REDOX REACTION at a glance

 

Oxidation–Reduction reaction (Redox)

 Oxidation–Reduction(redox) reactions are two opposite and complementary reactions which occur simultaneously

Redox reactions have been defined in several ways before attaining a more general and simplified definition.

1        In term of addition of oxygen:  Oxidation is defined as the addition of oxygen to a substance while reduction is defined as the removal of oxygen from a substance. E.g.

            
              Reduction              CuO + C_(s) → Cu_(s) +  CO_(g)     O.A    R.A                       
   

the example above, carbon (C) is oxidized to carbon (II) oxide (CO) while Cupper (II) oxide (CuO) is reduced to metallic Copper (Cu).

   Carbon is removing oxygen from CuO and so is the reducing agent because it causes CuO to become reduced to Cu. CuO  supplies the oxygen atom that causes carbon to become oxidized to CO and so CuO  is the oxidizing agent.

2       In terms of removal of hydrogen:  Oxidation is defined as the removal of hydrogen from a substance while reduction is defined as the addition of hydrogen to a substance                    

                    ——————
               ↓Reduction   ↓
H₂S   +   Cl₂ → S_(s) + HCl(aq)
 R A        O.A    
 ↑  Oxidation  ↑
     ——————

                     

          Similarly in this reaction, H₂S is oxidized to atomic S_(s) due to the removal of hydrogen as chlorine is reduced by gaining or addition of hydrogen. H₂S is action as the reducing agent while Cl₂ is the oxidizing agent.

3        In terms of change in the oxidation number of an element: Oxidation is defined as the increase in oxidation number of an element while reduction is the decrease in oxidation number of an element. 

4.  Definition in terms of electronegative elements: - Oxidation is the addition of an electronegative element to a substance or the removal of electropositive element from a substance while reduction is the removal of electronegative element from a substance or the addition of electropositive element to a substance 

4        In terms of electron transfer: Oxidation is defined as the loss of electrons while reduction is defined as the gain of electrons. 

  When an element loses an electron to become an ion; the O.N increase to a higher number while a gain of electrons by an element will lead to a decrease in O.N of an element. For example,

        ₂₀Ca   → 20Ca²⁺+ 2e^–
        20 protons      20 protons
   20 electrons         18 electrons
    O.N 0 (zero)        O.N (+2)

        

        ₁₇Cl    +   e^–   → 17Cl^–
     17 protons            17 protons
     17 electrons         18 electrons
     O.N 0 (zero)         O.N( –1)

          In other words, as noted from the above examples, loss of electrons means a higher O.N while gain of electrons means a decrease in O.N of an element.

                            Oxidation
               Mg + Cl   →MgCl2       
               R.A   O.A     reduction

Example of redox reactions that occurs generally around us include

1. Photosynthesis

2. Rusting of iron

  Fe(s) + nH₂O → Fe₂O₃.nH₂O_(s)

3.   Combustion

             –——————
            ↓ Oxidation      ↓
          C₄H₁₀ + O2 → CO₂ + H₂O
                              red   ↑
                            ——————

Redox reactions always involve the movement of electrons ( i.e loss and gain of electrons) For example Pb° →Pb²⁺ + 2e^–  

Pb²⁺→Pb⁴⁺ + 2e^–. O.N increased from 0 to +2 to +4 in Pb. i.e., oxidation involves loss of electrons which will lead to increase in O.N.

In contrast reduction involves the gains of electrons which will lead to a decrease in the O.N of the element for example                   S° + e → S^– + e^– → S^2–.

Some examples of redox recitations are

1        Fe_(s)+ S_(s)→ FeS

          R.A        O.A

          In the above reaction Fe losses electrons to become iron (II) ions (Fe²⁺), there is an increase in its O.N from 0 to +2 and so it is the reducing agent. Sulphur on the other hand, gains electrons from the iron, its O.N decreases from 0 to (–2) and so it is the O.N

2. Pb⁺²O^-2 + C⁺²O^-2 → Pb⁰+ CO₂

          In the above example, the O.N of Pb decreased from +2 to 0, so PbO is reduced and so PbO is the oxidizing agent.

          The O.N of C increased from +2 to +4, so CO is oxidized and so CO is the reducing agent.

3        H_2(g)+ O_2(g) → H₂O_(l) 
          R.A           O.A

          O.N of H increased from 0 to (+1) i.e. H is oxidized.

          The above reaction is a combustion reaction, and in this reaction. It is important to note that all combustion reactions are redox reaction with oxygen as the oxidizing agent.

i        AgNO_3(aq) + NaCl_(aq) → AgCl_(s) + NaNO_3(aq)

          The above reaction, is a double decomposition reaction, it is not a redox reaction as there is no change in the O.N number of all the element s involved. Another non-redox reaction is a neutralization reaction, there is no change in the O.N of element involved in the reaction.

  ii  KOH_(aq) + HCl_(aq) → NaOH_(aq) + H₂O_(l)    (neutralization reaction)

          

Oxidizing and reducing agents (in summary)

OXIDIZING AGENT 1.      Supplies Oxygen 2.      Removes Hydrogen 3.      Decreases in oxidation number 4.      Gains electrons

REDUCING AGENT Supplies hydrogen Removes Oxygen Increases in oxidation number Loss electrons

Test for oxidizing agents

To common test or reactions that are used to test for an oxidizing agent involves the action on iron (II) chloride and hydrogen sulphide.

a)       Reaction with FeCl₂

          When an oxidizing agent is added to green iron (II chloride; the iron (II) ions become oxidized to yellow or brown Fe³⁺.

          Fe²⁺ →     Fe³⁺ + e^–

          green         yellow/brown

b)      Reaction with hydrogen sulphide

          When hydrogen sulphide is bubbled through a solution of an oxidizing agent, the sulphide ions S^2– becomes oxidized to elemental sulphur; and this is seen or observed as yellow deposits sulphur,                    i.e. S^2– → S(s) + 2e^–.

Test for reducing agents

Two commonest reagents that are used to test for a reducing agent are

1        Acidified potassium tetraoxomanganate(VI) (KMnO₄) and acidified potassium heptaoxodichromate(I) (K₂Cr₂O₇).

a)   Action of potassium hyptaoxodichromate (VI) (K₂Cr₂O₇)

  When acidified potassium heptaoxodichromate (VI) (K₂Cr₂O₇) is added to a sample of a reducing agent, its colour changes from orange to green, due to the reduction of the dichromate (VI) ion  (Cr⁶⁺) (orange) to chromium (III) (Cr³⁺) ion green

     Cr⁶⁺  +  3e^– → Cr³⁺
  Orange                green

 b) Test using acidified potassium tetraoxomangane(VI) (KMnO₄)

  When acidified potassium tetraoxomanganate (VII) to a sample of reducing agent, the purple colour changes to colourless: due to the reduction of the manganate ion from (+7) which is purple to (+2) which is colourless and a more stable oxidation state.

MnO₄^- + 8H⁺ + 5e^– →Mn²⁺ + 4H₂O

                  purple                       colourless         

Mn⁷⁺ + 5e^– → Mn²

  purple              colourless

                   This reaction is reversible as the purple colour is restored when an oxidizing agent is reintroduced into the mixture.

                  Mn²⁺ + 5e →Mn⁷⁺   
               colourless        purple

THEORY QUESTIONS 

1(a)State what you will see on

i) bubbling SO2 into acidified KMnO4 solution 

ii). 

(b)(i).Write the ionic equation for the reaction between zinc powder and silver trioxonitrate (V) solution 

(ii). Which substance in bi above is I. Oxidized  II. Reduced