Introduction

Qualitative analysis refers to the methods used in determining the identity of a chemical compound as opposed to its amount. For example, a majority of this semester has been attributed to determining how much of a particular analyte is present through a variety of quantitative techniques. However, in this experiment we switch things up a little. Instead of concentrating upon how much of something is present, you will now only be concerned with what is present. Overall, the qualitative procedure uses known reactions of a given chemical species and interprets the obtained results using deductive reasoning. In other words, there are a variety of chemical analyses existing for different elements or types of a compound, and combining this with your understanding of acid-base equilibria, redox reactions, and solubility, it becomes fairly elementary to determine what particular analyte is found in an unknown mixture.

To increase the accuracy of the analysis process, the procedure for a typical qualitative assay generally breaks down several collections of unknown ions into groups of ions that follow the same chemical pattern. In fact, the groups are formed from the basis of the reactivity of the ions, not their groups in the periodic table! Similar ions are then placed in the same group and are then identified by use of confirmatory reactions.

Identifying Anions and Cations

Qualitative analysis is used to separate and identify the cations and anions present in an unknown chemical mixture. Ions are grouped according to their reactivity to specific compounds. First, ions are removed in groups from the initial aqueous solution. After each group has been separated, then testing is conducted for the individual ions in each group. What follows here is a common grouping of cations:

Group I cations are those which form insoluble chlorides when reacted with dilute HCl solutions:

Group I cations include Silver (Ag⁺),Mercury I (Hg₂²⁺),and Lead II (Pb²⁺) and are precipitated out by addition of 1 M HCl to the unknown cation mixture. The Group I cations form the precipitates AgCl, Hg₂Cl₂, and PbCl₂ that can then be removed from the mixture by centrifugation.

Group II cations, while soluble in dilute HCl, are insoluble when reacted with a dilute H₂S solution at low (acidified) pH <1:

Group II cations include Bismuth III (Bi³⁺),Cadmium II (Cd²⁺), Copper II (Cu²⁺), Antimony III (Sb³⁺), Antimony V (Sb⁵⁺), Tin II (Sn²⁺) and Tin IV (Sn⁴⁺). The Group II cations are precipitated out by reaction with 0.1 M H₂S at a low pH, normally around 0.5. The precipitates that are formed are also removed from the remaining mixture by centrifugation.

Group III cations are cations that are soluble in acidic H₂S but are insoluble in basic H₂S.

Group III cations which include Aluminum (Al³⁺), Chromium III (Cr³⁺), Iron II (Fe²⁺), Iron III (Fe³⁺), Manganese II (Mn²⁺), and Zinc II (Zn²⁺ ) are precipitated out of the mixture by changing the pH of the previous solution (already saturated with 0.1 M H₂S) to a pH of 9 or greater.

All that is left at this point in the unknown solution should be the Group IV and V cations. These are the alkali and alkaline earth metal cations and the ammonium cation. These cations can be separated from each other by their varying reaction to carbonate solution.

Group IV and V cations include Barium (Ba²⁺), Calcium (Ca²⁺), and Magnesium (Mg²⁺) and Potassium (K⁺), Sodium (Na⁺), and Ammonium (NH⁴⁺) respectively. The alkaline earth metal cations Ba²⁺, Ca²⁺, and Sr ²⁺ are precipitated in 0.2 M (NH₄)₂CO₃ solution at pH 10 and then removed from the others by centrifugation. The alkali metal cations K⁺ and Na⁺ are soluble in most all solutions and thus cannot be separated by precipitation. Rather their presence is confirmed by a colorimetric flame test. The ammonium ion is also soluble and may also have been lost through earlier tests as ammonia. In order to test for the presence of ammonium you should take a small amount of the original mixture and add NaOH until the mixture becomes basic. At that point, the smell of ammonia (NH₃) will confirm the presence of ammonium ion (NH⁴⁺).

The portion of qualitative analysis described above simply separates the groups from each other. In order to determine the presence of individual ions in each group a variety of other reactions must be performed. Each reaction is specifically designed to react strictly with only one cation in each group thereby confirming only its presence. As with the main group separations, the reactions with each cation rely on those chemical properties which are unique to that ion. The discussion of the separation and confirmation of those ions is quite extensive and best broken up into several parts. The links below connect to detailed pages regarding the reactions used to confirm the presence of each cation above listed by group:

1. Group I Cations
2. Group II Cations
3. Group III Cations
4. Group IV Cations
5. Group V Cations

Anions

As with cations, the anions in a mixture are similarly determined using confirmation reactions based on the unique chemical reactivity of the anion. The most common anions found are: Carbonate (CO₃^2-), Sulfate and Sufite(SO₄^2-, SO₃^2-), Phosphate (PO₄^3-), the halogens (Cl^-, Br^- and I^-), Nitrate and Nitrite (NO₃^-, NO₂^-), acetate (CH₃COO^-), Chromate and Dichromate (CrO₄^2-, Cr₂O₇^2-). A list of the reactions used to confirm the presence of each of these anions can be found here.

The Experiment: As was mentioned in the purpose, qualitative analysis today is predominantly performed by machines and is thus somewhat historical in nature. As such, we have decided not to perform experiments running for weeks in order to cover all of the separations described above. Rather, the experiment you will be performing will only take a single lab period but should give you experience in the types of reactions and techniques you would typically perform during a qualitative analysis scheme.

In your experiment you will be investigating seven cations (Na⁺, Mg²⁺, Ni²⁺, Cr³⁺, Zn²⁺, Ag⁺, and Pb²⁺) as well as four anions (NO₃^-, Cl^-, I^- and SO₄^2-). You will use the reactivities of these ions to both separate and then identify them. The first part of this experiment asks you to investigate the cations and anions with the goal of gathering enough information about their reactivity and physical properties to use that data as a reference guide later in the experiment when given an unknown mixture of some of the same ions. Thus it is extremely important that you take meticulous notes with respect to every aspect of a solution or precipitates appearance. For example, don't just write down 'a white precipitate', comment on the structure of the precipitate, e.g. is it flaky or crystaline or powdery etc. as there may be several white precipitates in the experiment before you are through.

Part I: Reactions of Cations with NaOH and NH₃:

Most hydroxide salts are insoluble and thus precipitate out of solution when formed by reaction. The precipitates that form from these reactions often have distinctive colors or textures lending to their identification. Some hydroxide salts undergo further reaction to form oxides. For example:

If an excess of hydroxide is added, however, there are some hydroxide salts that will exhibit amphoteric character and re-dissolve, forming complex ions.

For example, aluminum in dilute hydroxide solution forms a white gelatinous powder, Al(OH)₃. But if more hydroxide is added the aluminum hydroxide dissolves and the soluble complex ion Al(OH)₄^- forms instead.

The simplest way to determine the structure of a complex ion is to look it up in a reference book or your text (Appendix K). In this experiment, all of the complex ions containing OH^- will contain four OH^- ions. This generalization should be used strictly for this experiment. The overall charge on the complex ion will therefore be the charge on the metal cation minus the number of OH^- ions.

In addtion to Aluminum, you will also be looking at the reaction of hydroxide with aqueous ammonia, NH₃.

NH₃ solutions are basic enough to create insoluble hydroxides, but do not supply an excess of OH- ions in the solution. If you add a small amuont of ammonia, insoluble hydroxides will precipitate.

Ammonia can also form complex ions with some metals. These complex ions are soluble and will generally have four NH₃ ligands bound to the metal, Ex. Zn(NH₃)₄²⁺. Silver is the only exception to the 4 ligand rule for this experiment, it has only 2 ammonia groups in its complex, Ag(NH₃)₂⁺.

A Stability Sequence

A stability sequence is a sequence of salts and ion for a given cation ranked on their stability from more stable to less stable . In the following series the most stable salts are the ones farthest to the right. While you can perform reactions to make more stable salts (left to right), you cannot go in the opposite direction.

For Ag⁺: Ag₂O(s) < AgCl(s), Ag(NH₃)₂⁺(aq) < AgI(s)

For Pb²⁺: PbCl₂(s) < PbSO₄(s) < PbI₂(s) < Pb(OH)₂(s) < Pb(OH)₄^2-(aq) in xs OH^-

For Zn²⁺: Zn(OH)₂ < Zn(NH₃)₄2+(aq) < Zn(OH)₄^2-(aq) in xs OH

For Ni²⁺: Ni(OH)₂(s) < Ni(NH₃)₄²⁺(aq) < Ni(OH)₂(s) in xs OH-

What these sequences indicate is that if, for example, you were to add aqueous sulfate ions to solid lead chloride, PbSO₄ would form. You could not, however, go in the other direction. But you could add iodide to lead sulfate or lead chloride and get lead iodide to form. Any species above shown as an ion is soluble in water. The neutral compounds are all solids and will precipitate.

Anion Precipitation Reactions

In this experiment only two cations will form precipitates when combined with the anions we are using. Silver forms two precipitates and lead forms 3 precipitates. You will need to test these anions (NO₃^1-, Cl^1-, I^1-, and SO₄^2-) with each of the two reactant cations and take careful notes to allow you to determine the identity of the anions in your unknown.