This demonstration shows a colorful reaction that is a good introduction to Le Châtelier’s principle, solubility, stoichiometry, neutralization reactions, and reaction rates. The demonstration uses readily available chemicals, is easy to set up and perform, and will engage your students with brilliant colors. Each color change is a visual indicator of a shift in equilibrium. This demonstration can also be used as a limiting reactant demonstration and kinetics demonstration. See the product links for instructions.
Wear safety goggles and chemical-resistant gloves while performing this demonstration. Universal indicator solution may stain skin and clothing. It contains ethanol, which is a flammable solvent. Keep flames and sources of ignition away from this solution. Sulfuric acid is corrosive to skin and eyes.
The final solution will be a slightly acidic solution of magnesium sulfate and can be disposed of down a drain with excess water.
Milk of Magnesia, an over-the-counter medication frequently used as a laxative and antacid, is a suspension of magnesium hydroxide, Mg(OH)2, in water. Magnesium hydroxide is only slightly soluble in water (0.012 g/L). During the demonstration, sulfuric acid (H2SO4) is added to a saturated solution of magnesium hydroxide. The hydroxide ion in the solution is immediately consumed when it reacts with hydrogen ions to produce water. The stress of disrupted equilibrium is relieved as more magnesium hydroxide dissociates. According to Le Châtelier’s principle, as the hydroxide ions react with hydrogen ions to produce molecular water, the hydroxide ions (OH–) decrease, placing a stress on the system. The response to the decrease in concentration is for more Mg(OH)2 to dissociate, re-establishing equilibrium. The process continues until all the magnesium hydroxide is neutralized, leaving only a clear, red solution of magnesium sulfate (due to the low pH resulting in excess sulfuric acid).
When the acid is added, the solution changes from purple (pH 10) to red (pH 4). The color of the solution then slowly changes to orange, yellow, green, blue, and finally back to violet as the magnesium hydroxide slowly dissolves. This process continues until all the magnesium hydroxide is neutralized, leaving only a clear solution of magnesium sulfate, which will be red due to the low pH, a result of the excess sulfuric acid. The final solution is also clear, indicating a solution rather than a suspension.
Mg(OH)2(s) Mg2+(aq) + 2OH–(aq)
Mg(OH)2(aq) + H2SO4(aq) → MgSO4(aq) + 2H2O(l)
Why does the solution change color very quickly when the acid is added? Explain the reactions or processes that are occurring.
The neutralization reaction between the hydroxide ions in solution and the acid is very fast.
Why does the solution slowly return to the original color? Explain the reactions or processes are occurring?
Once all the hydroxide ions in solution are consumed, the magnesium hydroxide slowly dissolves to relieve the disrupted equilibrium as described by Le Châtelier’s principle.
Using Le Châtelier’s principle, explain the color changes observed in the demonstration.
The hydroxide ions (OH–) in solution are neutralized, causing a disruption of the solubility equilibrium for magnesium hydroxide. This causes more of the solid magnesium hydroxide to dissolve. The stress of disrupted equilibrium is relieved as described by Le Châtelier’s principle, which states that if a stress is placed upon a reaction at equilibrium, the reaction shifts to relieve that stress.
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