Question

Challenge Identify the oxidizing agent and the reducing agent in each reaction. a. \(M g+I_{2} \rightarrow M g l_{2}\) b. \(H_{2} S+C l_{2} \rightarrow S+2 H C l\)

Short answer

For reaction a (\(Mg + I_2 \rightarrow MgI_2\)), the reducing agent is \(Mg\) and the oxidizing agent is \(I_2\). For reaction b (\(H_2S + Cl_2 \rightarrow S + 2HCl\)), the reducing agent is \(H_2S\) and the oxidizing agent is \(Cl_2\).

Step-by-step solution

Determine the oxidation states of the elements in the reactants and the products

Determine the oxidation states of the different elements in the reactants and products: \(Mg\) is in its elemental state and has an oxidation state of 0. \(I_2\) is in its elemental state and has an oxidation state of 0. In \(MgI_2\), \(Mg\) has an oxidation state of +2 and each \(I\) has an oxidation state of -1. Now let's compare the changes in oxidation state for each element.

Identify the changes in oxidation states and their roles as oxidizing or reducing agents

Analyze the changes in oxidation states for each element: \(Mg\): 0 → +2 (oxidation) \(I\): 0 → -1 (reduction) \(Mg\) loses electrons and undergoes oxidation, which means that \(Mg\) is the reducing agent in this reaction. On the other hand, \(I_2\) gains electrons and undergoes reduction, so \(I_2\) is the oxidizing agent. For reaction b: \(H_2S + Cl_2 \rightarrow S + 2HCl\)

Determine the oxidation states of the elements in the reactants and the products

Determine the oxidation states of the different elements in the reactants and products: In \(H_2S\), each \(H\) has an oxidation state of +1 and \(S\) has an oxidation state of -2. In \(Cl_2\), both \(Cl\) atoms have an oxidation state of 0. In the products, \(S\) is in its elemental state and has an oxidation state of 0. In \(HCl\), each \(H\) has an oxidation state of +1 and each \(Cl\) has an oxidation state of -1. Now let's compare the changes in oxidation state for each element.

Identify the changes in oxidation states and their roles as oxidizing or reducing agents

Analyze the changes in oxidation states for each element: \(S\): -2 → 0 (oxidation) \(Cl\): 0 → -1 (reduction) \(S\) loses electrons and undergoes oxidation, which means that \(H_2S\) is the reducing agent in this reaction. On the other hand, \(Cl_2\) gains electrons and undergoes reduction, so \(Cl_2\) is the oxidizing agent.

Key concepts

Oxidation States

Understanding oxidation states is fundamental in identifying the roles substances play in redox reactions. An oxidation state, often referred to as the oxidation number, is a measure of the degree of oxidation of an atom in a chemical compound. It represents the number of electrons that an atom either gains or loses when it forms a chemical bond.

In essence, it’s like the atom’s 'charge' if all the bonds to it were completely ionic. For example, in a molecule like water (H₂O), oxygen has an oxidation state of -2, while hydrogen has an oxidation state of +1. These numbers help us track electrons during a reaction which is crucial for understanding redox processes.

Oxidation State Changes

When a substance moves from a neutral element to a part of a compound, as we saw in the exercises with magnesium (Mg) and iodine (I₂), their oxidation states change. Magnesium starts at 0 and goes to +2, losing electrons, and iodine goes from 0 to -1, each atom gaining one electron. This gives us a clear picture of the reaction’s flow of electrons.

Oxidizing and Reducing Agents

In the dance of redox reactions, oxidizing and reducing agents are the key partners. An oxidizing agent takes electrons from another substance, therefore it is 'reduced'. Conversely, a reducing agent donates electrons to another substance, and thus it is 'oxidized'. These agents are always found in pairs; one cannot exist without the other in a redox reaction.

Role Identification

To spot these agents, look at the changes in oxidation states. For instance, with reaction (a), magnesium (Mg) loses two electrons and increases its oxidation state from 0 to +2 which makes it the reducing agent. Iodine (I₂), on the other hand, accepts those electrons, decreasing its oxidation number from 0 to -1, becoming the oxidizing agent. In our second example, sulfur (S) is part of hydrogen sulfide (H₂S) and it loses electrons, thus H₂S acts as the reducing agent, while chlorine (Cl₂) gains those electrons, marking it as the oxidizing agent.

Chemical Reactions

Chemical reactions involve the transformation of substances through the breaking and forming of chemical bonds, leading to new substances. Redox reactions are a special type of chemical reaction where oxidation and reduction occur simultaneously.

Every chemical reaction follows the law of conservation of mass, where the mass of the reactants equals the mass of the products. In a redox process, it's critical to ensure that the number of atoms and the charge are balanced on both sides of the reaction equation. This allows us to accurately represent the physical transactions occurring at the atomic level.

In the textbook examples, when magnesium reacts with iodine, the product is magnesium iodide, and sulfur reacting with chlorine gas produces elemental sulfur and hydrogen chloride. These transformations showcase the principles of redox reactions, with the transfer of electrons leading to significant chemical change.