In the Classroom edited by
Overhead Projector Demonstrations
Doris K. Kolb Bradley University Peoria, IL 61625
W
A Colorful Look at the Chelate Effect Donald C. Bowman Department of Science, Math, and Engineering, Central Virginia Community College, Lynchburg, VA 24502;
[email protected] The chelate effect is a generalization stating that a complex containing one or more five- or six-membered chelate rings is more stable than a structurally similar complex that lacks these chelate rings (1). This greater stability is experimentally evidenced by the preferential formation of a chelate in solution in which a metal ion exists with a polydentate ligand and a monodentate ligand of equal concentrations (2). The chelate effect can be demonstrated visually by noting color changes (explained by crystal field theory) that occur during ligand-exchange reactions involving monodentate and polydentate ligands and transition-metal ions such as copper(II) and nickel(II). Solutions and Equipment • 0.025 M CuSO4 is an aqueous solution containing 0.62 g of copper sulfate pentahydrate per 100 mL of solution. • 0.10 M NiSO4 is an aqueous solution containing 2.63 g of nickel sulfate hexahydrate per 100 mL of solution. • 1 M NH3 and 6 M NH3 are aqueous solutions prepared by diluting a stock solution of concentrated ammonia (14.8 M). • 5% Ethylenediamine is an aqueous solution containing 5 g of ethylenediamine and 95 g of water. This is approximately a 1 M solution. • 30% Ethylenediamine is an aqueous solution containing 30 g of ethylenediamine and 70 g of water. This is approximately a 6 M solution. • 15% Triethylenetetramine is an aqueous solution containing 15 g of triethylenetetramine and 85 g of water. This is approximately a 1 M solution. • 1 M Pyridine is an aqueous solution containing 8 mL of pyridine and 92 mL of water. • 1 M 2,2´-Bipyridine is an ethanolic solution containing 15.6 g of 2,2´-bipyridine per 100 mL of solution. • (3) Six-well polystyrene reaction plates or three sets of six 50-mL beakers • Medicine droppers or dropper bottles • Stirring rod
Procedure Place a six-well polystyrene reaction plate (or a set of six 50 mL-beakers) on the overhead projector and prepare three control solutions by adding 4 mL of 0.025 M copper(II) sul-
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fate solution to three wells and stirring in (one solution to a well) 20 drops of 1 M NH3, 20 drops of 5% ethylenediamine and 20 drops of 15% triethylenetetramine. Now add 4 mL of 0.025 M copper(II) sulfate solution to the other three wells. The solutions appear pale blue. Add 20 drops of 1 M NH3 to each of these three wells and gently stir. The solutions now appear deeper blue. Add 20 drops of 5% ethylenediamine (en) solution to two of the three wells containing copper– ammonia complex and stir until the solutions are a dark purple. Add 20 drops of 15% triethylenetetramine (trien) solution to one of the two wells now containing copper–(en) complex and stir until the color becomes a deep royal blue. Demonstrate that a more stable complex will not easily revert to a less stable complex by adding an excess of the ligand that formed a less stable complex. Add 20 drops of 5% ethylenediamine solution to the well containing the copper– (trien) complex. Gently stir the solution and note that no color change is observed. Add 20 drops of 1 M NH3 to the other copper–(en) complex and note that again no color change is observed. To show that the effect is not exclusively a phenomenon of ammonia and simple alkyl amine ligands, perform a similar experiment with copper(II) ions and aromatic amine ligands, again preparing control solutions (4 mL of 0.025 M copper(II) sulfate with 20 drops of 1 M pyridine and 4 mL of 0.025 M copper(II) sulfate with 20 drops of 1 M 2,2´bipyridine) in two wells of a second six-well reaction plate on the overhead projector. Add 4 mL of 0.025 M copper(II) sulfate solution to two of the other wells. Add 20 drops of 1 M pyridine to each of the two wells. Gently stir and the solutions become a deep sky blue. Add 20 drops of 1 M 2,2´bipyridine (bipy) to one of the wells and stir until the solution turns to an aquamarine color. Adding 20 additional drops of 1 M pyridine to the copper–(bipy) complex will not change the color back to sky blue. Finally, to show that the effect is not specific to copper(II) ions, perform a demonstration using nickel(II) ions instead. Prepare control solutions (4 mL of 0.10 M nickel(II) sulfate with 20 drops of 6 M NH3 and 4 mL of 0.10 M nickel(II) sulfate with 20 drops of 30% ethylenediamine) in two wells of a third six-well reaction plate. Then, add 4 mL of 0.10 M nickel(II) sulfate solution to two of the other wells. The solutions have a very light green color. Add 20 drops of 6 M NH3 to each of the two wells. The solutions turn sky blue with gentle stirring. Add 20 drops of 30% ethylenediamine to one of the wells, and upon stirring, the solution becomes magenta. When 20 additional drops of 6 M NH3 is added to the nickel–(en) complex, the solution will not return to the sky blue nickel–ammonia complex.
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In the Classroom
Hazards
Discussion
Caution is needed in preparing all solutions and they should be prepared in a fume hood. Copper(II) sulfate pentahydrate may cause eye and skin irritation. Nickel(II) sulfate hexahydrate may cause an allergic skin reaction and may cause eye and respiratory tract irritation. According to the National Toxicology Program 11th Annual Report on Carcinogens, all compounds of nickel are carcinogenic. Ammonia is corrosive and irritating to skin and eyes. Breathing ammonia irritates the lungs causing coughing and shortness of breath. Ethylenediamine is corrosive and irritating to skin and eyes. Breathing ethylenediamine irritates the lungs causing coughing and shortness of breath. It is harmful if swallowed, inhaled, or absorbed through the skin possibly causing liver, kidney, and respiratory system damage. It also may cause an allergic skin reaction. Triethylenetetramine is corrosive, causing burns to any area of body contact. It is harmful if swallowed, inhaled, or absorbed through the skin. It may cause liver, kidney, and respiratory system damage. Pyridine is irritating to the eyes, nose, and throat. Pyridine causes local irritation on contact with the skin, mucous membranes, and cornea. 2,2´-Bipyridine is very hazardous in case of ingestion or inhalation. It is an irritant to eyes and a slight hazard to skin upon contact. Safety goggles and gloves should be worn when preparing solutions. Safety goggles should be worn when presenting the demonstration.
The original solutions of copper(II) sulfate contain the complex ion [Cu(H2O)4]2+. By adding the 1 M NH3 solution, a ligand-exchange reaction occurs forming the new complex ion [Cu(NH 3 ) 4 ] 2+ . Upon adding the bidentate ethylenediamine ligand, another ligand exchange occurs according to the reaction
NH
NH Cu 2ⴙ
Figure 1. The structural formula for the copper–(trien) complex.
NH2 H 2N H 2N
ⴙ
Ni2
H2 N
Table 1. Stability Constants for Complex Ions Formed in the Demonstration Log β
[Cu(NH3)4]
12.6
[Cu(en)2]2+
19.6
[Cu(trien)]
20.6
2+
2+
[Cu(pyridine)4]2+
The nickel complexes have octahedral geometry. The structural formula for [Ni(en)3]2+ illustrates this geometry in Figure 2. The stability constant, β, of a complex is defined as the overall formation constant that results from the addition of n ligands to a metal ion according to the relationship MLn ; β =
[ ML n ] [ M][L ]n
(3)
The stability constants for the complexes formed in the demonstration are given in Table 1 (3–5). A complex with a larger formation constant is favored thermodynamically by the relationship (4) ∆G o = −2.3 RT log β The standard free energy has an enthalpy component and an entropy component according to the relationship (5) ∆G o = ∆H o − T ∆ S o
N NH2 H2
Figure 2. The structural formula for the nickel–(en) complex.
Complex
[Cu(en)2]2+ + 4NH3 (1)
When the copper–(en) complex reacts with triethylenetetramine, a tetradentate ligand, yet another new complex is formed, [Cu(trien)]2+. All of the copper complexes formed in the demonstration have square planar geometry. The structural formula for one example of the copper complex series, the copper–(trien) complex, is shown in Figure 1. In parallel fashion to the ammonia–(en) ligand exchange, a copper–pyridine complex forms a new complex when the bidentate ligand 2,2´-bipyridine is added to the solution. Aqueous solutions of nickel(II) sulfate contain the complex ion [Ni(H2O)6]2+. By adding 6 M NH3, the complex ion [Ni(NH3)6]2+ is formed. When ethylenediamine is added to this complex, a new complex forms according to the reaction [Ni(NH3)6]2+ + 3 en [Ni(en)3]2+ + 6NH3 (2)
M + nL
N H2
N H2
[Cu(NH3)4]2+ + 2 en
10.2
2+
[Cu(bipy)2]
17.9
[Ni(NH3)6]2+
08.6
[Ni(en)3]2+
18.6
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The typical explanation for why chelate complexes have larger negative free energy values than similar monodentate ligand complexes focuses on only entropy differences. A number of general and inorganic texts use this approach (6–12). This emphasis on entropy may result from the fact that the enthalpy difference contribution for this type of ligand-exchange reaction is sometimes very small [as evidenced in complexes involving cadmium(II) and zinc(II) ions] (13). What produces the often large entropy contribution? Note that in eq 1 there is an increase in the number of particles as the ligand exchange proceeds (3 reactant particles produce 5 product particles). A similar increase occurs in eq 2. This increase in the number of particles represents an increase in disorder of the system. Another explanation for chelate formation preference comes from recognizing how the reactions
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might proceed. To form a copper–ammonia complex requires four separate favorable collisions between the metal ion and the ligand molecules. Forming the copper–(en) complex requires an initial collision for the first ligand to attach by one arm. However, the other arm is nearby and only requires a rotation of the other end to enable the ligand to form the chelate ring; two such sequences produce the complex. This action would favor the formation of the complex with bidentate groups rather than monodentate groups (14). Although reasonable, these explanations are insufficient to explain the chelate effect. There are reactions where the enthalpy difference is a factor and may even rival or exceed the size of the entropy factor. For the ligand-exchange reaction between [Cu(NH3)4]2+ and two molecules of ethylenediamine, ∆H ⬚ = ᎑22.6 kJ兾mol, roughly 55% of the total contribution to the ∆G ⬚ at 25 ⬚C value of ᎑39.9 kJ兾mol (as calculated from β value difference in Table 1) (15). Possibly copper(II) ions form stronger bonds as a result of chelate ring formation and involvement of unfilled d orbitals allowing for crystal field effects not available for cadmium(II) and zinc(II) ions (filled d orbitals) or manganese(II) ions (half-filled d orbitals) (16). Regardless of the relative significance of the enthalpy term, it is still an oversimplification to explain the magnitude of the entropy value using only the concept of increase in disorder owing to increased particle count. Myers, and later Chung, showed that the entropy effect of chelation is much more involved, having at least five separate factors (17–19). For a general chemistry audience, instead of trying to examine the multitude of possible equilibrium and thermodynamic factors contributing to the chelate effect for a particular system, it may be best for the instructor to emphasize the experimental reality of the effect, and only briefly mention possible enthalpy and entropy factors. WSupplemental
Material
Digital pictures of the solutions of the complex ions formed in the demonstration are available in this issue of JCE Online. In addition, structures of the complexes discussed in this article are available in fully manipulable Jmol and Chime format as JCE Featured Molecules in JCE Online (see p 1248).
Literature Cited 1. Cotton, F. Albert; Wilkinson, G.; Gaus, P. Basic Inorganic Chemistry, 3rd ed; Wiley and Sons: New York, 1995; p 186. 2. Frausto da Silva, J. J. R. J. Chem. Educ. 1983, 60, 390. 3. Pflaum, R. T.; Brandt, W. W. J. Am. Chem. Soc. 1955, 76, 6216. 4. Martell, A. E.; Chaberek, S. Anal. Chem. 1954, 26, 1692. 5. Stability, Chelation, and the Chelate Effect. http:// agrss.sherman.hawaii.edu/courses/Soil640/Chelate.html (accessed May 2006). 6. Brown, T. L.; LeMay, H. E., Jr.; Bursten, B. E.; Burdge, J. R. Chemistry: The Central Science, 9th ed; Prentice-Hall: Upper Saddle River, NJ, 2003; p 956. 7. Hill, J. W.; Petrucci, R. H.; McCreary, T. W.; Perry, S. S. General Chemistry, 4th ed; Prentice-Hall: Upper Saddle River, NJ, 2005; p 931. 8. Ebbing, D. D.; Gammon, S. D. General Chemistry, 7th ed; Houghton-Mifflin Co.: Boston, 2002; p 1071. 9. Brady, James E.; Russell, Joel W.; Holum, John R. Chemistry: Matter and Its Changes, 3rd ed; Wiley and Sons: New York, 2000; p 917. 10. Cotton, F. Albert; Wilkinson, G.; Murillo, Carlos A.; Bochmann, Manfred. Advanced Inorganic Chemistry, 6th ed; Wiley and Sons: New York, 1999; p 27. 11. Shriver, D. F.; Atkins, P. Inorganic Chemistry, 3rd ed; Freeman and Company: New York, 1999; p 243. 12. Swaddle, T. W. Inorganic Chemistry—An Industrial and Environmental Perspective; Academic Press: San Diego, 1997; p 248. 13. Spike, C. G.; Parry, R. W. J. Am. Chem. Soc. 1953, 75, 3770. 14. Cotton, F. A.; Wilkinson, G. Advanced Inorganic Chemistry— A Comprehensive Text, 4th ed.; Wiley and Sons: New York, 1980; p 71. 15. Leussing, D. L.; Gallagher, P. K. J. Phys. Chem. 1960, 64, 1632. 16. Atkinson, G.; Bauman, J. E. Inorg. Chem. 1962, 2, 66. 17. Myers, R. Thomas. Inorg. Chem. 1978, 17, 952. 18. Chung, Chung-Sun. Inorg. Chem. 1979, 18, 1321. 19. Chung, Chung-Sun. J. Chem. Educ. 1984, 61, 1062.
Editor’s Note This is the last Overhead Projector Demonstrations column edited by Doris Kolb prior to her death on December 20, 2005. Doris began this column in 1987 ( J. Chem. Educ. 1987, 64, 348–351) and was its sole editor. Since 1987, 71 columns and 126 demonstrations have appeared. She also edited the Chemical Principles Revisited column from 1977 to 1979 and served as a member and as chair of the JCE Board of Publication over a 20-year period. In her inaugural Overhead Projector Demonstrations column, Doris mentioned that she could remember the first time she saw a demonstration on an overhead projector. She
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was impressed by its “clarity, beauty, and simplicity”. That experience inspired her to begin doing overhead demonstrations. Doris’s memory was better than mine. I cannot remember the first time I saw a demonstration on an overhead, but I can vividly remember seeing Doris do demonstrations on an overhead. She was a true master whose main goals were to help students to learn and to help teachers to achieve similar mastery. In her memory, reactions from this article are shown on the cover of this issue as they would appear on the stage of an overhead projector. Doris, all of us miss you.
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