Electrochemistry Study Guide

📖 Core Concepts Electrochemistry – studies how electrical potential differences drive chemical changes (redox reactions). Redox (Reduction‑Oxidation) – electron transfer that changes oxidation states. Oxidation: loss of electrons → oxidation state ↑ (reducing agent). Reduction: gain of electrons → oxidation state ↓ (oxidizing agent). Mnemonics: OIL RIG (Oxidation Is Loss, Reduction Is Gain) / LEO GER (Lose Electrons Oxidation, Gain Electrons Reduction). Electrochemical Cell – converts chemical energy ↔ electrical energy. Anode: site of oxidation, electron source. Cathode: site of reduction, electron sink. Electrolyte: ionically conducting, electronically insulating medium that completes the internal circuit. Standard Hydrogen Electrode (SHE) – reference electrode with defined potential 0.00 V (1 atm H₂, [H⁺]=1 M, pH 0). Standard Electrode Potential (E°) – measured relative to SHE; sign tells you whether a half‑cell is a stronger oxidizer (positive) or reducer (negative). Cell Potential (E\cell) – difference between cathode and anode potentials: $$E{\text{cell}} = E{\text{cathode}} - E{\text{anode}}$$ Gibbs Free Energy Relation – $$\Delta G = -n F E{\text{cell}}$$ (negative ΔG ⇔ spontaneous, electrical work obtainable). Nernst Equation (half‑reaction): $$E = E^{\circ} - \frac{RT}{nF}\ln Q$$ At 25 °C (298 K) → $$E = E^{\circ} - \frac{0.0592\ \text{V}}{n}\log{10} Q$$ Faraday’s Laws – (1) mass deposited ∝ charge passed; (2) amount deposited ∝ equivalent weight. $$m = \frac{Q\,M}{nF}$$ --- 📌 Must Remember OIL RIG / LEO GER – oxidation = loss, reduction = gain of electrons. Anode = Oxidation, Cathode = Reduction (both galvanic and electrolytic cells). Cell EMF sign: Positive \(E{\text{cell}}\) → spontaneous (galvanic); negative → non‑spontaneous (electrolysis). Standard potentials: SHE = 0.00 V. More positive E° → stronger oxidizing agent, will act as cathode in a spontaneous cell. More negative E° → stronger reducing agent, will act as anode. ΔG–E\cell relationship: \(\Delta G = -nF E{\text{cell}}\). Nernst factor at 25 °C: 0.0592 V per electron for base‑10 log. Faraday constant: \(F = 96\,485\ \text{C mol}^{-1}\). Electrolysis voltage requirement: must exceed the absolute value of the cell’s standard EMF plus any overpotential. Corrosion protection: sacrificial anode (e.g., Zn) must have more negative E° than the metal to be protected. --- 🔄 Key Processes Calculating Cell Potential (Galvanic) Write half‑reactions (as reductions). Look up E°(reduction) for each. Identify cathode (more positive E°) and anode (more negative E°). Compute \(E{\text{cell}}^{\circ}=E{\text{cathode}}^{\circ}-E{\text{anode}}^{\circ}\). Apply Nernst equation if concentrations differ. Using the Nernst Equation Determine n (electrons transferred). Write the reaction quotient Q (productsⁿ / reactantsⁿ). Plug into \(E = E^{\circ} - \frac{0.0592}{n}\log Q\). Adjust for temperature if not 25 °C: replace 0.0592 V with \(\frac{RT}{nF}\ln10\). Faraday’s First Law – Mass from Charge Calculate total charge: \(Q = I t\) (current × time). Use \(m = \frac{Q M}{nF}\) to find deposited mass or amount of gas evolved. Electrolysis of Aqueous NaCl (Laboratory Cell) Cathode (reduction): water → H₂ (E° = –0.83 V) (favored over Na⁺ reduction). Anode (oxidation): Cl⁻ → Cl₂ (E° = +1.36 V) (favored over water oxidation). Overall: \(\ce{2 NaCl + 2 H2O → H2 + Cl2 + 2 NaOH}\). Corrosion Prevention by Sacrificial Anode Attach metal with more negative E° (e.g., Zn, Mg) to structure. The sacrificial metal oxidizes (anodic), protecting the structural metal (cathodic). --- 🔍 Key Comparisons Galvanic vs. Electrolytic Cell Spontaneity: Galvanic = spontaneous (positive \(E{\text{cell}}\)); Electrolytic = non‑spontaneous (requires external voltage). Electron flow: Galvanic – from anode to cathode through external circuit; Electrolytic – forced by power source opposite direction. Standard Reduction vs. Oxidation Potentials Reduction potential: tabulated as written for reduction. Oxidation potential: simply the negative of the reduction potential. Corrosion of Iron vs. Passivation of Al/Ti Iron: forms porous Fe₂O₃·xH₂O (rust) → continues to corrode. Aluminium/Titanium: form thin, adherent oxide layers → stop further oxidation (self‑passivating). Concentration Cell vs. Standard Cell Standard cell: non‑zero \(E^{\circ}\) from different electrode materials. Concentration cell: \(E^{\circ}=0\); voltage arises solely from concentration difference via Nernst equation. --- ⚠️ Common Misunderstandings “Anode is always positive.” True for galvanic cells (anode negative, cathode positive). In electrolytic cells the anode is positive because the external source drives electrons to the cathode. “The larger the E° value, the larger the current.” E° determines spontaneity, not current magnitude; current depends on resistance, overpotentials, and electrode area. “Water cannot be electrolyzed because its reduction potential is very negative.” Overpotential and catalysts (e.g., Pt) lower the practical voltage; water electrolysis occurs at 1.8–2.0 V. “All metals corrode at the same rate.” Corrosion rate depends on metal’s E°, protective oxide formation, and environmental conductivity (e.g., salts accelerate rust). --- 🧠 Mental Models / Intuition “Electron Highway”: Think of the external wire as a one‑way highway for electrons from the oxidation (anode) “factory” to the reduction (cathode) “store”. “Redox Ladder”: Oxidation state changes are like climbing up or down a ladder; each electron moves you one rung. “Battery as a Hill”: A positive cell potential is like a hill that drives electrons downhill spontaneously; a negative potential means you need to push them uphill (electrolysis). “Concentration Cell as a Balance”: The side with higher ion activity “wins” (gets reduced) while the lower side “loses” (gets oxidized) to restore balance – Le Chatelier in action. --- 🚩 Exceptions & Edge Cases Overpotential: Real electrodes often need extra voltage (overpotential) beyond thermodynamic values, especially for gas evolution (e.g., H₂, O₂, Cl₂). Hydrogen Evolution vs. Metal Deposition: Even though Na⁺ has a very negative reduction potential, water reduction dominates in aqueous solution because of lower overpotential. Passivation: Metals like Al and Ti may appear inert, but under aggressive conditions (e.g., strong acids) the protective oxide can break down. Temperature Dependence: Nernst term \((RT/F)\) grows with temperature, making cell potentials more sensitive to concentration changes at higher T. --- 📍 When to Use Which Choose a galvanic cell when you need to harvest electrical energy from a spontaneous redox reaction (e.g., batteries). Choose an electrolytic cell when you must drive a non‑spontaneous reaction (e.g., metal plating, water splitting, chlorine production). Use the Nernst equation whenever ion concentrations deviate from 1 M or when pressure/partial pressure of gases differs from 1 atm. Apply Faraday’s first law for quantitative predictions of deposited mass or gas volume given charge passed. Select a sacrificial anode for corrosion protection of metals that cannot form a stable passive layer (e.g., steel pipelines). --- 👀 Patterns to Recognize Positive \(E{\text{cell}}\) → spontaneous → galvanic (look for more positive reduction half‑cell). Redox couples with large potential gap → strong driving force (e.g., Zn/Cu). In aqueous solutions, water reduction/oxidation often outranks extreme metal potentials (Na⁺, K⁺ reduction rarely observed). Presence of a salt bridge or porous frit → indicates a galvanic cell designed to keep solutions separate while allowing ion flow. Gas evolution at electrodes → check standard potentials: H₂ (−0.83 V), O₂ (1.23 V), Cl₂ (1.36 V). --- 🗂️ Exam Traps “The anode is always positive.” – Wrong for galvanic cells; correct only for electrolytic cells. Confusing standard reduction potential with oxidation potential – Remember to flip the sign for oxidation. Using the Nernst equation with activities but plugging concentrations directly – In dilute solutions this is acceptable, but in concentrated solutions activity coefficients matter. Assuming the most negative reduction potential will be reduced in aqueous electrolysis – Water’s reduction occurs first because of kinetic favorability. Treating a concentration cell as having non‑zero \(E^{\circ}\). – Its standard potential is zero; all voltage comes from concentration ratio. Neglecting overpotential for gas‑evolving reactions – Leads to underestimating required voltage (e.g., water electrolysis ≈ 2 V, not 1.23 V). ---