In the Classroom edited by
Overhead Projector Demonstrations
Doris K. Kolb Bradley University Peoria, IL 61625
Chemical Equilibria Involving Copper(II) Ethylenediamine Complexes Roberto Zingales Department of Inorganic and Analytical Chemistry, University of Palermo, Viale delle Scienze—Parco d’Orleans II, I 90128 Palermo, Italy;
[email protected] Copper(II) ions and ethylenediamine molecules form two fairly stable complexes in aqueous solution (1). If the ligand is added dropwise to a solution containing a copper(II) ion, the stepwise formation of two species can be recognized by their different colors. If these reactions are carried out on the overhead projector, they can be shown to a large class of students. Procedure Place a transparency sheet on the overhead projector and arrange three Petri dishes in a row at the top. Pour a thin layer of 0.2 M copper(II) sulfate solution into each Petri dish. Leave the solution in the first dish unchanged. Add five drops of aqueous 25% ethylenediamine solution to the second dish. Swirl gently to mix the solutions until the color turns from pale blue to dark blue, thus confirming the formation of the [Cu(en)(H2O)2]2⫹ complex, according to the reaction [Cu(H2O)4]2⫹ ⫹ en
[Cu(en)(H2O)2]2⫹⫹ 2H2O
(1)
Add ten drops of the ethylenediamine solution to the solution in the third dish, rotating to mix the solutions, until the dark violet [Cu(en)2]2⫹ complex is formed, according to the reaction [Cu(H2O)4]2⫹ ⫹ 2en
[Cu(en)2]2⫹ ⫹ 4H2O
(2)
Write on the transparency sheet, beneath each dish, the chemical formula of the corresponding copper(II) species, highlighting the different colors of the copper(II) aquo complex and of its ethylenediamine complexes. Put a fourth Petri dish on the projector underneath the first row of Petri dishes and pour a thin layer of the copper(II) sulfate solution into it. Add ten drops of the ethylenediamine solution, gently swirling to mix the solutions, so that the mixture becomes dark violet. Then cautiously add 6 M sulfuric acid solution one drop at a time, swirling the dish after each addition, and taking note of the number of drops needed to change the dark violet color to dark blue. Then add the same number of drops of sulfuric acid solution till the pale blue color of the copper(II) aquo complex is restored. Finally, cautiously add 6 M sodium hydroxide solution dropwise. Each drop colors the solution violet and then the color fades as the solutions are thoroughly mixed. Count the number of drops needed to give the solution the same dark blue color as the [Cu(en)(H2O)2]2⫹ complex in the top row. Then add an equal number of sodium hydroxide drops, until the dark violet color of the [Cu(en)2]2⫹ complex has been restored (check with its solution in the top row).
Discussion A goal of this demonstration is to show complex formation reactions with ethylenediamine as a ligand and copper(II) as the ion to be coordinated. Crystal field theory (2) can be used to justify the different colors of the aquo complex and of the two ethylenediamine complexes or, conversely, the colors can be used to give a visual demonstration of the theory. Moreover, the colors of the complexes can be used to show the stepwise formation of two different products, according to eqs 1 and 2. Finally, the number of drops of ligand solution required to obtain each of the complexes confirms, in a semi-quantitative way, that the second complex contains twice as many ligand molecules as the first one. In the second part of the demonstration, the two complexes are destroyed and then restored by addition of the proper reactants. This may be used to illustrate simultaneous equilibria, masking of reactants, or Lewis acid–base reactions. Although copper(II) and ethylenediamine form stable complexes, the addition of sulfuric acid causes a color change from dark violet to dark blue to pale blue. As the colors of the solutions in the dishes on the top row confirms, the [Cu(en)2]2⫹ complex has been converted to [Cu(en)(H2O)2]2⫹ and then to the copper(II) aquo complex, according to reactions [Cu(en)2]2⫹ ⫹ H⫹ ⫹ 2H2O
[Cu(en)(H2O)2]2⫹ ⫹ enH⫹ (3)
[Cu(en)(H2O)2]2⫹ ⫹ H⫹ ⫹ 2H2O
[Cu(H2O)4]2⫹ ⫹ enH⫹ (4)
Reactions 3 and 4 are the result of two complex formation equilibria (eqs 1 and 2) and of the protonation of ethylenediamine molecules, which show a basic behavior, according to the equation en ⫹ H⫹
enH⫹
(5)
Logarithms of equilibrium constants for eqs 1–5 are reported in Table 1. According to the laws governing chemical equilibria, the addition of an increasing amount of hydrogen ions shifts reactions 3, 4, and 5 to the right. In other words, ethylenediamine molecules are increasingly involved in the formation of the protonated enH⫹ species, which is more stable than the copper complexes. Thus copper(II) ions revert back stepwise to their aquo complex, as confirmed by the characteristic pale blue color of the solution. Addition of sodium hydroxide gradually neutralizes the hydrogen ion excess, reversing the previously described sequence of reactions. Ethylenediamine molecules are set free and their copper(II) complexes restored, as shown from their colors.
JChemEd.chem.wisc.edu • Vol. 80 No. 5 May 2003 • Journal of Chemical Education
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In the Classroom
with the protons, leaving the ligand free to coordinate copper(II) ions.
Table 1. Equilibrium Constants of Reactions Involved in the Demonstrations Reaction
Log K1
1. [Cu(H2O)4]
2⫹
2⫹
⫹ en
Cu(en)(H2O)2] ⫹ 2H2O
2. [Cu(H2O)4]2⫹ ⫹ 2en 3. [Cu(en)2]
2⫹
[Cu(en)2]2⫹ ⫹ 4H2O
⫹
⫹ H ⫹ 2H2O 2⫹
⫹
enH⫹
19.99 2⫹
[Cu(en)(H2O)2]
4. [Cu(en)(H2O)2] ⫹ H ⫹ 2 H2O 5. en ⫹ H⫹
10.66
⫹
⫹ enH
[Cu(H2O)4]
2⫹
0.60 ⫹
⫹ enH
-0.73 9.93
Reactions involved in this demonstration can also be explained according to the acid–base theory (3) proposed by Lewis to unify all kinds of chemical reactions into a general pattern. Copper(II) ions show an acidic behavior, while water and ethylenediamine molecules behave as bases. By sharing their lone pairs of electrons with copper(II) ions, water and ethylenediamine form coordinate covalent bonds or dative bonds. Ethylenediamine is a stronger base than water, so that it removes copper(II) ions from the aqueous complex, according to reactions 1 and 2. Hydrogen ions show the same acidic properties as copper(II) ions, so that they compete for the lone pairs of electrons in ethylenediamine. Of course, the stronger acid prevails and the copper complex is destroyed in favor of the protonated ethylenediamine species. Conversely, hydroxide ion is a stronger base than ethylenediamine, and thus reacts
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Hazards Caution is needed in preparing the solutions, especially the sulfuric acid. Concentrated sulfuric acid is both a strong oxidant and powerful dehydrating agent. It can severely burn and blister the skin, and it puts holes in clothes. In preparing dilute solutions always pour the acid into the water. Sodium hydroxide pellets are quite caustic. Avoid contact with skin, especially near the eyes. To make up a solution, add the weighed pellets to cool water with stirring. When NaOH dissolves in water heat is evolved. Wear safety glasses when preparing or presenting the demonstration. Acknowledgment The author is greatly indebted to Roberto Triolo for his revision of the English text. Literature Cited 1. Harris, D. C. Quantitative Chemical Analysis, 4th ed.; W. H. Freeman: New York, 1995; pp AP24, AP43. 2. Brown, T. L.; Le May, H. E., Jr.; Bursten, B. E. Chemistry. The Central Science, 6th ed.; Prentice Hall International: London, 1994; pp 943–950. 3. Lewis, G. N. Valence and the Structure of Atoms and Molecules; The Chemical Catalog Co.: New York, 1923.
Journal of Chemical Education • Vol. 80 No. 5 May 2003 • JChemEd.chem.wisc.edu
