Hydrolysis of Salts: pH

Solutions of various salts turn different colors when universal indicator solution is added. Demonstrates the abilities of some salts to alter the pH of an aqueous solution. 

Curriculum Notes 

This demonstration is usually performed when discussing hydrolysis of salts during acid-base equilibrium. Sodium acetate is the salt of a weak acid, acetic acid and a strong base, sodium hydroxide.  The acetate ion reacts with water to establish an equilibrium system 

CH3COO- (aq) + H2O <=> CH3COOH(aq)  + OH-(aq)    

The resultant solution is basic.

This demonstration might also be used for introducing the concept that the equivalence point of a weak acid-strong base titration may not occur at pH 7, but at pH 8.6.  and treating how this might affect the selection of an indicator. Allow about 10 minutes for this demo.

Lead Time 
One day of lead time is required for this project.
Discussion 
  • The four solutions are all 0.1M. They are NaCl, NaNO2, NH4Cl, and NH4NO2.
  • The universal indicator is like a rainbow. Going from acid to base it changes red, orange, yellow, green, blue, and violet, with green being neutral.
  • The color differences in this demo are discernable, but not great. Please keep in mind that approximately 20% of the males in your class will be red-green colorblind so be sure to describe the color differences verbally.
  • Sodium ions and chloride ions are not strong enough acids or bases respectively to hydrolyze water to any observable extent. Therefore, the sodium chloride solution is neutral. (The solution is boiled to remove any acidity caused by dissolved carbon dioxide.)
  • Ammonium ion is a weak acid. (Ka = 5.7 x 10-10) Therefore, the ammonium chloride solution is acidic and turns the indicator orange.
  • Nitrite ion is a weak base (Kb = 1.4 x 10-11) so the sodium nitrite solution turns the indicator blue-green.
  • Ammonium nitrite contains both a weak acid and a weak base, but since the Ka of the ammonium ion is greater than the Kb of the nitrite ion, the indicator shows that the solution is very slightly more acidic than the neutral sodium chloride solution.
Materials 
  • indicator reference bottles containing universal indicator in each of the following buffer solutions: pH 2, pH 7, and pH 10
  • 150 mL bottles containing 0.1 M solutions of each of the following solutions: NaCl, NaNO2, NH4Cl, and NH4NO2
  • 4 600 mL beakers
  • dropper bottle containing universal indicator
  • 4 stir rods
Procedure 
  • Exhibit the three indicator reference bottles. Lay them on the overhead projector. Explain to the class the range of colors that can be expected from universal indicator solution.
  • Pour each of the four solutions into separate beakers. Fill each beaker about halfway.
  • Add 2-3 drops of indicator solution to each beaker and place it on the overhead projector.
  • Compare each of the beakers with each other and the indicator reference bottles to determine its relative pH.
Safety Precautions 

None of the solutions used is particularly hazardous. Wear safety goggles.

© Copyright 2012 Email: Randy Sullivan, University of Oregon Chemistry Department and UO Libraries Interactive Media Group