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.
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.
|Sodium (Na+), Potassium (K+), Ammonium (NH4+)||No common exceptions|
|Fluorides (F-)||Insoluble: MgF2, CaF2, SrF2, BaF2, PbF2|
|Chlorides (Cl-)||Insoluble: AgCl, Hg2Cl2
Soluble in hot water: PbCl
|Bromides (Br-)||Insoluble: AgBr, Hg2Br2, PbBr2
Moderately soluble: HgBr2
|Iodides (I-)||Insoluble: many heavy-metal iodides|
|Sulfates (SO42-)||Insoluble: BaSO4, PbSO4, HgSO4
Moderately soluble: CaSO4, SrSO4, Ag2SO4
|Nitrates (NO3-), Nitrites (NO2-)||Moderately soluble: AgNO2|
|Chlorates (ClO3-), Perchlorates (ClO4-)||Moderately soluble: KClO4|
|Acetates (CH3CO2-)||Moderately soluble: AgCH3CO2|
|Sulfides (S2-)||Soluble: those containing NH4+, Na+, K+, Mg2+, Ca2+|
|Oxides (O2-), Hydroxides (OH-)||Soluble: Li2O, LiOH, Na2O, NaOH, K2O,
KOH, BaO, Ba(OH)2
Moderately soluble: CaO, Ca(OH)2, SrO, Sr(OH)2
|Carbonates (CO32-), Phosphates (PO43-), Arsenates (AsO43-)||Soluble: those containing NH4+, Na+, K+|
The trick, of course, is to pick a set of conditions so that only a single ion or a small subset of ions precipitate while leaving all of the other ions 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 ion in the presence of almost any other ions. Together these procedures are known as Qualitative Analysis.
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.
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+:
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
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 blue cone (hottest portion) of a Bunsen burner flame until there is no color produced from the wire. Then the wire is dipped into the solution and heated in the flame. If the sodium ion is present, a bright yellow/orange flame will be produced.
Like the sodium ion, the potassium ion is also soluble and so cannot be precipitated. Its presence is also determined by a flame test. However, the yellow/orange flame from sodium is so intense that it obscures the less intense lavender flame of potassium. To block the sodium's yellow/orange 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.
Like sodium and potassium, the ammonium ion is also soluble and so cannot be precipitated. To confirm the presence of ammonium, we will take advantage of the equilibrium between ammonia (a weak base) and water:
NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)
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 damp red 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:
Ag+(aq) + Cl-(aq) AgCl(s)
However, you must be careful, because addition of too much HCl will result in the equilibrium formation of AgCl2- ion which is soluble:
AgCl(s) + Cl-(aq) AgCl2-(aq)
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:
AgCl(s) + 2 NH3(aq) Ag(NH3)2+(aq) + 2 Cl-(aq)
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 re precipitate.
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:
Cu2+(aq) + H2S(aq) CuS(s)
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 and you can no longer smell it (all of the receptors in the nose are occupied), and at levels in excess of 1000 ppm, H2S is lethal. However, since the nose is already desensitized at this level, you will have no warning of the 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):
CH3CSNH2(aq) + 2 H2O(l) CH3CO2-(aq) + H2S(aq) + NH4+(aq)
Hydrogen sulfide is a weak, diprotic acid that reacts with water to produce the sulfide ion, S2-:
H2S(aq) + 2 H2O(l) 2 H3O+(aq) + S2-(aq)
To confirm the presence of copper, the precipitate must first be dissolved with hot nitric acid:
3 CuS(s) + 8 H+(aq) + 2 NO3-(aq) 3 Cu2+(aq) + 2 NO(g) + 3 S(s) + 4 H2O(l)
The presence of Cu2+ can be confirmed by two tests. A solution containing Cu2+ is light blue in color. If concentrated ammonium hydroxide is added a deep blue copper ammonia complex is formed:
Cu2+(aq) + 4 NH3(aq) Cu(NH3)4+(aq)
The second confirmation test is the addition of potassium hexacyanoferrate(II), which forms a reddish brown precipitate:
Cu(NH3)4+(aq) + [Fe(CN)6]4-(aq) Cu2[Fe(CN)6](s)
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:
Bi3+(aq) + H2S(aq) Bi2S3(s)
To confirm the presence of bismuth, the precipitate must first be dissolved with hot nitric acid:
Bi2S3(s) + 4 H+(aq) + NO3-(aq) 2 Bi3+(aq) + NO(g) + 3 S(s) + 2 H2O(l)
The presence of Bi3+ is confirmed by two tests. The addition of concentrated ammonium hydroxide to a solution of bismuth ions produces a precipitate of white bismuth hydroxide:
NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)
Bi3+(aq) + 3 OH-(aq) Bi(OH)3(s)
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:
2 Bi(OH)3(s) + 3 [Sn(OH)4]2-(aq) 2 Bi(s) + 3 [Sn(OH)6]2-(aq)
(Updated 7/13/13 by C.R. Snelling)