Chapter 14. Qualitative Determination of Selected Metal Cations

Objectives

Laboratory 10
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  • To investigate the chemical reactivity of different metal cations in different types of chemical reactions: precipitation, oxidation–reduction, and complex ion formation.
  • To make and record observations of chemical reactions with known metal cations, and then repeat these procedures to identify their presence in unknown samples.
  • To report the procedures, observations, and conclusions in a concise and convincing manner.

Introduction

In this experiment you are the primary investigator at an art crime scene. An eccentric art collector has offered $5 million for Johannes Vermeer’s 1656 painting, “A Maid Asleep.” Not surprisingly, many people have come forward claiming they have the original painting. Authorities have narrowed it down to six paintings, all of which look identical. Several of these paintings must be forgeries!

The only way to positively identify the actual Vermeer painting is through chemical analysis of the pigments used to create the different works of art. Six different blue pigment samples have been obtained, one from each of the six paintings. All of the pigment samples have a similar blue hue, so the actual sample from “A Maid Asleep” is difficult to distinguish from the others by visual identification.

Figure 10.1 Vermeer's "A Maid Asleep."1

The blue pigment samples have been dissolved into six different aqueous solutions. The resulting solutions contain a mixture of cations originating in the pigments used in the paint sample. You will be analyzing one of these samples. Your task is to analyze the unknown solution and identify the metal cations present. This will be accomplished by 1) performing a series of experiments and making observations with samples of known metal cations, then 2) repeating these procedures with your unknown sample.

1 Johannes Vermeer (Dutch, 1632–1675) The Metropolitan Museum of Art, New York. Bequest of Benjamin Altman.

Background information on Chemical Tests

Reactions of metal ions in aqueous solution are generally of three types: precipitation, complex ion formation, and oxidation–reduction.

Precipitation, or formation of an insoluble product, is an important driving force for a metathesis reaction. An example of a precipitation reaction occurs when a solution of silver nitrate is added to a solution containing sodium chloride.

AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)

This is called the molecular equation because all formulas are written as though the compounds are molecules, even though all compounds are in fact ionic in this equation. Molecular equations are useful when doing stoichiometry calculations, where the complete formula must be used.

The ionic equation uses formulas that reflect the form of each substance in the solution. The reaction between solutions of silver nitrate and sodium chloride would be represented as follows.

Ag+(aq) + NO3–(aq) + Na+(aq) + Cl–(aq) → AgCl(s) + Na+(aq) + NO3–(aq)

Both reactants and the sodium nitrate product are soluble ionic compounds that exist in solution as ions surrounded by layers of water molecules. The silver chloride product is an insoluble ionic compound; it precipitates from the solution. The ions are closely associated in the solid as indicated by writing them together in the formula. Observation of formation of the precipitate is evidence that the reaction has occurred.

If you look carefully at the ionic equation, you will notice that Na+(aq) and NO3–(aq) appear as both reactants and products; they did not undergo any chemical change. Ions that do not participate in the reaction are called spectator ions and can be canceled, producing the following net ionic equation.

Ag+(aq) + Cl–(aq) → AgCl(s)

The net ionic equation contains only those ions that are involved in chemical reaction and has the advantage of focusing attention on the change that is taking place.

Consider combining solutions that contain potassium nitrate and sodium chloride. The molecular equation again represents a metathesis reaction:

KNO3(aq) + NaCl(aq) → KCl(aq) + NaNO3(aq)

The ionic equation is:

K+(aq) + NO3–(aq) + Na+(aq) + Cl–(aq) → K+(aq) + Cl–(aq) + Na+(aq) + NO3–(aq)

If we cancel spectator ions here, there is nothing left. In other words, the result of combining these two solutions produces simply a mixture of ions; there is no chemical reaction:

KNO3(aq) + NaCl(aq) → N.R. (no reaction)

There is no driving force, since both products are soluble. The observation is a clear, colorless solution.

Another possibility for reaction in solution is complex ion formation. A complex ion contains a central metal ion surrounded by and bonded to a number of ions or groups. The whole collection has an overall charge. Two complexing agents you will encounter in this experiment are NH3 and OH–. Examples of reactions that produce complex ions are:

Cu2+(aq) + 4 NH3(aq) → Cu(NH3)42+(aq)
Zn(OH)2(s) + 2 OH–(aq) → Zn(OH)42+(aq)

Evidence of reaction in the first case is a color change. Copper(II) ion in solution is light blue and the Cu(NH3)42+ complex ion is dark blue. Notice that no solid is present so the solution will be clear, though highly colored.

For the second equation, evidence that reaction has occurred is the disappearance of the white solid, Zn(OH)2. Reaction of solid Zn(OH)2 with excess OH– produces a clear, colorless solution containing the complex ion Zn(OH)42–.

Question 10.1: Two other examples of complexing agents (besides NH3 and OH–) are Cl− and CN−. Write the net ionic equation for each of the following reactions that form complex ions.

a. Cobalt(II) ions in aqueous solution react with concentrated hydrochloric acid to produce the complex ion, CoCl42−(aq).

b. Zinc ions react with cyanide ions, CN−, in aqueous solution to form the complex ion, Zn(CN)42−(aq).

For another set of tests in the procedure you will initially add a limited amount of NaOH(aq), and then an excess amount. A similar strategy will be used with the addition of NH3(aq). What is the purpose of adding first a limited amount of a reagent, and then a greater amount? Also, what do NaOH and NH3(aq) have in common?

Sodium hydroxide is a strong base and ammonia is a weak base, and so both are sources of hydroxide ion:

NaOH(s) → Na+(aq) + OH–(aq)
NH3(aq) + H2O(l) → NH4+(aq) + OH–aq)

If a precipitate forms upon addition of either of these reagents to a solution of a metal ion, the metal hydroxide has formed:

Zn2+(aq) + 2 OH–(aq) → Zn(OH)2(s)

If the precipitate dissolves when excess base is added, a complex ion has formed. In excess NaOH(aq), hydroxide concentration is high:

Zn(OH)2(s) + 2 OH–(aq) → Zn(OH)42–(aq)

In excess NH3(aq), concentration of molecular NH3 is high:

Zn(OH)2(s) + 4 NH3(aq) → Zn(NH3)42+(aq) + 2 OH–(aq)

Notice in both of these examples the initial metal hydroxide is insoluble (a precipitate forms) but the complex is soluble.

The third possibility for reaction in aqueous solution is an oxidation–reduction reaction. One reaction of this type that you will observe is the reaction between iron(III) ions and iodide ions:

2 Fe3+(aq) + 2 I–(aq) → 2 Fe2+(aq) + I2(aq)

Iron(III) ion has been reduced and iodide ion has been oxidized. Evidence for this reaction is the yellow to brown color of the product I2 in aqueous solution. I2 also gives a distinctive purple color in dichloroethane.

Question 10.2: Write the net ionic equation for the following oxidation–reduction reaction. Substances labeled (aq) are soluble and ionic.

4 Zn(s) + 2 HNO3(aq) + 8 HCl(aq) → 4 ZnCl2(aq) + N2O(g) + 5 H2O(l)

The experiment consists of a series of reactions between metal ions in solution and a second reagent. This system of qualitative analysis will then be applied for identification of metal cations in an unknown sample.

The procedure has six parts:

  1. Reaction with sulfate ion, SO42–.
  2. Reaction with iodide ion, I–.
  3. Reaction with carbonate ion, CO32−.
  4. Reaction with limited, then with excess, hydroxide ion, OH–.
  5. Reaction with limited, then with excess, ammonia, NH3.
  6. Testing of unknown sample.

It will be necessary to observe the solutions closely and carefully record your observations. Note color changes including intensity of colors for both solutions and solids. Note formation of a precipitate and describe the texture of the solid; e.g., grainy, gelatinous, powdery, cloudy.

Background Information on Paint Pigments

Artists commonly mix white pigments with blue pigments to achieve lighter shades, so assume that your unknown solution contains cations from both a white and a blue pigment mixed together. Use the Pigment Handbook (given on the following page) to determine which pigments are possibilities for your sample. You will determine the procedure to identify the cations on your own. Use the observations from Parts A–E to guide your decisions. In this testing, a positive result or a negative result may be used to confirm the identity of a cation.

Once you have identified the cations, the corresponding pigments can be identified using the Pigment Handbook that follows. Many of the pigments used during the Renaissance period (from the 14th through the 17th centuries) are no longer in use, having been replaced by synthetic pigments. Therefore, the identity of the pigments should provide important evidence as to whether the samples from each painting can be traced to “A Maid Asleep,” painted in 1656, or a forgery. Based on the dates when the pigments were either introduced or replaced, it should be possible to estimate the range of dates when the painting was most likely produced (i.e., not before 1810, prior to 1725, etc.). If the paint sample contains a combination of pigments that were not both used during the Renaissance period, the painting should be considered a forgery.

Remember that your work will be used as evidence in court. In order to be valid, you must keep careful details of your experimental procedures and describe the reasoning used to conclusively identify the pigments present and the time interval when the painting is likely to have been produced.

Pigment Handbook (White and Blue Pigments)

The following are pigment possibilities for the six paintings. Each pigment contains only one cation. Remember, your sample will have the cation from one blue pigment mixed with the cation from one white pigment.

White Pigments

Pigment Name: Lead White Chemical Formula: 2PbCO3–Pb(OH)2

Used since antiquity, lead white was the only white used in European easel paintings until the 19th century. It is one of the oldest man-made pigments, and its use dates back to the Ancient Greeks and Egyptians, but continues in the modern day.

Pigment Name: Zinc White Chemical Formula: ZnO

First introduced in 1840, this white is colder in appearance than lead white and doesn’t cover nearly as well, yet it is far less expensive.

Blue Pigments

Pigment Name: Azurite Chemical Formula: 2CuCO3–Cu(OH)2

Azurite was the most important blue pigment in European painting during the middle ages and Renaissance. It was replaced when “Prussian blue” was discovered in the 18th century.

Pigment Name: Prussian Blue Chemical Formula: Fe[Fe3+Fe2+(CN)6]3

The first modern, artificially manufactured color was Prussian blue. It was made by Diesbach in Berlin around 1704. Diesbach accidentally formed the blue pigment when experimenting with the oxidation of iron. The pigment was available to artists by 1724 and has been extremely popular since its discovery through to the modern day.

Pigment Name: Cobalt Blue Chemical Formula: CoAl2O4

Gahn and Wenzel found cobalt aluminum oxide during research on cobalt compounds in 1777. However, the color was not manufactured commercially until late in 1803 or 1804.

Materials Required

Equipment

  • 11–4-mL test tubes
  • 2–25-mL test tubes
  • small test tube rack
  • medicine dropper

Chemicals

  • 0.1 M aluminum nitrate, Al(NO3)3
  • 0.1 M barium nitrate, Ba(NO3)2
  • 0.1 M cobalt(II) nitrate, Co(NO3)2
  • 0.1 M copper(II) nitrate, Cu(NO3)2
  • 0.1 M iron(III) nitrate, Fe(NO3)3 in 1 M HNO3
  • 0.1 M magnesium nitrate, Mg(NO3)2
  • 0.1 M manganese(II) nitrate, Mn(NO3)2
  • 0.1 M lead nitrate, Pb(NO3)2
  • 0.1 M zinc nitrate, Zn(NO3)2
  • 0.1 M sodium sulfate, Na2SO4
  • 0.1 M potassium iodide, KI
  • 0.1 M sodium carbonate, Na2CO3
  • 6 M sodium hydroxide, NaOH
  • 6 M ammonia, NH3
  • 0.1 M unknown cation solutions
  • 1,2-dichloroethane, C2H4Cl2

Common Equipment

  • centrifuge

Cautions

Salts of Ba2+ and Pb2+ are toxic; be sure to wash your hands thoroughly. 1,2-dichloroethane is a suspect carcinogen and may be irritating to the respiratory tract. Fume hoods must be on.

Procedure

Label a 25-mL test tube with your initials and give it to your lab instructor to be filled with your unknown solution (contains two cations).

A. Reactions of metal ions with sulfate ion.
  1. Clean nine 4-mL test tubes and rinse thoroughly with distilled water. Put 1 mL (approx. 20 drops, estimate ¼ full test tube) of a different metal ion solution in each of the nine test tubes. Make sure that there is no precipitate visible at this point.
  2. Add 1 mL of 0.1 M Na2SO4 to each of the nine test tubes (estimate ½ full test tube).
  3. Stopper and shake each test tube. In your lab notebook, record when a precipitate forms, noting the color and appearance of the precipitate. For example: a thick, white ppt. (precipitate). Be as specific as possible. Check carefully for filmy or gelatinous precipitates by holding the test tube up to the light—it may also be helpful to centrifuge the test tube. Note that some solutions change color without the formation of precipitate. Record that a colored solution formed.
B. Reactions of metal ions with iodide ion.
  1. Repeat step 1.
  2. Add 1 mL of 0.1 M KI to each of the nine test tubes.
  3. Repeat step 3. If a precipitate forms, it is necessary to centrifuge the test tube to clearly observe the color of the solution and the precipitate. If the solution is colored, add 10 drops of dichloroethane and record the color of the dichloroethane layer also.
C. Reactions of metal ions with carbonate ion.
  1. Repeat step 1.
  2. Add 1 mL of 0.1 M Na2CO3 to each of the nine test tubes.
  3. Repeat step 3.
D. Reactions of metal ions with hydroxide ion.
  1. Repeat step 1.
  2. Add one drop of 6 M NaOH to each of the nine test tubes. Add three extra drops to the acidic iron(III) nitrate solution (limited amount).
  3. Repeat step 3.
  4. Add 1 mL of 6 M NaOH to each of the nine test tubes (excess amount).
  5. Stopper and shake each test tube. Check for any change in appearance, especially if the initial precipitate dissolves.
E. Reactions of metal ions with ammonia.
  1. Repeat step 1.
  2. Add one drop of 6 M NH3 to each of the nine test tubes. Add three extra drops to the acidic iron(III) nitrate solution (limited amount).
  3. Repeat step 3.
  4. Add 1 mL of 6 M NH3 to each of the nine test tubes (excess amount).
  5. Stopper and shake each test tube. Check for any change in appearance, especially if the initial precipitate dissolves.
F. Analysis of unknown solution.
  1. Use observations in Parts A–E to develop a procedure that will allow you to determine the cations found in the unknown. Record the procedure in your lab notebook. Be specific as to amounts and identity of each solution used—this procedure will be included in your formal report. Complete the procedure and record your observations.

Waste Disposal

All solutions containing Pb2+, Co2+, Ba2+, dichloroethane, and unknowns should be collected in a beaker at your desk and transferred to the inorganic waste beakers in the hood upon completion of the experiment. You must fill in the waste disposal sheet. Your lab instructor will dispose of the total volume in the appropriate container.

Laboratory Report

The procedure section of your report should accurately record the procedure you devised to identify the cations in your unknown. This section should be concise without unnecessary detail or “fluff.” The reader should be able to duplicate this part of the experiment using only this section. Include a reference to the procedure for Parts A–E, as you did not create this part of the experimental procedure.

Report your observations from Parts A–E on the Report Form. When a precipitate formed, use ppt. as an abbreviation and report the color and appearance of the precipitate. Also report the color of solutions that formed. Use NR for No Reaction.

Create your own report format for Part F. It should contain all observations from your unknown tests. You may format your observations any way you like, as long as the information is displayed in a complete, accurate, and organized manner. Clearly identify both cations present in your unknown. (Include any additional information required in the handout.) Describe and write equations for the reactions you used to identify your cations. Include the identity of any precipitates formed.

In the conclusion section, discuss the implications of your results. Include how your unknown results exclude certain pigments and to what extent they confirm the presence or absence of the cations. Discuss the relevance of the time frame of the painting. Based on your results, discuss the likelihood that your sample was from the original painting or from a forgery. Remember, these results will be testified to in court, so be thorough and accurate.