Question

For each of the following reactions, indicate which elements are oxidized and which are reduced: a. \(2 \mathrm{Cr}^{3+}+3 \mathrm{Mg} \longrightarrow 2 \mathrm{Cr}+3 \mathrm{Mg}^{2+}\) b. \(\mathrm{Ni}+\mathrm{Cl}_{2} \longrightarrow \mathrm{Ni}^{2+}+2 \mathrm{Cl}^{-}\) c. \(2 \mathrm{Cl}^{-}+\mathrm{F}_{2} \longrightarrow 2 \mathrm{~F}^{-}+\mathrm{Cl}_{2}\)

Short answer

For reaction (a), Cr is reduced and Mg is oxidized. For reaction (b), Ni is oxidized and Cl is reduced. For reaction (c), Cl is oxidized and F is reduced.

Step-by-step solution

Identify the oxidation states before and after the reaction for (a)

Look at the oxidation states of chromium and magnesium before and after the reaction. Chromium goes from an oxidation state of +3 in Cr^{3+} to 0 in Cr, while magnesium goes from 0 in Mg to +2 in Mg^{2+}.

Determine which elements are oxidized and reduced in (a)

An element is oxidized if it loses electrons (increases its oxidation state), and it is reduced if it gains electrons (decreases its oxidation state). For reaction (a), Cr is reduced (from +3 to 0), and Mg is oxidized (from 0 to +2).

Identify the oxidation states before and after the reaction for (b)

Evaluate the charges on nickel and chlorine before and after the reaction. Nickel goes from 0 in Ni to +2 in Ni^{2+}, and chlorine goes from 0 in Cl_2 to -1 in Cl^{-}.

Determine which elements are oxidized and reduced in (b)

For reaction (b), Ni is oxidized (from 0 to +2), and Cl is reduced (from 0 to -1).

Identify the oxidation states before and after the reaction for (c)

Inspect the oxidation states of chlorine and fluorine before and after the reaction. Chlorine goes from -1 in Cl^{-} to 0 in Cl_2, while fluorine goes from 0 in F_2 to -1 in F^{-}.

Determine which elements are oxidized and reduced in (c)

For reaction (c), Cl is oxidized (from -1 to 0), and F is reduced (from 0 to -1).

Key concepts

Oxidation States

Understanding oxidation states is crucial for deciphering the changes that occur during oxidation-reduction (redox) reactions. An oxidation state, often referred to as oxidation number, is an indicator of the degree of oxidation of an element within a compound or ion. It represents the hypothetical charge an atom would have if all bonds to atoms of different elements were completely ionic.

For instance, in the reaction where chromium (Cr) and magnesium (Mg) react to form Cr and Mg²⁺, the oxidation states change. Initially, chromium has an oxidation state of +3 in Cr³⁺ and magnesium is in its elemental form with an oxidation state of 0. After the reaction, chromium's oxidation state is 0 (since it's in its elemental form) and magnesium's is +2 in Mg²⁺. This change in oxidation state indicates the transfer of electrons between atoms and thus, the occurrence of a redox reaction.

Redox Chemistry

Redox chemistry deals with the chemical processes involving electron transfer, which leads to changes in oxidation states of the elements involved. In a redox reaction, there is always one species that donates electrons (oxidized) and one that accepts electrons (reduced).

Using the provided textbook examples:

• In reaction (a), Mg is oxidized as it loses electrons (oxidation state increases from 0 to +2), and Cr is reduced as it gains electrons (oxidation state decreases from +3 to 0).
• In reaction (b), Ni is oxidized (oxidation state increases from 0 to +2), while Cl is reduced (oxidation state decreases from 0 to -1).
• Finally, in reaction (c), Cl is oxidized (oxidation state increases from -1 to 0), and F is reduced (oxidation state decreases from 0 to -1).

These examples illustrate that redox chemistry is fundamentally about the movement of electrons from one element or compound to another.

Electrochemistry

Electrochemistry is a branch of chemistry that studies the relationship between electricity and chemical changes, which often involve redox reactions. During an electrochemical process, oxidation occurs at the anode (positive electrode) and reduction at the cathode (negative electrode).

For example, an electrochemical cell can be constructed using the reactions from the exercise, where metal electrodes (such as Cr, Mg, Ni, or Cl) are connected by an electrical conductor. As the reaction progresses, one metal would get oxidized while the other gets reduced, driving the flow of electrons through the conductor, thus generating an electric current. This principle is the basis for batteries, where spontaneous redox reactions provide electrical energy.

Balancing Chemical Equations

Balancing chemical equations is essential to accurately represent the conservation of mass in a chemical reaction. It ensures that the number of atoms for each element is the same on both sides of the equation. When dealing with redox reactions, one must also ensure that the charge is balanced, which means the total charges on the reactant side equal the total charges on the product side.

For balancing redox reactions, the half-reaction method can be particularly useful. It involves writing separate equations for the oxidation and reduction processes and balancing them for mass and charge. These balanced half-reactions are then combined to give the balanced overall chemical equation. This method can simplify the balancing process for complex redox reactions and is crucial for understanding electrochemistry where each half-reaction occurs at a different electrode.