Before delving into the crux of this question, let’s first understand what redox reactions are. Redox, short for reduction-oxidation, refers to chemical reactions in which there is a transfer of electrons from one species to another.
Consider the following generic redox reaction:
Oxidant + Reductant → Reductant + Oxidant
In the reaction above, the oxidant is the species that is reduced, while the reductant is the species that is oxidized. In other words, the oxidant gains electrons, while the reductant loses electrons.
Now, coming back to the primary question, it is a well-established fact that “In all redox reactions, the total increase in oxidation number must be equal to the total decrease in oxidation number.”
To understand why this statement holds true, let’s first define what oxidation number is. Oxidation number is the formal charge that an atom in a molecule or ion would have if all the bonds were 100% ionic.
For example, in H2O, the oxidation number of hydrogen is +1, while that of oxygen is -2. Similarly, in NaCl, the oxidation number of sodium is +1, while that of chlorine is -1.
The Conservation of Charge
Now, let’s try to answer the original question. Why must the total increase in oxidation number equal the total decrease in oxidation number in redox reactions?
The answer to this question lies in the conservation of charge. According to this fundamental principle, charges can neither be created nor destroyed. The total charge of a system before a reaction is equal to the total charge of the same system after the reaction.
Let’s use the example of a redox reaction between iron (II) ions and permanganate ions to illustrate this:
In the reaction above, iron (II) ions are oxidized to iron (III) ions, while permanganate ions are reduced to manganese (II) ions.
The oxidation states of iron and manganese in the reactants side are +2 and +7, respectively. The oxidation states of these elements in the product side are +3 and +2, respectively.
The increase in oxidation state is 1, and the decrease is 5. However, when we add up the increase and decrease in oxidation states, we find that they are equal:
5 × 2 iron atoms = 10
1 × 2 manganese atoms = 2
Overall, there is a decrease of 10 units of charge, which must be matched by an increase of 10 units of charge. This shows how the conservation of charge plays out in redox reactions.
Applying The Same Principle to Other Examples
It is important to note that the principle of conservation of charge holds true for all redox reactions, not just the one we used as an example above.
If we were to take any two species that undergo a redox reaction, we would find that the total increase in oxidation number of one species is equal to the total decrease in oxidation number of the other species.
For example, consider the reaction between copper (II) ions and zinc. Copper (II) ions are reduced to copper metal, while zinc is oxidized to zinc ions.
The oxidation state of copper in copper (II) ions is +2, while the oxidation state of copper in copper metal is 0. The increase in oxidation state is 2.
The oxidation state of zinc in zinc is 0, while that in zinc ions is +2. The decrease in oxidation state is also 2.
Thus, we can see that the increase and decrease in oxidation states of these two elements is equal, and is determined by the conservation of charge.
Conclusion
In conclusion, we can say that the statement “In all redox reactions, the total increase in oxidation number must be equal to the total decrease in oxidation number” is based on the fundamental principle of the conservation of charge. This principle holds true for all redox reactions, and is a result of the transfer of electrons from one species to another.
Why Must the Total Increase In Oxidation Numbers Equal the Total Decrease of Them In Redox Reactions?
Before delving into the crux of this question, let’s first understand what redox reactions are. Redox, short for reduction-oxidation, refers to chemical reactions in which there is a transfer of electrons from one species to another.
Consider the following generic redox reaction:
Oxidant + Reductant → Reductant + Oxidant
In the reaction above, the oxidant is the species that is reduced, while the reductant is the species that is oxidized. In other words, the oxidant gains electrons, while the reductant loses electrons.
Now, coming back to the primary question, it is a well-established fact that “In all redox reactions, the total increase in oxidation number must be equal to the total decrease in oxidation number.”
To understand why this statement holds true, let’s first define what oxidation number is. Oxidation number is the formal charge that an atom in a molecule or ion would have if all the bonds were 100% ionic.
For example, in H2O, the oxidation number of hydrogen is +1, while that of oxygen is -2. Similarly, in NaCl, the oxidation number of sodium is +1, while that of chlorine is -1.
The Conservation of Charge
Now, let’s try to answer the original question. Why must the total increase in oxidation number equal the total decrease in oxidation number in redox reactions?
The answer to this question lies in the conservation of charge. According to this fundamental principle, charges can neither be created nor destroyed. The total charge of a system before a reaction is equal to the total charge of the same system after the reaction.
Let’s use the example of a redox reaction between iron (II) ions and permanganate ions to illustrate this:
In the reaction above, iron (II) ions are oxidized to iron (III) ions, while permanganate ions are reduced to manganese (II) ions.
The oxidation states of iron and manganese in the reactants side are +2 and +7, respectively. The oxidation states of these elements in the product side are +3 and +2, respectively.
The increase in oxidation state is 1, and the decrease is 5. However, when we add up the increase and decrease in oxidation states, we find that they are equal:
5 × 2 iron atoms = 10
1 × 2 manganese atoms = 2
Overall, there is a decrease of 10 units of charge, which must be matched by an increase of 10 units of charge. This shows how the conservation of charge plays out in redox reactions.
Applying The Same Principle to Other Examples
It is important to note that the principle of conservation of charge holds true for all redox reactions, not just the one we used as an example above.
If we were to take any two species that undergo a redox reaction, we would find that the total increase in oxidation number of one species is equal to the total decrease in oxidation number of the other species.
For example, consider the reaction between copper (II) ions and zinc. Copper (II) ions are reduced to copper metal, while zinc is oxidized to zinc ions.
The oxidation state of copper in copper (II) ions is +2, while the oxidation state of copper in copper metal is 0. The increase in oxidation state is 2.
The oxidation state of zinc in zinc is 0, while that in zinc ions is +2. The decrease in oxidation state is also 2.
Thus, we can see that the increase and decrease in oxidation states of these two elements is equal, and is determined by the conservation of charge.
Conclusion
In conclusion, we can say that the statement “In all redox reactions, the total increase in oxidation number must be equal to the total decrease in oxidation number” is based on the fundamental principle of the conservation of charge. This principle holds true for all redox reactions, and is a result of the transfer of electrons from one species to another.