Oxidizing and reducing agents Knowledge-Based Study

Study notes

1. Oxidation and Reduction

   Oxidation-reduction (redox) reactions are chemical processes that involve the net transfer of electrons from one chemical species to another. To fully understand these processes, two complementary half-reactions must be defined:

   • Oxidation: The loss of electrons by a chemical species, resulting in an increase in its oxidation state.
   • Reduction: The gain of electrons by a chemical species, resulting in a decrease in its oxidation state.

   Because electrons cannot exist freely in an isolated chemical system, oxidation and reduction always take place simultaneously. The total quantity of electrons released during oxidation must equal the quantity of electrons accepted during reduction.

   The Roles of Chemical Agents

   In every redox transformation, the participating substances act as functional agents based on their complementary roles in the electron transfer process:

   • Oxidizing Agent (Oxidant): The substance that gains electrons and undergoes reduction. By removing electrons from another reactant, it causes that reactant to be oxidized. Consequently, the oxidation state of the key atom within the oxidizing agent decreases.
   • Reducing Agent (Reductant): The substance that loses electrons and undergoes oxidation. By providing electrons to another reactant, it causes that reactant to be reduced. Consequently, the oxidation state of the key atom within the reducing agent increases.

   Important Convention: The terms oxidizing agent and reducing agent always refer to the complete chemical formula of the reactant species in the balanced equation, rather than just the specific element undergoing an oxidation state change.

   Systematic Analysis Method

   To correctly identify which reactant acts as an oxidizing agent and which acts as a reducing agent, follow these sequential steps:

   1. Assign oxidation numbers to every atom on both the reactant and product sides of the equation.
   2. Compare the initial and final oxidation states to identify which elements underwent an increase or decrease.
   3. Identify the oxidized and reduced substances: the species whose element increased in oxidation state was oxidized; the species whose element decreased in oxidation state was reduced.
   4. Assign the complementary agent roles: the oxidized species is the reducing agent, and the reduced species is the oxidizing agent.

   Illustrative Example

   Consider the single displacement reaction:

   $\text{Zn}(s) + \text{Cu}^{2+}(aq) \rightarrow \text{Zn}^{2+}(aq) + \text{Cu}(s)$

   • Zinc starts at an oxidation state of 0 and increases to +2. Because zinc loses two electrons, $\text{Zn}(s)$ is oxidized and serves as the reducing agent.
   • Copper starts at an oxidation state of +2 and decreases to 0. Because copper(II) gains two electrons, $\text{Cu}^{2+}(aq)$ is reduced and serves as the oxidizing agent.

   Polyatomic Ions and Common Laboratory Agents

   In reactions containing polyatomic ions, determine the oxidation state of the central atom before assigning agent roles. Consider the reaction between iron(II) and the permanganate ion in acidic solution:

   $\text{MnO}_4^-(aq) + 5\text{Fe}^{2+}(aq) + 8\text{H}^+(aq) \rightarrow \text{Mn}^{2+}(aq) + 5\text{Fe}^{3+}(aq) + 4\text{H}_2\text{O}(l)$

   • Manganese changes from +7 in $\text{MnO}_4^-$ to +2 in $\text{Mn}^{2+}$. Because it gains electrons, the permanganate ion ($\text{MnO}_4^-$) is reduced and functions as the oxidizing agent.
   • Iron changes from +2 in $\text{Fe}^{2+}$ to +3 in $\text{Fe}^{3+}$. Because it loses an electron, $\text{Fe}^{2+}$ is oxidized and functions as the reducing agent.
   • Hydrogen and oxygen maintain oxidation states of +1 and -2, respectively, functioning as spectator entities with respect to electron transfer.

   Key Summary Rules

   • Oxidized species = Undergoes electron loss = Oxidation number increases = Reducing agent
   • Reduced species = Undergoes electron gain = Oxidation number decreases = Oxidizing agent