Question

Explain why \(\mathrm{SO}_{2}\) can be used as a reducing agent but \(\mathrm{SO}_{3}\) cannot.

Short answer

In summary, SO₂ can act as a reducing agent because it has the potential to be further oxidized, as its sulfur atom has an oxidation state of +4, which can increase to +6, forming SO₃. However, SO₃ cannot act as a reducing agent because its sulfur atom is already at its maximum oxidation state of +6, making it unable to donate electrons to other species.

Step-by-step solution

Understanding Reducing Agents

A reducing agent is a species that donates electrons to another species in a redox reaction, thereby getting oxidized itself. The reducing agent typically has a low oxidation state, with the potential for further oxidation.

Identifying Oxidation States

In SO₂, the sulfur atom has an oxidation state of +4 as each of the two oxygen atoms has an oxidation state of -2. In contrast, the sulfur atom in SO₃ has an oxidation state of +6 as each of the three oxygen atoms has an oxidation state of -2.

Evaluating the Redox Potential

Examining the oxidation states, it is clear that SO₂ has the potential to be further oxidized since sulfur is only at a +4 oxidation state. The maximum oxidation state sulfur can reach is +6, which is the same oxidation state sulfur has in SO₃. This means that if SO₂ undergoes further oxidation, it can form SO₃, as seen in the balanced equation below: \[2\,\text{SO}_2 + \text{O}_2 \rightarrow 2\,\text{SO}_3\]

Comparing Redox Stability

For SO₂, since it can be further oxidized, it has the ability to donate electrons to other species, hence, acting as a reducing agent. In contrast, SO₃ cannot act as a reducing agent because the sulfur atom has already reached its maximum oxidation state (+6), and it does not have the ability to donate electrons to other species.

Conclusion

Based on the oxidation state comparison and the redox potential of SO₂ and SO₃, we can conclude that SO₂ can act as a reducing agent, as it can donate electrons and be further oxidized to SO₃. However, SO₃ cannot act as a reducing agent because its sulfur atom is already at its maximum oxidation state, and it does not have the ability to donate electrons to other species.

Key concepts

Understanding Oxidation States

Getting to grips with oxidation states is crucial for anyone tackling redox chemistry. Oxidation states indicate the degree of oxidation of an atom within a compound. Think of it like a team jersey number that helps pinpoint each atom's role in a redox process.

An atom's oxidation state can often clue you into its potential to act in a reaction. Generally speaking, an atom with a lower oxidation state has more electrons that it 'could' lose, making it a prime candidate to be a reducing agent—a chemical Samaritan that gives electrons away to another atom in need. In our exercise, sulfur dioxide (SO₂) has a sulfur oxidation state of +4, which means it's halfway up the ladder—it can climb up to a max of +6 or step down. On the other side, sulfur trioxide (SO₃) has a sulfur oxidation state of +6, basically shouting 'I'm all out of electrons to give!'

Redox Reactions Demystified

Ever seen a party where some guests take the snacks, and others bring them? That’s kind of how redox reactions work. These reactions involve the transfer of electrons between substances. Reducing agents are the generous ones offering some of their electrons, while oxidizing agents are those taking the electrons.

Let's make it super simple: reducing agents reduce the charge of other atoms by giving out electrons, while they themselves get oxidized, or lose electrons. Caught in a redox reaction, sulfur dioxide (SO₂) can hand over electrons to become sulfur trioxide (SO₃) as shown in the chemical equation from the exercise. SO₂ is like the person at the party who brings an extra pizza for everyone—it's a hit because it's so giving!

Sulfur Dioxide as a Reducing Agent

Considering sulfur dioxide as a reducing agent isn't just about chemistry; it's like understanding why a half-charged battery still has juice to power a device. SO₂ is electron-rich — its +4 oxidation state of sulfur means it's prepared to give up electrons and climb to a +6 state, transforming into SO₃.

Why can't SO₃ do the same? Think of SO₃ as the fully charged battery; it has nowhere to go — no more electrons to give. SO₂ is ready to be the superhero, swinging into the reaction to rescue other atoms by donating electrons. Meanwhile, SO₃ is the retired hero, having already given its all. Remember, a great reducing agent likes to live life on the edge, ready to lose some electrons and jump to a higher oxidation state, just like SO₂ is set to do.