1. Dissociation of weak electrolytes

a) Fill two test tubes with 3-4 drops of diluted acetic acid and add 1 drop of an indicator (methylorange). Add some solid sodium acetate into one of the test tubes. Compare colors of the solutions in the test tubes. Explain the results of the experiment.

b) Repeat the experiment with the solution of ammonium hydroxide. Use phenolphtalein as indicator and add ammonium chloride.

2. Formation of weak electrolytes

a) Take 3-4 drops of a CH₃COONa solution in a test tube and add some diluted HCl. Check the smell before and after the experiment. Explain the results of the experiment.

b) Repeat the experiment with NH₄Cl and NaOH respectively.

3. Amphoteric hydroxides

a) Take 5-6 drops of a solution of any chromium (III) salt and add dropwise a diluted solution of sodium hydroxide until the precipitate of chromium hydroxide is formed. Divide the precipitate of Cr(OH)₃ into two parts and dissolve them in NaOH and HCl respectively.

b) Repeat the same experiment using any zinc salt.

QUESTIONS AND PROBLEMS

1. The dissociating constant of butyric acid C₃H₇COOH is 1.510. Calculate the degree of its dissociation in a 0.005 M solution.

2. What is the hydrogen ion concentration [H⁺] in an aqueous solution of formic acid if  = 0.03?

3. Calculate the concentration of acetate ions in 0.1 M solution of acetic acid in presence of 0.01 M HCl.

4. Calculate the ionic strength and the activities of the ions in a solution containing 0.01 molel^-1 of Ca(NO₃)₂ and 0.01 mol/l of CaCl_2.

Experiment 2

HETEROGENOUS EQUILIBRIUM

In a saturated solution of a sparingly soluble strong electrolyte, an equilibrium sets between the precipitate (solid phase) of the electrolyte and its ions in the solution:

Ag₃PO₄ 3 Ag⁺ + PO

This equilibrium state may be characterized by a constant called a solubility product, K_sp:

K_sp (Ag₃PO₄) = [Ag⁺]³ [PO]

When the concentration of one of the ions of an electrolyte in its saturated solution is increased (for instance, by introducing another electrolyte containing common ions), the product of the concentrations of the electrolyte ions becomes greater than K_sp. In this case the equilibrium between the solid phase and the solution shifts towards the direction of formation of a precipitate.

Consequently, the condition for formation of a precipitate is the greater value of the product of the concentrations of the ions belonging to a sparingly soluble electrolyte in comparison with its solubility product.

Conversely, if the concentration of one of the ions in a saturated solution is diminished (for example by combining it with another ion), the product of the ion concentrations will be lower than the value of K_sp, the solution will become unsaturated and the equilibrium between the solid phase and the solution shifts towards the dissolution of the precipitate.

The value of K_sp can be used to calculate the solubility of sparingly soluble electrolytes in water and in solutions containing other electrolytes (solubility can be determined as a molar concentration of a dissolved substance).

The shift of an equilibrium of a chemical reaction can be determined theoretically by calculation of an equilibrium constant of a chemical reaction:

K=

If the K value exceeds 1, the equilibrium of a chemical reaction is shifted towards the forward reaction. In case of K<1 the equilibrium shifts to the direction of a backward reaction.

EXPERIMENTAL PART