Before learning how to identify a Redox reaction, one must understand what is meant by Redox reaction. Redox reactions are considered as electron transfer reactions. It is included in both Organic Chemistry and Inorganic Chemistry. It got its name ‘Redox’ because a redox reaction consists of an oxidation reaction and a reducing reaction. Determining the oxidation number is the key point in identifying a redox reaction. This article discusses the types of redox reactions, giving examples for each redox reaction, the half reactions in a redox reaction, and also explains the rules in determining oxidation numbers and the variations in oxidation numbers.
What is a redox reaction
Acid base reactions are characterized by a proton transfer process, similarly oxidation-reduction or redox reactions involve an electron transfer process. A redox reaction has two half reactions, namely oxidation reaction and the reduction reaction. Oxidation reaction involves the loss of electrons and the reduction reaction involves the acceptance of electrons. Therefore, a redox reaction contains two species, oxidizing agent undergoes the oxidization half reaction and the reducing agent undergoes the reducing half reaction. The extent of reduction in a redox reaction is equals to the extent of oxidization; that means, the number of electrons lost from the oxidizing agent equals to the number of electrons accepted by the reducing agent. It is a balanced process in terms of electron exchanging.
How to Identify a Redox Reaction
Find the Oxidation number:
To identify a redox reaction, first we need to know the oxidation status of each element in the reaction. We use the following rules to assign oxidation numbers.
• The free elements, which are not combined with others, have the oxidation number zero. Thus, atoms in H2, Br2, Na, Be, Ca, K, O2 and P4 have the same oxidation number zero.
• For ions that are composed of only one atom (monoatomic ions), the oxidation number equals to the charge on the ion. For example:
Na+, Li+ and K+ have the oxidation number +1.
F–, I–, Cl– and Br– have the oxidation number -1.
Ba2+, Ca2+, Fe2+ and Ni2+ have the oxidation number +2.
O2- and S2- have the oxidation number -2.
Al3+ and Fe3+ have the oxidation number +3.
• The most common oxidation number of oxygen is -2 (O2-: MgO, H2O), but in hydrogen peroxide it is -1 (O22- : H2O2).
• The most common oxidation number of hydrogen is +1. However, when it is bonded to metals in group I and group II, the oxidation number is -1 (LiH, NaH, CaH2).
• Fluorine (F) shows only -1 oxidation status in all its compounds, other halogens (Cl–, Br– and I–) have both negative and positive oxidation numbers.
• In a neutral molecule, the sum of all oxidation numbers equals to zero.
• In a polyatomic ion, the sum of all oxidation numbers equals to the charge on the ion.
• Oxidation numbers need not be only integers.
Example: Superoxide ion (O22-) – Oxygen has the -1/2 oxidation status.
Identify the oxidation reaction and reduction reaction:
Consider the following reaction.
2Ca + O2(g) —> 2CaO(s)
Step 1: Determine the oxidizing agent and the reducing agent. For this, we need to identify their oxidation numbers.
2Ca + O2(g) —> 2CaO(s)
0 0 (+2) (-2)
Both of the reactants have the oxidation number zero. Calcium increases its oxidation state from (0) —> (+2). Therefore, it is the oxidizing agent. Conversely, in Oxygen the oxidation state decreases from (0) —> (-2). Therefore, Oxygen is the reducing agent.
Step 2: Write half-reactions for the oxidation and the reduction. We use electrons to balance the charges in both sides.
Oxidation: Ca(s) —> Ca2+ + 2e ——(1)
Reduction: O2 + 4e —> 2O2- ——(2)
Step 3: Obtaining the redox reaction. By adding (1) and (2), we can obtain the redox reaction. Electrons in the half reactions should not appear in the balanced redox reaction. For this, we need to multiply reaction (1) by 2 and then add it with reaction (2).
(1)*2 + (2):
2Ca(s) —> 2Ca2+ + 4e ——(1)
O2 + 4e —> 2O2- ——(2)
———————————————————————————-
2Ca + O2(g) —> 2CaO(s)
Identifying redox reactions
Example: Consider the following reactions. Which one resembles a redox reaction?
Zn(s) + CuSO4(aq) —> ZnSO4(aq) + Cu(s)
HCl(aq) + NaOH(aq) —> NaCl(aq) + H2O(l)
In a redox reaction, oxidation numbers change in reactants and products. There should be an oxidizing species and a reducing species. If the oxidation numbers of elements in the products do not change, it cannot be considered as a redox reaction.
Zn(s) + CuSO4(aq) —> ZnSO4(aq) + Cu(s)
Zn(0) Cu (+2) Zn (+2) Cu (0)
S (+6) S (+6)
O (-2) O (-2)
This is a redox reaction. Because, zinc is the oxidizing agent (0 —> (+2) and Copper is the reducing agent (+2) —> (0).
HCl(aq) + NaOH(aq) —> NaCl(aq) + H2O(l)
H(+1),Cl (-1) Na(+1), O(-2), H(+1) Na(+1) , Cl (-1) H(+1), O(-2)
This is not a redox reaction. Because, the reactants and the products have the same oxidation numbers. H (+1), Cl (-1), Na(+1) and O(-2)
Types of redox reactions
There are four different types of redox reactions: combination reactions, decomposition reactions, displacement reactions and disproportionation reactions.
Combination reactions:
Combination reactions are the reactions in which two or more substances combine to form a single product.
A + B —> C
S(s) + O2(g) —> SO2(g)
S(0) O(0) S(+4),O(-2)
3 Mg(s) + N2(g) —> Mg3 N2(s)
Mg(0) N(0) Mg(+2), N(-3)
Decomposition reactions:
In decomposition reactions, a compound breaks down into two of more components. It is the opposite of combination reactions.
C —> A + B
2HgO (s) —> 2Hg(l) + O2(g)
Hg(+2), O(-2) Hg(0) O(0)
2 NaH (s) —-> 2 Na (s) + H2 (g)
Na(+1), H(-1) Na(0) H(0)
2 KClO3(s) —> 2KCl(s) + 3O2(g)
Displacement reactions:
In a displacement reaction, an ion or atom in a compound is replaced by an ion or an atom of another compound. Displacement reactions have a wide range of applications in industry.
A + BC —> AC + B
Hydrogen displacement:
All alkali metals and some alkaline metals (Ca, Sr and Ba) replace by hydrogen from cold water.
2Na( s) + 2H2O (l) —> 2NaOH (aq) + H2(g)
Ca( s) + 2H2O (l) —> Ca(OH)2 (aq) + H2(g)
Metal displacement:
Some metals in the elemental state can displace a metal in a compound. For example, Zinc replaces Copper ions and Copper can replace Silver ions. Displacement reaction depends on the place activity series (or electrochemical series).
Zn(s) + CuSO4(aq) —> Cu(s) + ZnSO4(aq)
Halogen displacement:
Activity series for halogen displacement reactions: F2 > Cl2 > Br2 > I2. As we go down the halogen series, the power of oxidizing ability decreases.
Cl2(g) + 2KBr (aq) —> 2KCl (aq) + Br2(l)
Cl2(g) + 2KI (aq) —> 2KCl (aq) + I2(s)
Br2(l) + 2I– (aq) —> 2Br–(aq) + I2(s)
Disproportionation reactions:
This is a special type of the redox reaction. An element in one oxidation state is simultaneously oxidized and reduced. In a disproportionation reaction, one reactant should always contain an element that can have at least three oxidation states.
2H2O2(aq) —> 2H2O (l) + O2(g)
Here the oxidation number in the reactant is (-1), it increases to zero in O2 and decreases to (-2) in H2O. Oxidation number in Hydrogen does not change in the reaction.
HOW TO IDENTIFY A REDOX REACTION – Summary
Redox reactions are considered as electron transfer reaction. In a redox reaction, one element is oxidizing and it releases electrons and one element is reducing by gaining the released electrons. The extent on oxidation is equals to the extent of reduction in terms of electrons exchanging in the reaction. There are two half reactions in a redox reaction; they are called oxidation half reaction and the reduction half reaction. There is an increase in oxidation number in oxidation, similarly the oxidation number decreases in the reduction.
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