Oxidation Study Guide

📖 Core Concepts Redox reaction – chemical change where oxidation states shift; oxidation (loss of e⁻) and reduction (gain of e⁻) occur together. Oxidation – electron loss or increase in oxidation number. Reduction – electron gain or decrease in oxidation number. Redox pair – the reducing agent + the oxidizing agent that react. Redox couple – a specific species and its oxidized form (e.g., Fe²⁺/Fe³⁺). Half‑reaction – separate oxidation or reduction equation; two half‑reactions combine to give the overall redox equation. Oxidant (oxidizing agent) – accepts electrons, is reduced. Reductant (reducing agent) – donates electrons, is oxidized. Standard electrode potential, \(E^\circ\) – voltage of a half‑reaction under standard conditions (1 M, 1 atm, 25 °C) measured vs the standard hydrogen electrode (SHE, 0 V). Cell voltage (EMF) – \(E{\text{cell}} = E^\circ{\text{cathode}} - E^\circ{\text{anode}}\). Positive \(E{\text{cell}}\) → spontaneous. Inner‑sphere electron transfer – a bridging ligand links oxidant and reductant during electron transfer. Outer‑sphere electron transfer – reactants stay separate; electrons “hop” through space. --- 📌 Must Remember Oxidation = loss of e⁻; reduction = gain of e⁻. By convention \(E^\circ{\text{H}^+ + e^- \rightarrow \tfrac12 \text{H}2}=0\) V. Positive \(E^\circ\) → stronger oxidizing agent than H⁺ (e.g., F₂ + 2 e⁻ → 2 F⁻, +2.866 V). Negative \(E^\circ\) → weaker oxidizing agent (e.g., Zn²⁺ + 2 e⁻ → Zn, –0.763 V). Cell voltage equation: \(E{\text{cell}} = E^\circ{\text{cathode}} - E^\circ{\text{anode}}\). Electron‑transfer redox reactions are usually fast; atom‑transfer can be slow (rust) or fast (combustion). Sacrificial anode = metal with lower \(E^\circ\) that preferentially oxidizes to protect another metal (cathodic protection). NAD⁺ ↔ NADH and GSH ↔ GSSG are key biological redox couples. --- 🔄 Key Processes Write half‑reactions Balance atoms (except O/H), then balance charge with electrons. Combine half‑reactions Multiply to equalize electrons, add, cancel electrons, simplify. Calculate cell voltage Identify cathode (reduction) and anode (oxidation). Plug their \(E^\circ\) values into the cell voltage equation. Inner‑sphere ET mechanism Form a bridged complex → electron moves through bridge → complex dissociates. Outer‑sphere ET mechanism Reactants approach → electron tunnels → products separate; no bond formation. Cathodic protection setup Attach sacrificial anode → current flows from anode (oxidizing) to protected metal (cathode). Electroplating Connect metal to be plated as cathode, metal source as anode, apply voltage → metal ions reduce onto cathode surface. --- 🔍 Key Comparisons Oxidation vs Reduction – loss of e⁻ vs gain of e⁻. Oxidant vs Reductant – electron acceptor vs electron donor. Inner‑sphere vs Outer‑sphere ET – requires bridging ligand vs no direct bond. Electron‑transfer vs Atom‑transfer redox – transfer of electrons only (fast) vs transfer of whole atoms (can be fast or slow). Positive \(E^\circ\) vs Negative \(E^\circ\) – stronger vs weaker oxidizing agent than H⁺. --- ⚠️ Common Misunderstandings “Oxidant = oxidized species.” The oxidant causes oxidation; it itself is reduced. Assuming a positive \(E^\circ\) guarantees a spontaneous reaction. Only the cell voltage (difference) matters. Swapping cathode and anode in the voltage equation. Remember: cathode = reduction, anode = oxidation. All redox reactions are fast. Atom‑transfer processes (e.g., rust) can be kinetically slow. Treating NAD⁺ as a “metal.” It is an organic cofactor that shuttles electrons, not a metal ion. --- 🧠 Mental Models / Intuition Electron flow “downhill”: electrons travel from the half‑reaction with lower \(E^\circ\) (more negative) to the one with higher \(E^\circ\) (more positive). Redox pair as a two‑way street: the reduced form can give up an electron (become oxidized) and the oxidized form can take it back (become reduced). Bridging ligand = handshake: inner‑sphere ET only occurs when the two species can “shake hands” via a shared ligand. --- 🚩 Exceptions & Edge Cases Atom‑transfer speed: combustion of methane is atom‑transfer but very fast; rusting is also atom‑transfer but very slow. Reducing equivalents: a hydride ion (H⁻) counts as a one‑electron reducing equivalent, not just a free electron. Stepwise oxidation of hydrocarbons: multiple distinct half‑reactions (alcohol → aldehyde/ketone → acid → peroxide) may occur in one overall process. --- 📍 When to Use Which Inner‑sphere vs Outer‑sphere: look for a ligand that can bind both partners → inner‑sphere; otherwise assume outer‑sphere. Standard potentials: use \(E^\circ\) values to decide which species will be reduced (higher \(E^\circ\)) and which will be oxidized (lower \(E^\circ\)). Sacrificial anode: choose a metal with a more negative \(E^\circ\) than the metal you want to protect. Electroplating vs galvanic cell: use electroplating when you need a controlled deposition of metal; use a galvanic cell when you want spontaneous electricity generation. --- 👀 Patterns to Recognize Balanced electrons: the total electrons lost in the oxidation half‑reaction always equal those gained in the reduction half‑reaction. Positive cell voltage = spontaneous: if \(E{\text{cell}} > 0\), the redox reaction proceeds without external energy. Bridging ligand presence in mechanisms: formulas like \(\text{M–X–M'}\) signal inner‑sphere ET. Biological redox couples: NAD⁺/NADH and GSH/GSSG always appear in pairs (oxidized ↔ reduced). --- 🗂️ Exam Traps Sign reversal of \(E^\circ\): picking the wrong sign for a half‑reaction flips the predicted direction. Mix‑up of cathode/anode: many distractors list the anode potential in the cathode slot (or vice‑versa). Assuming rusting is fast: a question may describe a “slow redox” and expect you to identify it as an atom‑transfer process. Confusing reducing equivalents: a choice that lists only electrons and ignores hydride (H⁻) may be a trap. NAD⁺ vs NADH wording: remember NAD⁺ accepts electrons (oxidized form), NADH donates them (reduced form). ---