Chapter 21 Thanks to D Scoggin Cabrillo College Electrochemistry Chapter 21 Thanks to D Scoggin Cabrillo College

Electrochemistry and Redox Oxidation-reduction: “Redox” Electrochemistry: study of the interchange between chemical change and electrical work Electrochemical cells: systems utilizing a redox reaction to produce or use electrical energy

Redox Review Redox reactions: electron transfer processes Oxidation: loss of 1 or more e- Reduction: gain of 1 or more e- Oxidation numbers: imaginary charges (Balancing redox reactions)

Oxidation Numbers (O.N.) 1. Pure element O.N. is zero 2. Monatomic ion O.N. is charge 3. Neutral compound: sum of O.N. is zero Polyatomic ion: sum of O.N. is ion’s charge *Negative O.N. generally assigned to more electronegative element

Oxidation Numbers (O.N.) 4. Hydrogen assigned +1 (metal hydrides, -1) 5. Oxygen assigned -2 (peroxides, -1; OF2, +2) 6. Fluorine always -1

Oxidation-reduction Oxidation is loss of e- O.N. increases (more positive) Reduction is gain of e- O.N. decreases (more negative) Oxidation involves loss OIL Reduction involves gain RIG

Redox Oxidation is loss of e- causes reduction “reducing agent” Reduction is gain of e- causes oxidation “oxidizing agent”

Balancing Redox Reactions 1. Write separate equations (half-reactions) for oxidation and reduction 2. For each half-reaction a. Balance elements involved in e- transfer b. Balance number e- lost and gained 3. To balance e- multiply each half-reaction by whole numbers

Balancing Redox Reactions: Acidic 4. Add half-reactions/cancel like terms (e-) 5. Acidic conditions: Balance oxygen using H2O Balance hydrogen using H+ Basic conditions: Balance oxygen using OH- Balance hydrogen using H2O 6. Check that all atoms and charges balance

Examples Acidic conditions: Basic conditions:

Types of cells Voltaic (galvanic) cells: a spontaneous reaction generates electrical energy Electrolytic cells: absorb free energy from an electrical source to drive a nonspontaneous reaction

Common Components Electrodes: conduct electricity between cell and surroundings Electrolyte: mixture of ions involved in reaction or carrying charge Salt bridge: completes circuit (provides charge balance)

Electrodes Anode: Oxidation occurs at the anode Cathode: Reduction occurs at the cathode Active electrodes: participate in redox Inactive: sites of ox. and red.

Voltaic (Galvanic) Cells A device in which chemical energy is changed to electrical energy. Uses a spontaneous reaction.

Oxidation Reduction

Zn2+(aq) + Cu(s)  Cu2+(aq) + Zn(s) Zn gives up electrons to Cu “pushes harder” on e- greater potential energy greater “electrical potential” Spontaneous reaction due to relative difference in metals’ abilities to give e- ability of e- to flow

Cell Potential Cell Potential / Electromotive Force (EMF): The “pull” or driving force on electrons Measured voltage (potential difference)

Ecell = +1.10 V

Cell Potential, E0cell E0cell cell potential under standard conditions elements in standard states (298 K) solutions: 1 M gases: 1 atm

Standard Reduction Potentials E0 values for reduction half-reactions with solutes at 1M and gases at 1 atm Cu2+ + 2e  Cu E0 = 0.34 V vs. SHE SO42 + 4H+ + 2e  H2SO3 + H2O E0 = 0.20 V vs. SHE

E0cell and DG0 E0cell > 0 DG0 < 0 Spontaneous E0cell < 0 DG0 > 0 Not E0cell = 0 DG0 = 0 Equilibrium

E0cell = E0cathode - E0anode Calculating E0cell E0cell = E0cathode - E0anode Br2(aq)+2V3+ +2H2O(l)  2VO2+(aq)+ 4H+(aq)+ 2Br-(aq) Given: E0cell = +1.39 V E0Br2 = +1.07 V What is E0V3+ and is the reaction spontaneous?

E0 values More positive: Stronger oxidizing agent More readily accepts e- More negative: Stronger reducing agent More readily gives e- Stronger R.A. + O.A.  Weaker R.A. + O.A.

Free Energy and Cell Potential n: number of moles of e- F: Faraday’s constant 96485 C mol of e-

DG0, E0, and K At equilibrium: DG0 = 0 and K = Q At 298 K:

Nernst Equation Under nonstandard conditions

Concentration Cells . . . a cell in which both compartments have the same components but at different concentrations

Batteries A battery is a galvanic cell or, more commonly, a group of galvanic cells connected in series.

Fuel Cells Galvanic cells Reactants are continuously supplied. 2H2(g) + O2(g)  2H2O(l) anode: 2H2 + 4OH  4H2O + 4e cathode: 4e + O2 + 2H2O  4OH

Corrosion Some metals, such as copper, gold, silver and platinum, are relatively difficult to oxidize. These are often called noble metals.

Electrolysis Forcing a current through a cell to produce a chemical change for which the cell potential is negative.

Stoichiometry How much chemical change occurs with the flow of a given current for a specified time? current and time  quantity of charge  moles of electrons  moles of analyte  grams of analyte