Oxidation and reduction Knowledge-Based Study

Study notes

1. Understanding Oxidation and Reduction

   Historically, chemists defined oxidation as the addition of oxygen to a substance and reduction as the removal of oxygen. In modern chemistry, however, these terms are defined more broadly in terms of electron transfer.

   • Oxidation: The process in which an atom, ion, or molecule loses one or more electrons. This leads to an increase in oxidation state.
   • Reduction: The process in which an atom, ion, or molecule gains one or more electrons. This leads to a decrease (a "reduction") in oxidation state.

   Because electrons cannot exist freely in solution, oxidation and reduction always occur simultaneously in complementary processes known as redox reactions. A widely used mnemonic to recall electron movement is OIL RIG: Oxidation Is Loss, Reduction Is Gain of electrons.

   Tracking Electrons: Oxidation Numbers

   To determine whether an element has undergone oxidation or reduction, chemists assign an oxidation number (or oxidation state) to each atom in a chemical species. The oxidation number represents the hypothetical charge an atom would have if all shared electrons were assigned to the more electronegative element.

   Key Rules for Assigning Oxidation Numbers

   1. Free Elements: Any atom in its pure uncombined elemental form has an oxidation state of 0 (e.g., $\text{Na}$, $\text{Fe}$, $\text{O}_2$, $\text{Cl}_2$).
   2. Monatomic Ions: The oxidation number equals the actual electrical charge on the ion (e.g., $\text{Na}^+$ is +1, $\text{S}^{2-}$ is -2).
   3. Hydrogen: Usually +1 when bonded to nonmetals (such as in $\text{H}_2\text{O}$ and $\text{HCl}$), but -1 when bonded to metals in metal hydrides (such as $\text{NaH}$).
   4. Oxygen: Typically -2 in most compounds, with exceptions including peroxides (such as $\text{H}_2\text{O}_2$, where it is -1) and when bonded to fluorine (such as $\text{OF}_2$, where it is +2).
   5. Fluorine and Halogens: Fluorine is always -1 in its compounds. Other halogens are usually -1 unless bonded to oxygen or a more electronegative halogen.
   6. Sum of Oxidation Numbers: For a neutral compound, the sum of all oxidation numbers must equal 0. For a polyatomic ion, the sum must equal the net charge of the ion.

   Identifying Agents in a Reaction

   In any redox transformation, the participating species facilitate each other's electron transfers. They are classified into two complementary roles:

   • Oxidizing Agent (Oxidant): The reactant that brings about the oxidation of another substance. In doing so, the oxidizing agent gains electrons and undergoes reduction. Its oxidation number decreases.
   • Reducing Agent (Reductant): The reactant that brings about the reduction of another substance. In doing so, the reducing agent loses electrons and undergoes oxidation. Its oxidation number increases.

   Important distinction: The terms oxidized and reduced describe what happens to individual elements, whereas oxidizing agent and reducing agent typically refer to the entire reactant formula or species containing that element.

   Analyzing Redox via Half-Reactions

   A redox equation can be split into two individual components called half-reactions, explicitly showing electron movement.

   Consider the single-displacement reaction between solid zinc metal and aqueous copper(II) sulfate:

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

   Step-by-Step Breakdown

   • Oxidation Half-Reaction: Zinc metal loses two electrons:
     $$\text{Zn}(s) \rightarrow \text{Zn}^{2+}(aq) + 2e^-$$
     The oxidation state of zinc increases from 0 to +2. Therefore, $\text{Zn}$ is oxidized and serves as the reducing agent.
   • Reduction Half-Reaction: Copper(II) ions gain two electrons:
     $$\text{Cu}^{2+}(aq) + 2e^- \rightarrow \text{Cu}(s)$$
     The oxidation state of copper decreases from +2 to 0. Therefore, $\text{Cu}^{2+}$ is reduced and serves as the oxidizing agent.

   Notice that the number of electrons released in the oxidation process exactly balances the number of electrons accepted in the reduction process, preserving total electrical charge.