Introduction

When a salt is dissolved in water, the resulting solution can be either acidic, basic, or neutral depending on the ions contained in the salt. If the anion (negative ion) is that of a weak acid, then it will react with the water to form some un-ionized acid and hydroxide ion.

For example:
 

CH3COO- (aq)+ H2O ® CH3COOH(aq) + OH- (aq)

A basic solution with a pH greater than 7 will result. If the cation (positive ion) is that of a weak base, then it will react with the water to form un-ionized base and hydronium ion.

For example:
 

NH4 + (aq) + H2O ® NH3 (aq) + H3O+ (aq)

An acidic solution with a pH less than 7 will result. If the salt contains both the cation of a weak base and anion of a weak acid then the pH of the resulting solution will depend on the relative strengths or solubilities of the related acid and base. These reactions, i.e., the reaction of an ion with water, are called hydrolysis reactions. A solution of NaCl(aq) is neutral. Group I and Group cations are not Bronsted acids and will not undergo hydrolysis. Cl- is the conjugate base of the strong acid HCl, and it will not react with water. NaCl does not undergo hydrolysis, and its solution as well as any solution made from a Group I or Group II cation and the conjugate base of a strong acid will be neutral.

A buffer solution is a solution that resists a change in its pH upon the addition of small quantities of either an acid or a base. An example of such a solution is one containing a weak acid and the salt of the weak acid, or a solution of a weak base and the salt of the weak base. A solution containing NH₄ Cl with NH₃ will be a buffer solution.

We will look at a buffer consisting of a weak acid and its conjugate base, CH₃COOH(aq) and CH₃COO- (aq). The addition of a small amount of a strong acid will increase the hydronium ion present in the solution. This amount of hydronium ion will react with an equal amount of the anion of the weak acid to form un-ionized weak acid.
 

CH3COO- (aq) + H3O+ (aq) ®CH3COOH(aq) + H2O

This removes the added hydronium ion and changes the equilibrium amounts of the weak acid and its anion.

The addition of a small amount of strong base will increase the hydroxide ion present in the solution.
 

CH3COOH(aq) + OH- (aq) ® CH3COO- (aq) + H2O

This amount of hydroxide ion will react with an equal amount of the weak acid to form an increase in the amount of the anion of the weak acid (which is a weaker base than the hydroxide). This removes the added hydroxide ion and changes the equilibrium amounts of weak acid and its anion. In this way the increased amounts of hydrogen or hydroxide ion are replaced by the weaker acid molecule or weaker base ion.

Equipment
 

8 · 100 or 250 mL beakers	1· 600 mL beaker
analytical balance	100 mL graduated cylinder
pH meter	hot plate.

Chemicals
students 300 mL distilled or deionized water

Spill/Disposal
Spill/Disposal: A
 

[Al(H2O)6 ]Cl3(aluminum Chloride hexahydrate)	NH4Cl (Ammonium Chloride)
Na2CO3 (Sodium Carbonate)	CH3 COONa· 3H2O(Sodium Acetate trihyd)
NaHCO3 (Sodium Bicarbonate)

Spill/Disposal: B1
 

0.10 M CH3COOH (Acetic Acid)	0.10 M HCl (Hydrochloric Acid, aqueous)
0.10 M NaOH (Sodium Hydroxide, aqueous)

Procedure
Part I. Hydrolysis of Salts

1. Obtain 350 mL of distilled or deionized water in a 600 mL beaker and heat it to boiling on a hot plate. Remove the beaker from the hot plate and allow it to cool to room temperature without swirling or vigorously stirring the water.
2. While the water is cooling, calculate the weight of NaHCO₃, Na₂CO₃, [Al(H₂O)₆ ]Cl₃ , and NH₄Cl that must be dissolved in 50 mL of water to make a 0.10 M solution of each of these salts. Check your values with your instructor before you weigh out each of the quantities.

3. With a graduated cylinder, measure 50 mL of the cooled water and place it into a beaker that has been rinsed well with deionized water. Do this three more times. Into each of the 50 mL portions of water place one of the salt samples, being sure you mark which salt has been added to which beaker. Keep the rest of the cooled water for Part

4. Carefully stir each salt until it has completely dissolved in the water. Take the pH of each solution. Your instructor will have standardized the pH meter. Do not take a reading until the electrode is immersed in the salt solution. Turn the switch back to stand-by before you remove the electrode from the solution.

5. When you have finished all your pH measurements you can discard the solutions down the sink drain.

6. Show net ionic equations for the hydrolysis reactions.

Part II. Buffering Action of an Acetic Acid-Sodium Acetate Salt

1. Put 50 mLs of the cooled water into each of two 50 mL beakers that have been rinsed with deionized water. Take the pH of each solution. To one of the beakers add 5 mL of 0.10 M HCl solution and again take the pH. To the other beaker add 5.0 mL of 0.10 M NaOH solution and take the pH after the addition.

2. Measure out 100 mL of the 0.10 M CH₃COOH solution. Calculate the weight of CH₃COONa · 3 H₂O (s) that you will need to make this solution 0.10 M in CH₃COONa.

(C₂H₃O₂Na · 3 H₂O (s)) Note that the molecular weight of the salt must include any waters of hydration. Check your value with that of the instructor. Weigh out the needed amount of salt and add it to the solution. Carefully stir the solution until all of the salt has dissolved. Divide the solution into two 50 mL portions. You can discard any left over amount.

3. Take the pH of each solution. To one of the solutions add 5.0 mL of the 0.10 M HCl solution. After mixing the solution, take the pH. To the other solution of the buffer add 5.0 mL of the 0.10M NaOH solution. After mixing take its pH. When finished you can discard the solutions down the sink drain.

Disposal

All solutions from this experiment may be disposed of into the sink.