Which of the following oxides can act as oxidizing as well as reducing agent? (a) \(\mathrm{SO}_{2}\) (b) \(\mathrm{SO}_{2}\) (c) Both \(\mathrm{SO}_{2}\), and \(\mathrm{SO}_{3}\) (d) Neither \(\mathrm{SO}_{2}\) nor \(\mathrm{SO}_{3}\)

When we discuss chemical reactions, we refer to the process in which substances, known as reactants, transform into entirely new substances—products.
During these reactions, the atoms in reactants rearrange and form new chemical bonds. This process can involve the transfer of electrons between species, and when it does, we call these reactions redox reactions.

Redox reactions are a central concept in chemistry, short for reduction-oxidation reactions. Here's a quick way to remember what happens in these reactions:

• Oxidation is the loss of electrons.
• Reduction is the gain of electrons.

It's vital to recognize that oxidation and reduction always occur together; when one species is oxidized, another is reduced. This is because electrons are not destroyed or created in chemical reactions—they are simply transferred from one species to another.

Understanding oxidation states (also called oxidation numbers) is essential in analyzing redox reactions. Oxidation states are hypothetical charges that an atom would have if all bonds to it were ionic. They are a tool that helps chemists keep track of electron transfers.

Some key points to keep in mind are:

• The oxidation state of an element in its pure form is always zero.
• For ions composed of a single atom, the oxidation state is equal to the ion's charge.
• Oxidation states can help predict whether a substance will act as an oxidizing agent or a reducing agent.

An increase in oxidation state means oxidation, while a decrease signifies reduction. The substance that gets reduced (gain of electrons) is the oxidizing agent and the one that gets oxidized (loss of electrons) is the reducing agent.

Now, let's focus on sulfur dioxide, a chemical compound with the formula \(\mathrm{SO}_2\). Sulfur dioxide can act as both an oxidizing and reducing agent depending on the context of the reaction.

As we discussed earlier in the solution to the exercise, \(\mathrm{SO}_2\) has sulfur in a +4 oxidation state. This intermediate state allows sulfur dioxide to either gain or lose oxygen, leading to an increase or decrease in oxidation state, respectively.

Reducing Properties of SO2

When it acts as a reducing agent, \(\mathrm{SO}_2\) gets oxidized and can increase the oxidation state of sulfur to +6, as seen in \(\mathrm{SO}_3\). In this role, it donates electrons to another substance.

Oxidizing Properties of SO2

Contrastingly, as an oxidizing agent, \(\mathrm{SO}_2\) can accept electrons when another substance undergoes oxidation and decreases the sulfur's oxidation state from +4 to a lower number, such as +2 or 0.

This versatility makes \(\mathrm{SO}_2\) an interesting and important chemical, particularly in industrial processes where it can play a dual role in redox chemistry.