Animated Solution for Chemistry - Redox Reactions: The compound that cannot act both as oxidising and reducing agent is

The journey into the world of redox reactions is like exploring a grand cosmic exchange of currency, where the currency is electrons. Some elements are wealthy and willing to give, while others are greedy and want to take. But what happens when an element is caught right in the middle?

In this problem, we are asked to identify a compound that cannot act as both an oxidising and a reducing agent. To crack this, we need to deeply understand what these terms mean and how an element's oxidation state dictates its chemical behavior.

Imagine a ladder. This ladder represents the possible oxidation states an element can have. The bottom of the ladder is the minimum oxidation state, which corresponds to the maximum number of electrons an element can comfortably gain to complete its octet. The top of the ladder is the maximum oxidation state, representing the maximum number of valence electrons an element can lose.

When an element is sitting on a middle rung of this ladder—an intermediate oxidation state—it has options. It can climb up by losing more electrons (getting oxidised), which makes it a reducing agent. Or, it can climb down by gaining electrons (getting reduced), which makes it an oxidising agent.

However, if an element is at the very top of the ladder, it has no more valence electrons to lose. It cannot climb any higher. It can only go down. Therefore, an element in its maximum oxidation state can only undergo reduction, meaning it acts strictly as an oxidising agent.

Conversely, if it's at the very bottom, it can only go up, acting strictly as a reducing agent. Our mission is to find the compound where the central atom is stuck at the very top (or bottom) of its ladder.

Let's look at our first candidate, sulfurous acid, or ${\text{H}}_2{\text{SO}}_3$. We need to find the oxidation state of the central sulfur atom.

We know that hydrogen typically has an oxidation state of $+1$, and oxygen is usually $-2$. Setting up our algebraic equation for the neutral molecule:

Sulfur is in Group 16 of the periodic table. It has 6 valence electrons. This means it can lose up to 6 electrons (maximum oxidation state of $+6$) or gain 2 electrons to complete its octet (minimum oxidation state of $-2$).

Since $+4$ is right in the middle of $-2$ and $+6$, sulfur in ${\text{H}}_2{\text{SO}}_3$ is on a middle rung of the ladder. It can be oxidised to $+6$ (like in ${\text{H}}_2{\text{SO}}_4$) or reduced to lower states. Thus, it can act as both an oxidising and a reducing agent.

Next up is nitrous acid, ${\text{HNO}}_2$. Let's calculate the oxidation state of nitrogen.

Nitrogen belongs to Group 15, possessing 5 valence electrons. Its oxidation state ladder ranges from $-3$ (gaining 3 electrons, like in ammonia, ${\text{NH}}_3$) to $+5$ (losing all 5 valence electrons, like in nitric acid, ${\text{HNO}}_3$).

The calculated state of $+3$ is an intermediate value. Nitrogen here can still lose 2 more electrons or gain up to 6. Therefore, ${\text{HNO}}_2$ is perfectly capable of acting as both an oxidising and a reducing agent.

Now we encounter hydrogen peroxide, ${\text{H}}_2{\text{O}}_2$. This molecule is famous in redox chemistry.

Normally, oxygen has an oxidation state of $-2$. But peroxides are the exception. Because of the single bond between the two highly electronegative oxygen atoms (the peroxide linkage), the oxidation state of oxygen here is $-1$.

Oxygen's typical range is from $-2$ (in most oxides) to $0$ (in elemental ${\text{O}}_2$). The state of $-1$ is exactly halfway between.

Hydrogen peroxide can easily be reduced to water (${\text{H}}_2\text{O}$), where oxygen drops to $-2$. It can also be oxidised to oxygen gas (${\text{O}}_2$), where it rises to $0$. Because it can swing both ways, ${\text{H}}_2{\text{O}}_2$ is a versatile chemical that acts as both an oxidising and a reducing agent.

Phosphoric Acid

Finally, we arrive at phosphoric acid, ${\text{H}}_3{\text{PO}}_4$. Let's determine the oxidation state of phosphorus.

Phosphorus is in Group 15, just like nitrogen. It has exactly 5 valence electrons. This means its maximum possible oxidation state is $+5$.

In ${\text{H}}_3{\text{PO}}_4$, phosphorus has already lost all 5 of its valence electrons to the highly electronegative oxygen atoms. It is standing at the absolute top of its oxidation ladder. It has absolutely no more electrons to give away.

Because it cannot be oxidised any further, it cannot act as a reducing agent. It can only accept electrons and climb down the ladder (undergo reduction). Therefore, ${\text{H}}_3{\text{PO}}_4$ can only act as an oxidising agent.

This perfectly answers our question. Phosphoric acid is the compound that cannot act as both!