For each of the following balanced oxidation–reduction reactions, (i) identify the oxidation numbers for all the elements in the reactants and products and (ii) state the total number of electrons transferred in each reaction. (a) 2 MnO₄^-(aq) + 3 S^2-(aq + 4 H₂O(l) → 3 S(s) + 2 MnO₂(s) + 8 OH^-(aq)

Verified step by step guidance

1. Step 1: Assign oxidation numbers to each element in the reactants and products. For Mn in MnO_4^-, the oxidation number is +7. For S^2-, the oxidation number is -2. For H in H_2O, the oxidation number is +1, and for O, it is -2. In the products, S in S(s) is 0, Mn in MnO_2 is +4, and OH^- has O as -2 and H as +1.
2. Step 2: Identify the changes in oxidation numbers to determine which elements are oxidized and which are reduced. Mn goes from +7 in MnO_4^- to +4 in MnO_2, indicating a reduction. S goes from -2 in S^2- to 0 in S(s), indicating an oxidation.
3. Step 3: Calculate the change in oxidation numbers for each element involved in the redox process. Mn changes from +7 to +4, a change of 3 units per Mn atom. S changes from -2 to 0, a change of 2 units per S atom.
4. Step 4: Determine the total number of electrons transferred. Since there are 2 Mn atoms and 3 S atoms involved, calculate the total electron transfer by multiplying the change in oxidation number by the number of atoms: (2 Mn atoms * 3 electrons) + (3 S atoms * 2 electrons).
5. Step 5: Sum the total electrons transferred from both the oxidation and reduction processes to find the total number of electrons transferred in the reaction.

Key Concepts

Oxidation Numbers

Oxidation numbers are assigned to atoms in a compound to indicate their degree of oxidation or reduction. They help in identifying which atoms are oxidized and which are reduced during a redox reaction. The rules for assigning oxidation numbers include that the oxidation number of an element in its elemental form is zero, and for monoatomic ions, it equals the charge of the ion.

Redox Reactions

Redox (reduction-oxidation) reactions involve the transfer of electrons between two species. In these reactions, one species is oxidized (loses electrons) while another is reduced (gains electrons). Understanding the electron transfer is crucial for determining the total number of electrons transferred, which is essential for balancing the reaction and understanding its stoichiometry.

Balancing Redox Reactions

Balancing redox reactions requires ensuring that both mass and charge are conserved. This often involves separating the reaction into half-reactions for oxidation and reduction, balancing each half for mass and charge, and then combining them. The total number of electrons transferred can be determined from the balanced half-reactions, providing insight into the stoichiometry of the overall reaction.