Question

Which species is acting as a reducing agent in the following reaction? \(14 \mathrm{H}^{*}+\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}+3 \mathrm{Ni} \rightarrow 2 \mathrm{Cr}^{3+}+7 \mathrm{H}_{2} \mathrm{O}+3 \mathrm{Ni}^{2 *}\) (a) \(\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}\) (b) Ni (c) \(\mathrm{H}^{*}\) (d) \(\mathrm{H}_{2} \mathrm{O}\)

Short answer

Nickel (Ni) is acting as the reducing agent.

Step-by-step solution

Identify the Oxidation States of Elements in the Reactants and Products

Examine each element in the reaction to determine its oxidation state before and after the reaction. This can be done using the known oxidation states of elements and ions, and by applying the rules for assigning oxidation numbers.

Determine the Change in Oxidation States

Compare the oxidation states of elements in the reactants with their corresponding states in the products to find out which elements have undergone a change in oxidation state, which indicates either oxidation or reduction.

Identify the Reducing Agent

The species that is oxidized (loses electrons) is the reducing agent. It's the one that provides electrons to another substance and therefore its oxidation state increases.

Key concepts

Understanding Oxidation States

Learning to identify and understand oxidation states is crucial in chemistry, especially when studying redox reactions. Oxidation states, often called oxidation numbers, provide us with a way to keep track of electrons during chemical reactions.

Oxidation states are assigned to each atom based on a set of rules. The most important of which are: pure elements have an oxidation state of zero, the sum of oxidation states in a neutral compound is zero, and in ions, the sum of oxidation states equals the charge of the ion. These simple rules are the foundation for mastering redox chemistry.

For example, in the molecule water (H₂O), oxygen has an oxidation state of -2, while hydrogen has an oxidation state of +1. Knowing these, we can predict how these atoms might change during chemical reactions. Oxidation states play a pivotal role in identifying the reducing and oxidizing agents in a reaction.

Redox Reactions Demystified

The term 'redox' is a portmanteau for reduction-oxidation reactions. These reactions are all about the transfer of electrons between substances. An increase in the oxidation state indicates oxidation (loss of electrons), while a decrease signifies reduction (gain of electrons).

In any redox reaction, both processes occur simultaneously: the substance that gets oxidized loses electrons and is called the reducing agent, since it causes reduction in the other substance. Conversely, the substance that gets reduced gains electrons and is called the oxidizing agent.

The balance in electron exchange is key. For instance, in the textbook problem, when we analyze the reaction carefully, we'll notice that electrons move from one species to another, shifting oxidation numbers and identifying which species acts as a reducing agent.

Assigning Oxidation Numbers

To successfully engage with redox reactions, assigning oxidation numbers correctly is essential. It helps us track how electrons move during the reaction.

Let's apply this to practice. Take the provided textbook exercise where the reaction takes place between hydrogen ions (H⁺), dichromate ions (Cr₂O₇^2-), and nickel (Ni). By using the rules for assigning oxidation numbers, we can determine the oxidation states, before the reaction, as +1 for hydrogen, +6 for chromium (within the dichromate ion), and 0 for nickel (since it's in its elemental form).

After the reaction, the products have different oxidation states, with chromium becoming +3 in Cr³⁺ and nickel becoming +2 in Ni²⁺. With these observations, students can identify which element's oxidation state has increased - indicating it has been oxidized—and thus determine the reducing agent in the reaction.