Applied Electrochemistry Lecture Notes PDF

# Lec 1, 2 ## Corrosion - Destructive attack of metal by chemical or electrochemical reaction with its environment. - Relation between metal and environment: - Metal + Environment → Product - Loss in weight: Decrease in mechanical properties - Increase in weight: Due to the presence of...

# Lec 1, 2 ## Corrosion - Destructive attack of metal by chemical or electrochemical reaction with its environment. - Relation between metal and environment: - Metal + Environment → Product - Loss in weight: Decrease in mechanical properties - Increase in weight: Due to the presence of undesirable layer, passive layer, metal oxide layer. - Types of electrochemical cells: - Galvanic Cell (Primary Cell) - related to electrode - Example: Danial Cell → dissimilar electrode - Diagram: - Two half cells: Zn half-cell and Cu half-cell - Zn electrode in ZnSO₄ solution - Cu electrode in CuSO₄ solution - Overall reaction: Zn + Cu²⁺ → Zn²⁺ + Cu - Net equation: Zn/Zn²⁺ || Cu²⁺/Cu ## At Open Circuit - Active Zn - Zn Dissolution → Zn²⁺ + 2e^- (Discharge) - At the beginning, rate of oxidation = rate of reduction - Rate of dissolution = Rate of discharge then form electrode potential. - Cu²⁺ + 2e^- → Cu (Discharge) - Dissolution - Metal Ion in the solution will reduce → reduction in the concentration of metal ions in the solution. - At equilibrium, rate of dissolution = rate of discharge. ## Electrode Potential - Potential appears when the metal is in equilibrium with it's ions. - Standard electrode potential, potential when the electrolyte concentration is unity, pressure 1 atm, and temperature 298K - T = 25^oC - P = 1 atm - C = 1 molarity ## At Closed Circuit - Zn → Zn²⁺ + 2e^- (Anode, -ve) - Cu²⁺ + 2e^- → Cu (Cathode, +ve) - Net equation: Zn + Cu²⁺ → Zn²⁺ + Cu - Zn/ZnSO₄ // CuSO₄/Cu → Zn/Zn²⁺ || Cu²⁺/Cu - Direction of current is the movement of +ve ions, direction of electron withdrawal. - External Circuit → +ve to -ve, or Cathode to Anode - Internal Circuit → +ve to -ve, or Cathode to Anode ## Local Action Theory - Commercial Zn: Contains impurities that act as a separate cathode - Zn will corrode, creating local action current - Local action cell: 2H⁺ + 2e^- → H2↑ - Reduction can occur on the local electrode. - The value of local action current is very small. - Theory is wrong - Pure Zn in dil HCl → No Corrosion. - Commercial Zn in dil HCl → Corrosion ## Types of Galvanic Cells - Dry O2 + Metal → No Corrosion - Dry O2 + H2O(l) → Corrosion ### Dissimilar Electrodes: - Two electrodes "dissimilar". - Different electronegativity - Example: Danial Cell - Active electrode: Anode (Potential Difference) - Noble electrode: Cathode ### Strain Effect - Strained metal → Anode (Corrosion, ppt.) - Unstrained metal → Cathode (Deposition) - Example: Processing of metal ### Scratch Effect - Scratched metal → Anode - Unscratched metal → Cathode ### Grain Boundaries - Grain boundaries → Anode - Grain → Cathode ### Grain Size - Small grains → Anode - Large grains → Cathode ### Surface Condition - New section area → Anode - Old section area → Cathode ### Curvature Effect - Convex metal area → Anode - Straight metal area → Cathode ### Concentration Cells - Pt, H₂/HCl || HCl/H₂, Pt #### Different in Conc of electrolyte - Two electrodes "same" immersed in the same electrolyte with different concentration. - Lower Concentration → Anode - Higher Concentration → Cathode #### Different Aeration: - Same electrode in same electrolyte but on electrode is aerated, and another electrode is deaerated - Deaerated (Low one of O₂) → Anode - Aerated (High one of O₂) → Cathode ### Different in Motion - Higher velocity of electrolyte, Cathode - Lower velocity of electrolyte, Anode ### Different Illumination - Dark region of the metal will be anodic to the bright region - Dark region, Anode - Light region, Cathode ### Different Temperature - Low temp, Anode - High temp, Cathode ## Thermodynamics - ΔG represents the change in free energy of a spontaneous reaction, non-spontaneous reaction or reaction at equilibrium - ΔG < 0, spontaneous reaction - ΔG > 0, non-spontaneous reaction - ΔG = 0, equilibrium - Relation between potential and free energy: - ΔG = -ZFE - Standard Condition: ΔG^o = -ZFE^o - ΔG = Gib's Free energy - Z = number of electrons - F = Faraday's constant = 96500C - E = Electrode Potential - At equilibrium: - ΔG = ΔG^o + RTln Q_r/ Q_p - Nernt equation: - E = E^o - RTln Q_r/ Q_p/ ZF - If Red = +Ve: - ΔG = -Ve, Reduction is spontaneous - If Red = -Ve: - ΔG = +Ve, Reduction is non-spontaneous. - Oxidation is spontaneous. - E_cell = E_Cathode (Reduction) - E_Anode (Reduction) - Example: Steel is placed aerated seawater pH=7 - Corrosion of steel is possible or not? Why? - Anode, Fe → Fe²⁺ + 2e^-, E_red = -0.447V - Cathode, O₂ + 2H₂O + 4e^- → 4OH^-, E_red = 0.82V - Net equation: 2Fe + 0₂ + 2H₂O → 2Fe²⁺ + 4OH^- - E_cell = E_Cathode -E_Anode = 0.82 - (-0.447) = 1.267V - ΔG= -ZFE, ΔG -ve, Spontaneous reaction.