solubility guidelines, Bi₂S₃ is insoluble (as a matter of fact, the bismuth precipitated with the copper in the previous step).   Therefore if H₂S_(aq) is added to a solution containing Bi³⁺ ions, Bi₂S₃ will precipitate:

To confirm the presence of bismuth, the precipitate must first be dissolved with hot nitric acid:

The presence of Bi³⁺ is confirmed by two tests.  The addition of concentrated ammonia to a solution of bismuth ions produces a precipitate of white bismuth hydroxide:

The second confirmation test is the addition of freshly prepared sodium stannite to the solution containing the bismuth hydroxide precipitate.  The Bi³⁺ is reduced to black bismuth metal:

1. Prepare a fresh sodium stannite solution by adding 2 drops of 1 M SnCl₂ to a test tube.
2. Now add 6 M NaOH drop wise until the Sn(OH)₂ precipitate just dissolves.  Stir the solution to make sure all of the precipitate has dissolved.
3. Add several drops of this solution to the precipitate from Step 12 above.
4. The immediate formation of a black precipitate confirms the presence of Bi³⁺ in the solution.

Results/Calculations:

• Based on your results, draw a flow chart depicting the separation and confirmation the presence of all six ions.  Include the balanced chemical reactions, precipitates, and all observations.
• Based on your observations, what ions were present in your unknown?

(Updated 2/19/08 by C.R. Snelling)