A combination reaction may be denoted in the manner:
A + B -> C
Either A and B or both A and B must be in the elemental form for such a reaction to be a redox reaction. All combustion reactions, which make use of elemental dioxygen, as well as other reactions involving elements other than
dioxygen, are redox reactions.
Example
$CH_4(g) + 2O_2 -> CO_2(g) + 2H_2O(l)$
2.Decomposition reactions
This is opposite of the combination reaction. Decomposition reaction leads to the breakdown of a compound into two or more components at least one of which must be in the elemental state. Example
$2KClO_3 -> 2KCl + 3O_2$
Also all decomposition reactions are not redox reactions. For example, decomposition of calcium carbonate is not a redox reaction
$CaCO_3 -> CaO +CO_2$
3. Displacement reactions
In a displacement reaction, an ion (or an atom) in a compound is replaced by an ion (or an atom) of another element.
X + YZ → XZ + Y
Displacement reactions fit into two categories: metal displacement and non-metal displacement. (a) Metal Displacement
A Metal in a compound is replaced by another metal in the uncombined state
$CuSO_4(aq) + Zn(s) -> Cu(s) + ZnSO_4$
$V_2O_5 (s) + 5Ca (s) ->2V (s) + 5CaO (s)$
Here the reducing metal is a better reducing agent than the one that is being reduced which evidently shows more capability to lose electrons as compared to the one that is reduced.
Activity Series for Metals (Tendency to act as reducing agents)
Li > K > Na > Mg> Al > Zn> Fe > Pb > H > Cu> Ag > Hg> Au
(b) Non-metal displacement
The non-metal displacement redox reactions include hydrogen displacement and a rarely occurring reaction involving oxygen displacement.
All alkali metals and some alkaline earth metals (Ca, Sr, and Ba) which are very good reductants, will displace hydrogen from cold water
$2Na(s) + 2H_2O(l) -> 2NaOH(aq) + H_2(g)$
$Ca + 2H_2O -> Ca(OH)_2 + H_2$
Less active metals such as magnesium and iron react with steam to produce dihydrogen gas
$2Mg(s) + 2H_2O(g) -> Mg(OH)_2 + H_2(g)$
Many metals, including those which do not react with cold water, are capable of displacing hydrogen from acids. Dihydrogen from acids may even be produced by such metals which do not react with steam. Cadmium and tin are
the examples of such metals
$Zn(s) + 2HCl(aq) -> ZnCl_2 (aq) + H_2 (g)$
Here, the reactivity of metals is reflected in the rate of hydrogen gas evolution, which is the slowest for the least active metal Fe, and the fastest for the most reactive metal, Mg. Very less native state such as silver (Ag), and gold (Au) do not react even with hydrochloric acid. This is also clear from the below table
Activity Series for Metals (Tendency to act as reducing agents)
Li > K > Na > Mg> Al > Zn> Fe > Pb > H > Cu> Ag > Hg> Au
We have activity series for Halogens also
F > Cl > Br > I
Fluorine is so reactive that it attacks water and displaces the oxygen of water
$2H_2O (l) + 2F_2 (g) -> 4HF(aq) + O_2(g)$
It is for this reason that the displacement reactions of chlorine, bromine and iodine using fluorine are not generally carried out in aqueous solution. On the other hand, chlorine can displace bromide and iodide ions in an aqueous solution as shown below:
$Cl_2 (g) + 2KBr (aq) -> 2 KCl (aq) + Br_2 (l)$
$Cl_2 (g) + 2KI (aq) -> 2 KCl (aq) + I_2 (s)$
4.Disproportionation reactions
They are a special type of redox reactions.
In a disproportionation reaction an element in one oxidation state is simultaneously oxidised and reduced.
One of the reacting substances in a disproportionation reaction always contains an element that can exist in at least three oxidation states.
The element in the form of reacting substance is in the intermediate oxidation state; and both higher and lower oxidation states of that element are formed in the reaction
Example
+1 –1 +1 –2 0
$2H_2O_2 (aq) -> 2H_2O(l) + O_2(g)$
Here the oxygen of peroxide, which is present in –1 state, is converted to zero oxidation state in $O_2$ and decreases to –2 oxidation state in $H_2O$
$Cl_2 (g) + 2 OH^– (aq) -> ClO^-(aq) + Cl^-(aq) + H_2O (l)$
The above reaction describes the formation of household bleaching agents. The hypochlorite ion ($ClO^–$) formed in the reaction oxidises the colour-bearing stains of the substances to colourless compounds.
