Electrochemistry Chapter 21

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 elementO.N. is zero 2.Monatomic ionO.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; OF 2, +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 lossOIL Reduction involves gainRIG

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 H 2 O Balance hydrogen using H + Basic conditions: Balance oxygen using OH - Balance hydrogen using H 2 O 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.

OxidationReduction

Zn 2+ (aq) + Cu (s)  Cu 2+ (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)

E cell = V

Cell Potential, E 0 cell E 0 cell cell potential under standard conditions elements in standard states (298 K) solutions:1 M gases:1 atm

Standard Reduction Potentials E  values for reduction half-reactions with solutes at 1M and gases at 1 atm Cu e   Cu E  = 0.34 V vs. SHE SO 4 2  + 4H + + 2e   H 2 SO 3 + H 2 O E  = 0.20 V vs. SHE

E 0 cell and  G 0 E 0 cell > 0  G 0 0  G 0 < 0Spontaneous E 0 cell 0Not E 0 cell = 0  G 0 = 0Equilibrium

Calculating E 0 cell E 0 cell = E 0 cathode - E 0 anode Br 2(aq) +2V 3+ +2H 2 O (l)  2VO 2+ (aq) + 4H + (aq) + 2Br - (aq) Given:E 0 cell = V E 0 Br2 = V What isE 0 V3+ and is the reaction spontaneous?

E 0 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 C C mol of e - mol of e -

 G 0, E 0, and K At equilibrium:  G 0 = 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. 2H 2(g) + O 2(g)  2H 2 O (l) anode: 2H 2 + 4OH   4H 2 O + 4e  cathode: 4e  + O 2 + 2H 2 O  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  current and time  quantity of charge  moles of electrons  moles of analyte  grams of analyte