MODULE 1

Inorganic Chemistry

 

7.2.1.       Redox

 

 

Oxidation and reduction in terms of electron transfer.  The concept and use of oxidation numbers.  The idea that the oxidation number of an element changes in a redox reaction.

 

 

Oxidation and Reduction

Oxidation was originally definedas the gaining of oxygen or the loss of hydrogen from an element or compound.  Reduction was the reverse of this.

          PbO  +  CO    Pb  +  CO2

          2H2S  +  O2    2H2O  +  2S

 

N.B.  When oxidation takes place in a chemical reaction reduction will also take place at the same time in the reaction - hence the term redox (for reduction and oxidation).

The present view of redox is that involving electron transfer.

Oxidation involves a loss of electrons

Reduction involves a gain of electrons.

 

(Remember L.E.O.  or   O.I.L.R.I.G.)

 

                                                                OXIDATION

          Sn2+  +  2Fe3+    Sn4+  +  2Fe2+

                                                                                REDUCTION

The Sn2+ has lost 2 electrons and is therefore oxidized.

          Sn2+  Sn4+  +  2e- 

The Sn2+ is the electron donor or REDUCING  AGENT.

 

Each Fe3+ has gained an electron and is therefore reduced.

          Fe3+  +  e-    Fe2+

 

 The Fe3+ is the electron acceptor  or OXIDIZING  AGENT .

 

EXERCISE. 1       Complete the following:

                                                                OXIDATION

          Fe  +  Cu2+    Fe2+  +  Cu

                                      REDUCTION

  Cu2+ is the                              agent

  Fe is the                         agent

 

          2Mg  +  O2    2Mg2+  + O2-

 

  O2  is the                        agent

  Mg is the                         agent

 

Oxidation Number

These can provide a useful guide to the extent of oxidation or reduction in a species.

Oxidation numbers are determined as follows:

1.       The oxidation numbers of monatomic ions are the charges on the ions

ion

Na+

Cl-

Fe3+

S2-

 

Ox. Num.

+1

-1

+3

-2

 

 

2.       For non-ionic compounds the following rules may be applied.

A  Elements have oxidation numbers of zero (0)

 

B  Hydrogen is generally +1 in its compounds except when joined to a metal e.g. NaH when it is -1

                   NaH           

Ox. Num.    +1 -1

 

C  Oxygen is generally -2 in its compounds, except for H2O2 (where Ox. Num. is -1) and F2O (where Ox. Num. is +2).

 

D  In neutral molecules the sum of the Ox. Num.'s is zero and in ions the sum of the Ox. Num.'s is the same as the chargew on the ion.

         

 

CH4

NO2

SO42-

NH4+

Ox. Num.

-4 +1

+4 -2

+6 -2

-3 +1

 

E  In a compound the more electronegative atom has the negative oxidation number and the less electronegative atom has the positive oxidation number.

 

                   F2  +  1/2O2    F2O

 

Ox. Num.   0                 0                                   -1 +2

 

EXERCISE. 2       Add oxidation numbers to complete the following:

 

 

NaCl

HCl

CaCl2

Al2O3

CO

Ox. Num.

 

 

 

 

 

 

CO2

NaOH

OH-

MnO4-

Cr2O72-

Ox.Num.

 

 

 

 

 

 

Oxidation occurs when the oxidation number is increased  (or becomes less negative).

reduction occurs when the oxidation number is decreased (or becomes more negative).

 

                             S  +  O2    SO2

Ox. Num.              0        0                                 +4 -2

The oxidation number of S  has increased from 0 to +4 (i.e. it has been oxidized) while the oxidation number of O has decreased from 0 to -2 (i.e. it has been reduced).

 

EXERCISE. 3       Show using oxidation numbers whether the following is a redox reaction or not.

 

                   SO2  +  OH-    HSO3-

 

 

 

Balancing of redox equations from given half equations and for reactions where the reactants and products are specified (omitting spectator ions).

 

 

Redox Reactions

 

Many chemicals in aqueous solution are oxidizing agents (e.g. MnO4-, Cr2O72-, Cl2, Fe3+ etc.) while others are reducing agrents (e.g. SO32-, Sn2+, Fe2+ etc.).

It is therefore possible to have many redox reactions involving combinations of any pair of oxidant and reductant.

To balance these types of reactions the following procedure should be adopted.

 

1.       Write down the half equations for the oxidation and reduction reactions, including electrons.

 

2.       Adjust half equations so that the electrons lost = electrons gained.

 

3.       Add half equations together cancelling where appropriate.

 

e.g.  Reaction between Fe3+ and I- ions in solution.

 

1.       Fe3+  +  e-    Fe2+

          2I-    I2  +  2e-

 

2.       2Fe3+  +  2e-    2Fe2+

          2I-    I2  +  2e-

 

3.       2Fe3+  +  2I-    2Fe2+  +  I2

 

Oxidizing Action of KMnO4 and K2Cr2O7 in acid solution

 

A  KMnO4

Reaction between KMnO4 and Na2SO3 in acid conditions.

1.       MnO4-  +  8H+  +5e-    Mn2+  +  4H2O

          (purple)                                                     (colourless) 

          SO32-   +  H2O  SO42-  +  2H+  +  2e-

         

2.       2MnO4-  +  16H+  +  10e-    2Mn2+  +  8H2O 

          5SO32-   +  5H2O  5SO42-  +  10H+  +  10e-

3.  2MnO4-  +  5SO32-  +  6H+    2Mn2+  +  5SO42-  +  3H2O

 

K2Cr2O7

Reaction between K2CrO7 and H2S in acid conditions.

1.       CrO72-  +  14H+  +  6e-    2Cr3+  +  7H2O

          (orange)                                                       (green)

          H2S    S  +  2H+  +  2e-

 

2.       CrO72-  +  14H+  +  6e-    2Cr3+  +  7H2O

          3H2S    3S  +  6H+  +  6e-

 

3.       CrO72-  +  8H+  + 3H2S    2Cr3+  +  3S  +  7H2O

 

 

 

The use of standard redox potentials as a means of predicting the feasibility of reactions (no formal treatment of cells or the hydrogen electrode is required in this module).

 

 

Standard Electrode Potential  E¿

                  

 

When a metal is placed in a solution  of its ions an equilibrium is established between the tendency of the metal to loose electrons and pass into solution as ions, and the opposing tendency for the ions in solution to gain electrons and be deposited on the metal.

          M (s)    Mn+ (aq)  + ne-

The metal aquires an electrical charge (usually negative) and a potential difference is set up between the metal and the solution called the electrode potential of the metal.

The size of this electrode potential will depend on the position of this equilibrium and this depends on;

1.  the metal concerned

2.  the concentration of the metal ion in solution

3.  temperature

The temperature and concentration are therefore standardised at 25oC and 1 molar respectively.  Under these conditions the electrode potential is known as the standard electrode potential, E¿.  (For a gas/gas ion equilibrium the gas pressure is standardised as 1 atmosphere).

                   E¿  is in volts at 298K

By convention standard electrode potentials are written as

oxidised state  +  ne-    reduced state

i.e. as reduction potentials.  The voltage produced is for the reaction as written, left to right. 

 

Predicting the Feasibility of Reactions

 

A table of standard redox potentials can be used to predict the likelyhood that a reaction will take place.

Positive E¿ values indicate  systems that have a tendency to gain electrons (i.e. oxidising agents).  Negative E¿ values indicate systems with a tendency to lose electrons (i.e. reducing agents).

 

1.  Reaction of Fe2+ (aq) with halogens

          2Fe2+  + Cl2    2Fe3+  +  2Cl-

                                                                                      E¿

          2Fe2+    2Fe3+  +  2e-                                  -0.77volts

          Cl2  +  2e-   2Cl-                               +1.36volts

          2Fe2+  + Cl2    2Fe3+  +  2Cl- +0.59volts

 

Overall the value of E¿ is +ve so the reaction will take place.

 

•   Predict the feasibility of Fe2+ (aq) reacting with Br2 and I2.