In Practice Exercise 5.6 (Section 5.1), some of the uses of chlorine dioxide were described along with a reaction that could be used to make \(\mathrm{ClO}_{2}\). Another reaction that is used to make this substance is \(\mathrm{HCl}+\mathrm{NaOCl}+2 \mathrm{NaClO}_{2} \longrightarrow 2 \mathrm{ClO}_{2}+2 \mathrm{NaCl}+\mathrm{NaOH}\) Which element is oxidized? Which element is reduced? Which substance is the oxidizing agent and which is the reducing agent?

Understanding oxidation states is crucial when studying redox reactions. Oxidation states, also sometimes referred to as oxidation numbers, give us a hypothetical charge that an atom would have if all bonds were purely ionic. When determining the oxidation state of chlorine in different compounds, we follow certain rules, such as the sum of oxidation states in a neutral compound is zero, and for ions, it's equal to the charge of the ion.

For instance, in a molecule of HCl, hydrogen is assigned an oxidation state of +1, following the rule that hydrogen generally has an oxidation state of +1 when bonded to nonmetals. This leaves chlorine with an oxidation state of -1 to balance the molecule. In compounds like NaOCl and NaClO2, we use the fact that oxygen has a typical oxidation state of -2, and the sum of the oxidation states needs to balance with the overall charge of the compound. Through this example, it's clear that identifying the change in oxidation states helps us identify what is being oxidized and reduced.

Oxidizing and reducing agents are substances that help to drive redox reactions. An oxidizing agent, due to its tendency to gain electrons, ends up being reduced in the reaction. Conversely, a reducing agent donates electrons and is oxidized as the reaction progresses. In our example, NaOCl acts as the oxidizing agent as it contains chlorine that goes from an oxidation state of +1 to -1. Meanwhile, NaClO2 serves as the reducing agent because it contains chlorine that increases in oxidation state from +3 to +4, indicating a loss of electrons. Understanding the role these agents play helps students to trace the flow of electrons and comprehend redox reactions more thoroughly.

Chemical reactions are processes where reactants transform into products. In redox reactions, like the one used to produce chlorine dioxide, it's a matter of electron transfer between reactants. The oxidation states change as a result of this electron transfer. Analyzing the changes in oxidation states for chlorine across various compounds is key to understanding the reaction. Through careful examination, we can confirm that our chemical reaction adheres to the law of conservation of mass, meaning all atoms in the reactants are accounted for in the products, with different bonding patterns reflective of electron shifts.

Balancing chemical equations is an essential skill in chemistry, as it ensures that the Law of Conservation of Mass is respected. Each atom present in the reactants must also be present in the products. To balance the chemical equation provided, we count the number of atoms for each element on both sides and increment the coefficients of the compounds until we achieve a balanced equation. Balancing redox reactions can be complex, as it involves not only balancing the atoms but also ensuring that the transfer of electrons is accounted for. In the exercise, the equation provided is already balanced, which is why we're able to go directly into identifying oxidation and reduction processes without additional steps to balance it.