1. Determination of pH of solutions of some salts

1. Take some small amounts of solid salts Na₂CO₃, Al₂(SO₄)₃, Na₃PO₄, Na₂B₄O₇, NH₄Cl, (NH₄)₂CO₃ or others into different test tubes.

2. Add some distilled water into each test tube and stir until the salts are dissolved.

3. Using a universal indicator, determine pH values of each solution including pH of distilled water itself.

4. Write down the values of pH and the reactions of hydrolysis of the given salts.

2. Shift of the equilibrium of hydrolysis

1. Take 4 drops of antimony chloiride solution and add some distilled water. Pay attention to the precipitation of antimony oxochloride SbOCl. (The formation of SbOCl is due to the decomposition of Sb(OH)₂Cl).

2. Add some drops of HCl until the precipitate is dissolved.

3. Write down two steps of hydrolysis of SbCl₃ and explain the shift of the equilibrium of the process.

3. Irreversible hydrolysis

1. Take 4-5 drops of aluminum sulfate into a test tube and add the same amount of sodium carbonate solution.

2. Make sure that the obtained precipitate is not aluminum carbonate but aluminum hydroxide (use amphoteric properties of Al(OH)₃).

4. Temperature dependence of hydrolysis

Take 4-5 drops of a 1M solution of sodium acetate into a test tube and add 2 drops of an indicator phenolphtalein. Heat the solution. Explain the change in color of the indicator.

QUESTIONS AND PROBLEMS

1. Calculate pH of 0.1 M solution of NaOH (assume the dissociation to be complete).

2. Calculate pH of a 0.01 M solution of acetic acid if the degree of dissociation of the electrolyte equals 0.042.

3. Calculate pH of an ammonium buffer solution prepared by mixing of equal volumes of 0.1 M solution of NH₄OH and 0.01 M solution of NH₄Cl.

4. Which of the salts listed below undergo hydrolysis? Write the net ionic equations and indicate whether aqueous solutions of salts are neutral, acidic or basic. NaCN, KNO₃, K₂S, ZnCl₂, NH₄NO₂, MgSO₄.

5. Calculate hydrolysis constant and degree of hydrolysis in 0.1 M solutions of: a) NH₄Cl; b) Na₂CO₃ (only the first step of hydrolysis should be taken into consideration).

6. When aqueous solutions of Cr(NO₃)₃ and Na₂S are mixed together, a precipitate is formed and a gas is evolved. Write the molecular and net ionic equations of the reaction.

Experiment 4 oxidation-reduction reactions

In oxidation-reduction reactions (redox-reactions), the oxidation number of one or more elements in the reacting substances changes. The loosing of electrons by an atom attended by an increase in its oxidation number is called oxidation; the gaining of electrons by an atom attended by a decrease in its oxidation number is called reduction.

A substance containing an element that undergoes oxidation is called a reducing agent. These are almost all metals and some non-metals (C, H₂ and others, negatively charged ions of non-metals (S², I, N³ and others), cations in intermediate oxidation numbers (Sn²⁺, Fe²⁺ and others), ions containing elements in intermediate oxidation numbers (SO₃², NO₂, SnO₂² and others). In laboratories, such reducing agents as H₂, SO₂, KI, H₃PO₃, H₂S, HNO₂ are usually used.

A substance containing an element that undergoes reduction is called an oxidizing agent. These are atoms and molecules of some non-metals of high activity (O₂, O₃, Cl₂ and others) positively charged metallic ions (Fe³⁺, Cu²⁺, Hg²⁺ and others), particles containing ions in their highest oxidation numbers (MnO₄, NO₃, SO₄², Cr₂O₇², ClO₃ and others). The strongest oxidizing agent is electrical current (oxidation on anode). In laboratories, such oxidizing agents as KMnO₄, K₂Cr₂O₇, HNO₃, H₂SO_{4 (conc.)}, H₂O₂, PbO₂ are used.

Oxidizing and reducing properties of substances are described with the help of electrode potentials of systems.

The standard electrode potential (^is defined as the potential of a given electrode at concentrations (activities) of all the substances participating in the electrode process equal unity.

The dependence of an electrode potential on concentrations of substances participating in electrode processes and on temperature is expressed by the Nernst equation:

 = ^o + 2.3 log

where R - the molar gas constant; T - absolute temperature; F - the Faraday’s constant; n - number of electrons participating in the electrode process; [Ox] - concentration of the oxidized form of a substance; [Red] - concentration of the reduced form of a substance.

In case if T = 297 K (25^oC),

 = ^o + log

The more is the absolute value of redox potential, the stronger are oxidizing properties of the oxidized form.

The possibility of a redox-reaction can be determined from the electromotive force of the reaction (E):

E = _(ox) - _(red)

In case if E > 0, the direct redox-reaction is possible. In case if E < 0, the direct redox-reaction is impossible and the reaction proceeds in the backward direction.

The standard electromotive force E⁰ is related to the standard Gibbs energy G of the reaction by the expression

nFE= - G

On the other hand, G is related to an equilibrium constant K of a reaction:

G = - 2.3RT Iog K

It follows then that

nFE = 2.3RT log K

At 25C (298 K), the last equation acquires the form

log K =

To balance redox-reactions, the half-reaction method is used. In acidic media molecules of water and hydrogen-ions enter redox half-reactions. In alkaline media both water molecules and OH ions are available. In neutral media the left part of the half-equation contains water molecules and the right part contains either H⁺ or OH ions.

Some examples of redox half-reactions: