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 (j^o) 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:

j = j^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),

j = j^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 = j _(ox) - j _(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 DG of the reaction by the expression

nFE = - DG

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

DG = - 2.3RT Iog K

It follows then that

nFE = 2.3RT log K

At 25 C (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:

Concentrated sulfuric acid

SO₄² + 4H⁺ + 2e = H₂SO₃ + H₂O

SO₄² + 8H⁺ + 6e = S + 4H₂O

Nitric acid

NO₃ + 4H⁺ + 3e = NO + 2H₂O

NO₃ + 3H⁺ + 2e = HNO₂ + H₂O

NO₃ + 2H⁺ + e = NO₂ + H₂O

NO₃ + 10H⁺ + 8e = NH₄⁺ + 3H₂O

Manganese compounds

MnO₄ + 8H⁺ + 5e = Mn²⁺ + 4H₂O

MnO₄ + 2 H₂O + 3e = MnO₂ + 4OH

MnO₄ + e = MnO₄²

Chromium compounds

Cr₂O₇² + 14H⁺ + 6e = 2Cr³⁺ + 7H₂O

CrO₄² + 4H₂O + 3e = Cr(OH)₆³ + 2OH

Hydrogen peroxide

H₂O₂ + 2e = 2OH

H₂O₂ + 2H⁺ + 2e = 2H₂O

2H⁺ + O₂ + 2e = H₂O₂

2H₂O + O₂ + 2e = H₂O₂ + 2OH

EXPERIMENTAL PART

1. Transfer of an ion to a higher oxidation state

Take 6-8 drops of chromium (III) nitrate and add excess of NaOH. When the precipitate of chromium hydroxide dissolves add 3-4 drops of 3% H₂O₂ solution. Heat the mixture until the color turns yellow. Write down and balance a corresponding redox-reaction.

2. Redox properties of hydrogen peroxide

a). Take 3 drops of KI solution, add 2 drops of diluted sulfuric acid and 3% H₂O₂ solution. Add some starch solution to indicate the evaluation of iodine. Write down and balance a corresponding redox-reaction.

b). Take 6 drops of KMnO₄ solution, add 2 drops of diluted sulfuric acid and dropwise 3% H₂O₂. Observe the evaluation of oxygen. Write down and balance a corresponding redox-reaction.

3. Oxidizing properties of potassium permanganate

Fill 3 test tubes with 5 drops of KMnO₄ each. Add 2 drops of diluted sulfuric acid into the first test tube, 2 drops of distilled water into the second, and 2 drops of sodium hydroxide into the third one. Add Na₂SO₃ solution into each test tube until the change of their color. Write down and balance corresponding redox-reactions.

4. Oxidation of cations of d-elements

Take 2 drops of a manganese (II) salt in a test tube, then add 5-6 drops of nitric acid and some crystals of NaBiO₃. Observe the appearance of a pink color of HMnO_4.. Write down and balance a corresponding redox-reaction.

5. Reducing properties of p-elements

Take 3-4 drops of a solution of SnCl₂ in a test tube, add NaOH until the formed precipitate dissolves, and 2-3 of a bithmuth (III) salt solution. Observe the black precipitate of elementary bithmuth. Write down and balance a corresponding redox-reaction.

QUESTIONS AND PROBLEMS

1. Complete the equations of the following reactions and balance them:

(a) K₂S + KMnO₄ + H₂SO₄ = S + ....

(b) KI + K₂Cr₂O₇ + H₂ SO₄ = I₂ + ...

( c) K MnO₄ + H₂O₂ = ...

2. Indicate the direction in which the following reactions can proceed spontaneously:

(a) H₂O₂ + HClO = HCl + O₂ + H₂O

(b) H₃PO₄ + 2HI = H₃PO₃ + I₂ + H₂O

3. Can a salt of iron (III) be reduced to a salt of iron (II) in an aqueous solution by (a) potassium bromide, (b) potassium iodide?

4. Using the table of standard electrode potentials, calculate the equilibrium constants for the following reactions:

(a) Zn + CuSO₄ = Cu + ZnSO₄

(b) Sn + Pb(CH₃COO)₂ = Sn(CH₃COO)₂ + Pb

Experiment 5

Date: 2015-01-12; view: 926