Charles E. Ophadt Elmhurst College Elmhurst, Minois 60126
A Novel Approach for Qualitative Analysis
In order to present chemistry as a stimulating and fascinating pursuit, a second semester freshman laboratory program involving nontraditional qualitative investigations of metal ion properties was designed. The trend to omit traditional qualitative analysis and replace it with more quantitative experiments a t the freshman level results in a serious lack of understanding of ionic and molecular properties. I agree with Mel Gorman that we should "forget quant for freshman" (1). The lab program first involves providing some basic knowledge of reaction principles by conducting preliminary experiments. The students are able to discover for themselves the metal ion properties of interest. Secondly, the students are expected to use the information gathered to design further original experiments. A chemistry laboratory featuring the "discovery" of something unknown can be both exciting and intellectually rewarding. Several students were amazed to find that they experimentally discovered information not readily found in textbooks. For example, one student discovered that Cu(0H)z and Co(0H)z are both slightly amphoteric. The laboratory investigations are divided into three Parts Synthesisof a Metal Salt A salt is synthesized, identified, and partially analyzed. Properties of Metal Ions and Their Separations Metal ion equilibria associated with precipitation, weak acids and bases, amphoterism, and complex ions are utilized by the student to develop his own qualitative analysis separation scheme toidentify an unknown. Oxidation States of Metals Various oxidation states of metals are prepared and identified with the appropriate oxidizing and reducing agents selected by the student. (These experimental methods have been developed by the author.) Each student is given a unique set of five metals that he investigates throughly during the semester. The various properties of metal ions can he adequately demonstrated with only five metal ions without heing too repetitive and time consuming. The five metal sets are choosen from the following metals grouped by the number of their oxidation states: one oxidation state; AI(III), Ca(II), Cd(II), Mg(lI), Ni(lI), Pb(II), Zn(II), two oxidation states; Co(II,III), Cu(I,II), Fe(II,llI), Hg(I,II), Sn(II,IV), and several oxidation states; C r , I , V I Mn(II,IJI,N,VI,W). The following criteria are used in making up each five metal set: two metals of one oxidation state, two metals of two oxidation states, and either chromium or manganese are included to make five metals; a t least one metal must he amphoteric; a t least one metal must form an ammine complex with ammonium hydroxide; one metal must form a sulfide precipitate only in basic sulfide; and finally either lead or mercury must be included as a well known environmentally relevant metal. The lectures in the course which stress the application of equilibrium principles are closely correlated with the laboratory investigations. In addition to the laboratory investigations, library resources are utilized by the students to write short reports about the metallurgy and environmental aspects of the five metals and the effect of one
metal on hiotogical systems. The metallurgy report is supposed to include brief information about composition of ores and location, pretreatment of ores such as roasting and hydrometallurgy, how the metal ion is reduced to the metal, refining, uses, careless entry of the metal into the environment, and toxicity. The report on the effect of a metal on biological systems should contain information about the effects of chelation and both beneficial and harmful effects due to an excess or lack of a particular metal. I) Synthesis of a Metal Salt Synthesis is included in many laboratory hooks to teach various techniques. What kind of learning experience takes place if the student knows in advance the reaction equations, name, and exact formula of the compound heing synthesized? A more appropriate approach, which illustrates how chemists analyze new and unknown compounds and as utilized by Sanavely and Yoder (2), is to eive directions for the svnthesis and then let the student k a l y z e what and how much is present in the salt. Experimental directions are eiven for the ~ . r e.~ a r a t i oofn double salts of potassium or ammonium and divalent metal sulfates 12). Other experimental directions for the synthesis of certain coordination compounds are also given. The student must look up, use, or modify the qualitative tests in the text (3) to determine which ions and molecules present in the reaction mixture are present in the product. A conductometric analysis is carried out to determine the number of ions present in a solution of the salt. Quantitative analysis on water of hydration and colorimetric analysis of the metal ions are also suggested (2). Further quantitative analyses may be performed but there usually is not enough time. When the student reports what is present, he is given the elemental percentage composition and the molecular weight. This information is used by the students to check their quantitative work, find the molecular formula, write a balanced equation for the synthesis, and calculate the theoretical yield and percent yield. Finally, the students are asked to suggest a geometrical arrangement of the cations, anions, and molecules, knowing that a metal complex is present. The students are usually amazed by the crystals, color, and formula of the compound that they have synthesized. I I ) Properties of Metal Ions and Their Separations The method of qualitative analysis by observing the hehavior of metal ion with certain reagents and then having the student develop a separation scheme has been used effectively by others (4-6). A student develops a much greater understanding of reaction and equilihrium principles by developing his own qualitative analysis scheme than by following traditional "cookbook" procedures. The reagents used in this experiment are: sodium hydroxide, ammonium hydroxide, hydrogen sulfide and sulfide ions generated from tbioacetamide, and ammonium oxalate. Several preliminary experiments are carried out to develop the thinking and reasoning abilities of the student to eiahle him to properly interpret reactions that he observes. For example, he needs to know what specifically is present in a precipitate or causes the color of a solution to Volume 51, Number 6, June 7974
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change (formation of a complex ion) in order to write an equation to represent the reaction. The synthesis of a salt in Part I) illustrates one method of doing this. These preliminary experiments illustrate a second method for finding the products of a reaction. The first experiment is merely an exercise in reasoning to determine which ions or molecules cause the colors of several salts and solutions. In the next experiment, the student is asked to carry out reactions by substituting or eliminating ions and molecules to determine what causes the blue and red layers of a solution made by carefully pouring acetone on top of a cobalt(I1) chloride solution. Selected reagents are made available for testing. The tetracbloro- and hexaaquocobalt(10 complex ions, respectively, cause the two colors. The last preliminary experiment is carried out to find the differences and similarities in the action of limited and excessive amounts of sodium and ammonium hydroxide on copper(I1) ions. The reason for these experiments is that simply telling about hydroxide precipitates, ammine and hydroxo complexes has little meaning without some experimental verification by the student. The experimental directions are summarized as follows: Add 1or 2 drops of either 9 M NaOH or 6 M NHIOH to 10 drops of 0.1 M copper(II) nitrate solution and observe the precipitate. T o the precipitate, add 15-20 drops of either 9 M NaOH or 6 M NHaOH and observe that the precipitate dissolves in both cases and the resulting solutions are light blue and deep blue, respectively. The student is then asked to devise some experimental procedure to identify specifically the ions or molecule which first gives the precipitate and then causes it to dissolve. Some hints are given about weak and strong bases and the ions and molecules which are present in both bases. Further hints are given that they should try different sources of some ions or leave them out entirely to isolate certain ions which may be causing the reaction. Almost everyone is able to identify that both precipitates are copper(I1) hydroxides and that the hydroxide ion in NaOH causes the precipitate to dissolve forming the tetrahydroxocopper(I1) complex ion. About three-fourths of the students, if reminded of the difference in colors of the solutions with excess bases, are able to say by the process of elimination that NH3 causes the precipitate to dissolve when using excess NH40H forming the tetraaminecopper(I1) complex ion. These reactions provide "model reactions" for the other ions and begin to show how solubility rules and K , tables may effectively be used. Solubility rules and K,, tables are used directly to determine the products of the reactions of hydrogen sulfide or ammonium oxalate with metal ions. The student studies the behavior of his five metal ions (using both oxidation states of two metals and Cr(II1) or Mn(I1)) with the reagents previously listed. Knowing the behavior of two oxidation states of two metals is important for understanding Part III) of this experiment. A precautionary hint is given that not all precipitates formed under basic conditions with sulfide ions are metal sulfides. All precipitates are dissolved if possible using 6 M HCI or HN03. The results of the reactions of the five metals with the five reagents are summarized and then used by the student to develop a separation and identification scheme in the form of a flow chart. The only limitation given is that either NaOH or NHrOH must be used for the first separation. This prevents the direct use of traditional, classical sulfide schemes found in numerous textbooks. Encouragement is given to students to look up confirmatory tests and general information about the metals in "qual" texts made available and in the library. The proposed separation scheme is submitted to the instructor to be sure that gross deficiencies are not present. The student is made aware of possible difficulties that may be caused by incorrect p H and incomplete separations. The student is then supposed to make up a known mixture of the metal ions, test, and modify, if necessary, his separa416
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tion scheme. An unknown of three ions is then given. The final report submitted includes the flow chart, net ionic equations, and the reasoning for the choice of ions reported to he present. Ill) Oxidation States of Metals The purpose of studying the oxidation. states of metal ions is to become familar with the principles of oxidationreduction reactions. Many experiments on oxidation-reduction merely say to add certain reagents together and a certain oxidation state will result (7). Little mention is made of why the reaction works. In this experiment the student must determine for himself which oxidizing or reducing agent to use to prepare a certain oxidation state of a metal. These methods will he described in more detail than the previous two parts since these have not appeared elsewhere. The main tools for these experiments are tables of half reaction oxidation potentials in the text that I use (3) and other text books (8,9).The main principle followed in the use of these tables is: A reducing agent will react with an oxidizing agent below it in the series. This rule of course does not consider concentration effects or the rate of reaction. Applying this principle in test-tube reactions can give more of an understanding of oxidation-reduction processes than mechanical mathematical calculations. In the process of experimentation and interpretation of reactions. the terminolow of redox reactions becomes clear and meaningful to thgstudent. The general directions for the experiment are as follows: For thi metals that have two m a h oxidation states, dissolve the metal in either 6 M HCI, 6 M HN03, or 6 M NaOH, identify the resulting oxidation state, and use this solution to react with some other oxidizing or reducing agent to prepare and identify the second oxidation state. For chromium and manganese, start with the metal or any available salt. select an ao~ropriateoxidizine or reducing agent, andprepare and-identify all possible oxidation states including those used as starting materials. In order to surmount the initial c ~ n c ~ ~ t difficulties ual of the selection of an appropriate oxidizing or reducing agent by the application of the principle that a reducing agent will react with an oxidizing agent below it in the series, a preliminary experiment was designed and general directions and principles for redox reactions were written out. Preliminary Redox Experiment The preliminary experiment involves ,the reaction of hydrochloric and nitric acids and sodium hydroxide with the metals. The purpose is to show how the relative strengths (position in the oxidation potential series) of oxidizing and reducing agents will effect the product, if any, of the reaction. The directions are as follows: To 1 ml of water, add 50-100 mg of one metal (powdered or granular), then add 5-10 ml, with the first few drops added cautiously, of one of the following: 6 M HC1, 6 M HNOJ, or 6 M NaOH. Repeat until all five metals have been used and all three reagents are tested with each metal. Prolonged heating in a hot water bath may be necessary to get the reaction, if any, to take place. The oxidation state of the metal ions and the gas evolved (perhaps HZ, 0 2 , Clz. NO, NOz) are identified for the reactions which do occur. The metal ions may be identified through the properties investigated in Part II). The student is encouraged to look up properties and methods of identifying the gases, although simple tests are provided. In the interests of time, a full identification of the gases is made using only one metal with the results serving as "models" for the remaining metals. The reaction of metals with HNOB gives a good introduction to the properties of "NO," which are always considered in air pollution studies.
The results of the preliminary experiment are analyzed with the aid of a brief oxidation potential table (provided in the student write-up) containing the half reactions of the acids and hase which may be involved in the reaction. The half reactions of the metals involved are to be inserted into the table by the student. The student should be able to see how the main principle of the redox experiment is applied and answer the following questions: Why did some metals give no reaction a t all with HC1 but did react with HNOJ? Why are the oxidation states of metal ions sometimes different if HN03 was used rather than HCI? Why are O2 and C12 gases not formed? Why did some metals dissolve in hase and not others? General Directions and Principles for Redox Reactions Selection of Oxidizing and Reducing Agents. A brief oxidation potential series for the metals with all oxidation states involved should he written out by the student. The appropriate oxidizing or reducing agent is to be selected by the application of the main principle: A reducing agent will react witb an oxidizing agent below it in the series. The half reactions for the oxidizing or reducing agent being considered should be inserted into the series of metal half reactions. An oxidation potential series containing possible oxidizing and reducing agents is provided as a "starter" since the students have no idea which compounds are commonly available if the long tables in the textbooks (3, 8, 9) are used a t the beginning. Some of the reagents listed are: HzOz, NaC10, SO?, 108-, 0 2 , halogens, and halides. The equations for the hypothesized reactions are to be written out before any experimental work is attempted. Reactron Conditions. It is necessary that the reaction be carried out in acidic, neutral, or basic conditions as indicated by the H+, OH-, or absence of both in the half reactions because the potentials change with pH. The concentration of the solutions should be dilute or the reaction results will he ohsrured. Usually 50-100 mg of solids and 2 5 mlot liouidi added drouwise at tirst are adeauate. ~dentification of thh Oxidation State. ' Experimental proof should be given that the desired oxidation state of the metal was in fact obtained. These tests should be in addition to the observation of a desired color or precipitate in the primary reaction. This proof may come from further confirmatory reactions as studied in Part (II) or looked up in textbooks, or the product of the primary reaction may be further reacted with some other oxidizing or reducing agent. There are two possibilities which may hinder the identification procedures. First, a large or even a slight excess of umeacted oxidizing or reducing agent may react with the next reagent preventing the desired reaction. For example, HNOa would react with reducing agents above it in the series if they were added next. Whether this may occur will be found by lwking a t the tables of half reac-
tions and applying the main principle of this experiment. Secondly, in addition to the desired metal ion in the test tube, there will of might be the product of the oxidizing or reducing agent used in the reaction. This other product may also react witb the next reagent added. Again consult tables of half reactions. This complication is a second reason why some common oxidizing and reducing agents are suggested to the students. Most of these have rather "inert" or easily removed by-products. Instability of Metal Oxidation States. The desired metal oxidation state mav he unstable with resuect to reactions with water or oxygen in air or dissolved in water and thus prevent its preparation or observation. Although a given oxidation state may be unstable with respect t o water or oxygen, many of these may he prepared and studied anyway because the unstable oxidation state reacts very slowly. Water or H+ ion may react as an oxidizing agent at all pH's with a reducing agent and product Hz gas as one product. For example, sodium metal is unstable in water. Water or OH- ion may also react as a reducing agent a t all pH's with an oxidizing agent to produce Oz gas as one product. For example, cobalt(1II) ion is unstable in water. Oxveen in air or dissolved in water is a good oxidizing agent.-i familar example is the rusting of iron-the (0) and (II) states are unstable in this environment. The appropriate half reactions are indicated in the student write-uu. Finally, the possibilitv of making a complex ion or precipitate to stabilize a gi;en oxidation state is suggested. The student report is to include statements about the stability of all oxidation states toward water and oxygen. Conclusion Probably the hardest part of a laboratory program such as this is to get the student thinking and experimenting on his own. I feel that the preliminary experiments are able to provide a basic knowledge on which the student may build with confidence. At the completion of the course the student is thoroughly familar with the properties of five metals rather than having memorized detailed flow charts on about 20-25 metals or having no qualitative analysis whatsoever. In addition the student understands the principles of oxidation-reduction. Literature Cited
($ Thomplon, M . L . , mdBix1er.J. W.,J.CHEM.EDUC., 48 119(19711. (51 Sterretf. F. S. K.. Kennedy. S. E.. and Sparberg. E. B.. "A Laboratory Ihwntigatioll of Concepts in Chemistry." Hampr and Row, New York. 1468, p 149. 161 OConnor. R.. and Woelfd W. C.. "Freeman Library of Lsboratory Separates in Francisco. 1971. Vol.11. No. 1121. ChemiPtry." W . N . F ~ e e m a n C m p a n y Ssn . 171 Rantz. H . W.. and Mslm. L. E.. "Freeman Library of Lsboiatory Separates in Chemistry." W. H. FreemsnCompany, SanFrsncisco, 1970. V o l I, No. 1030. 181 Day, M . C.. and Selhin. .I.. "Theoretical Inorganic Chemistry." 2nd Ed.. Van Nostrand Reinhold Co.. New York, 1169. p 314. 191 Huhaoy, l. E.."lnorganieChemisiry."HerporsndRow. NewYork. 1972,p257.
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