In the Classroom edited by
Tested Demonstrations
Ed Vitz Kutztown University Kutztown, PA 19530
When A + B ⴝ B + A submitted by:
Erling Antony,* Lindsay Mitchell, and Lauren Nettenstrom Department of Chemistry, Arrowhead High School, 700 North Avenue, Hartland, WI 53029; *
[email protected] checked by:
David Speckhard Department of Chemistry, Loras College, Dubuque, IA 52001
Students frequently ask, “Does it matter which order you add the reactants?” During basic inorganic chemistry our answer is typically, “Always add a concentrated acid to water but otherwise the order usually does not matter.” This demonstration involves one instance when the same quantities of reactants are used in identical systems with very different results. The first portion of the demonstration gives results that are easily explained in a high school general chemistry class; later results may initially puzzle a college general chemistry or AP class, although students should achieve a satisfactory explanation after analyzing the system.
Discussion and Results When the sodium carbonate solution is placed in the flask and the HCl is added dropwise, the only gas collected is that displaced by the increasing solution volume. The equation for the reaction is as follows: H+ + CO32᎑ → HCO3᎑ When the hydrochloric acid is placed in the flask and the sodium carbonate is added dropwise, the results are quite
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Two standard gas collection apparatus are set up, each using a 250-mL Erlenmeyer flask, magnetic stirring device, 500-mL graduated cylinder, pneumatic trough, and connecting tubes (Fig. 1). A 50-mL buret is substituted for the thistle tube in each. The first flask contains 50 mL of 1.0 M HCl and the second contains 50 mL of 1.0 M sodium carbonate. The burets are filled with the opposing reactant and the systems are sealed for gas collection. Solution is slowly added to each flask while the contents are stirred. Pause the addition after 25 mL, observe the system, and record the volumes of gas collected. This is a good time for discussion of the processes and to predict future behaviors. Continue adding the solutions from the burets. After an additional 1 to 2 mL of solution have been added, observe the system and again record the volumes of gas. Continue adding the reactants until 50 mL of each has been added. Again observe the system and record volumes of gas.
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HCl added / mL Figure 2. Gas production vs volume of HCl added to Na2CO3 solution.
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NaCO3 added / mL Figure 1. Photograph of the apparatus.
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Figure 3. Gas production vs volume of Na2CO3 added to HCl solution.
Journal of Chemical Education • Vol. 77 No. 9 September 2000 • JChemEd.chem.wisc.edu
In the Classroom
different. Carbon dioxide is produced owing to the excess of acid (this is the common baking soda–vinegar reaction [1, 2]), which results in carbonic acid formation and decomposition as follows: 2 H+ + CO32᎑ → H2CO3 → H2O + CO2 Gas production for the first process is steady as expected, with a slight increase as one nears the end point (Fig. 2). This, again, would be expected. These results differ greatly from the reverse process. Figure 3 shows a steady production of gas until 25 mL of sodium carbonate solution has been added. At this point all acid has been neutralized and no additional carbon dioxide can be produced. The slope of the linear portion of the graph (21.0 mL of gas per 1 mL of sodium carbonate) is in fair agreement with the predicted value of 23.4 mL of gas per 1 mL of sodium carbonate (from atmospheric conditions during the experiment). The y-intercept probably reflects the gas’s combined solubility in the acid and the water to which it has been exposed. An interesting phenomenon occurs after 25 mL of sodium carbonate is added to the HCl. Water begins to siphon back into the Erlenmeyer flask. Upon seeing this, many students assume the reaction is starting to “go backwards”. After some consideration, however, they realize this would require the sodium carbonate solution to re-form and return to the buret. This results from the excess aqueous carbonate ions reacting with the water, which causes hydrogen carbonate ions and hydroxide ions to form. The equation is (3) HOH + CO32᎑ → HCO3᎑ + OH᎑ The hydroxide ions react with carbon dioxide as OH + CO2 → HCO3 ᎑
᎑
As the carbon dioxide is dissolved in the water and reacts with the hydroxide ions, the pressure decreases and water is forced back through the tube into the flask. This is similar to the reaction of calcium hydroxide with carbon dioxide from human breath, another common demonstration (4 ). The results clearly demonstrate that order of addition can be important while simultaneously demonstrating the stepwise neutralization of a dibasic ion and the solubility of
carbon dioxide in base. Another example of when the order of addition is important is the titration of phosphoric acid. Phosphoric acid may be titrated with sodium hydroxide to give two visible end points, but the reverse titration is impossible (5). Again, potassium nitrate may be reacted with perchloric acid to yield two visibly different products (6 ). However, few other processes are as easily observed and as simple to set up as the one in this demonstration. Comments We invite students to surround the demonstration table during the procedure. This helps focus their attention and improves visibility. In a large lecture hall, a colored water solution should be used in the pneumatic trough and graduated cylinder. An indicator may be added to the receiver flask if teachers wish to emphasize the pH dependence of the reactions. We like this demonstration because it not only illustrates the fact that the order of addition can be important, but it also assists with the concepts of excess or limiting reactant, the stepwise neutralization of an acid or base, and the unexpected but explainable siphoning phenomenon. This demonstration could be easily expanded to serve as a lab activity or added onto other lab experiments. If this is done, the volume of gas should be read after the addition of each milliliter of reactant. However, we find the procedure to be more thought provoking when used as a demonstration. Literature Cited 1. Shakhashiri, B. Chemical Demonstrations; University of Wisconsin Press: Madison, WI, 1992; Vol. 3, p 96. 2. Kelter, P.; Crouse, D. J. Chem Educ. 1985, 62, 1108. 3. Brown, T. L.; LeMay, H. E. Chemistry, 4th ed.; Prentice Hall: Englewood Cliffs, NJ, 1988; p 752. 4. Metcalfe, H. C.; Williams, J. E.; Castka, J. F. Modern Chemistry; Holt, Rinehart and Winston: New York, 1982; p 382. 5. Friesen, R. J. Chem 13 News 1981, Dec, 8 (answer to question 181). 6. Harris, W. E.; Kratochvil, B. An Introduction to Chemical Analysis; Saunders: Philadelphia, 1981; p 188.
JChemEd.chem.wisc.edu • Vol. 77 No. 9 September 2000 • Journal of Chemical Education
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