Animated Solution for Chemistry - Redox Reactions: The redox reaction among the following is

Have you ever wondered what happens at the atomic level when substances react? Sometimes, they just swap partners, like dancers changing pairs. But other times, there is a fundamental exchange of currency—the currency of electrons. This exchange is what we call a redox reaction.

In a redox reaction, one species gives up electrons (undergoes oxidation), while another species eagerly accepts them (undergoes reduction). The easiest way to track this flow of electrons is by looking at the oxidation states of the elements before and after the reaction. If the oxidation states change, we have a redox reaction on our hands!

Let's put on our detective hats and investigate the four chemical reactions given in the problem to find the true redox reaction.

Acid-Base Neutralization

Our first candidate is the reaction between sulfuric acid and sodium hydroxide:

This is a classic neutralization reaction. Let's check the oxidation states. Sulfur in ${\text{H}}_2{\text{SO}}_4$ is at $+6$, and it remains $+6$ in the sulfate ion of ${\text{Na}}_2{\text{SO}}_4$. Sodium is $+1$ in $\text{NaOH}$ and stays $+1$ in the product. Oxygen and hydrogen remain steadfast at $-2$ and $+1$, respectively. Since no element changes its oxidation state, no electrons were transferred. This is not a redox reaction.

Precipitation Reaction

Next up is the reaction of a cobalt complex with silver nitrate:

This looks intimidating, but it's just a double displacement reaction resulting in the precipitation of silver chloride. The cobalt ion is comfortably sitting at an oxidation state of $+3$ (since water is neutral and there are three chloride ions). It remains $+3$ in the product. Silver is $+1$ in ${\text{AgNO}}_3$ and stays $+1$ in $\text{AgCl}$. Once again, the oxidation states are completely unchanged. Not a redox reaction!

Allotropic Transformation

Let's skip to the fourth option for a moment—the formation of ozone from atmospheric oxygen:

Here, oxygen gas (${\text{O}}_2$) is absorbing ultraviolet light to form ozone (${\text{O}}_3$). Both of these are elemental forms (allotropes) of oxygen. By definition, the oxidation state of an element in its free, uncombined state is exactly zero. So, oxygen goes from $0$ to $0$. No change in oxidation state means this is definitely not a redox reaction.

The Combination Reaction

Finally, let's examine the combination of dinitrogen and dioxygen at a scorching $2000\text{ K}$:

Let's break down the oxidation states. On the reactant side, both nitrogen and oxygen are in their elemental diatomic forms. Therefore, their initial oxidation states are $0$.

Now, look at the product, nitric oxide ($\text{NO}$). Oxygen is more electronegative than nitrogen, so it dictates the terms and takes an oxidation state of $-2$. For the molecule to be electrically neutral, nitrogen must balance this out by taking an oxidation state of $+2$.

Let's summarize the changes:

Nitrogen: Goes from $0$ to $+2$. Its oxidation state increased, meaning it lost electrons. This is oxidation. Oxygen: Goes from $0$ to $-2$. Its oxidation state decreased, meaning it gained electrons. This is reduction.

Because both oxidation and reduction are occurring simultaneously, we have found our match! This combination reaction is a textbook example of a redox reaction.

Tracking oxidation states is like following the money trail in a financial investigation. It tells you exactly where the electrons are going. By systematically checking the oxidation states of all elements in a reaction, you can confidently identify whether a redox process is taking place. Keep practicing this skill, and you'll be a redox detective in no time!