Electrochemistry Applications of Redox Review l Oxidation reduction

Electrochemistry Applications of Redox

Review l Oxidation reduction reactions involve a transfer of electrons. l OIL- RIG l Oxidation Involves Gain l Reduction Involves Loss l LEO-GER l Lose Electrons Oxidation l Gain Electrons Reduction

Applications l Moving electrons is electric current. l 8 H++Mn. O 4 -+ 5 Fe+2 +5 e® Mn+2 + 5 Fe+3 +4 H 2 O l Helps to break the reactions into half reactions. l 8 H++Mn. O 4 -+5 e- ® Mn+2 +4 H 2 O l 5(Fe+2 ® Fe+3 + e- ) l In the same mixture it happens without doing useful work, but if separate

l Connected this way the reaction starts l Stops immediately because charge builds up. H+ Mn. O 4 - Fe+2

Galvanic Cell Salt Bridge allows current to flow H+ Mn. O 4 - Fe+2

travels in a complete circuit l Instead of a salt bridge e- l Electricity H+ Mn. O 4 - Fe+2

Porous Disk H+ Mn. O 4 - Fe+2

e- e- Anode e. Reducing Agent Cathode e. Oxidizing Agent

Cell Potential l Reducing agent pushes the electron. l Oxidizing agent pulls the electron. l The push or pull (“driving force”) is called the cell potential Ecell l Also called the electromotive force (emf) l Unit is the volt(V) l = 1 joule of work/coulomb of charge l Measured with a voltmeter

0. 76 H 2 in Anode Zn+2 SO 4 -2 1 M Zn. SO 4 Cathode H+ Cl - 1 M HCl

Standard Hydrogen Electrode l This is the reference all other oxidations are compared to H 2 in l. E º = 0 lº indicates standard + Cl states of 25ºC, 1 H atm, 1 M solutions. 1 M HCl

Cell Potential + Cu+2 (aq) ® Zn+2(aq) + Cu(s) l The total cell potential is the sum of the potential at each electrode. l Zn(s) l Eº cell = EºZn® Zn+2 + Eº Cu+2 ® Cu l We can look up reduction potentials in a table. l One of the reactions must be reversed, so change it sign.

Cell Potential l Determine the cell potential for a galvanic cell based on the redox reaction. l Cu(s) + Fe+3(aq) ® Cu+2(aq) + Fe+2(aq) l Fe+3(aq) + e-® Fe+2(aq) Eº = 0. 77 V l Cu+2(aq)+2 e- ® Cu(s) Eº = 0. 34 V l Cu(s) ® Cu+2(aq)+2 e. Eº = -0. 34 V l 2 Fe+3(aq) + 2 e-® 2 Fe+2(aq) Eº = 0. 77 V

Line Notation solid½Aqueous½solid l Anode on the left½½Cathode on the right l Single line different phases. l Double line porous disk or salt bridge. l If all the substances on one side are aqueous, a platinum electrode is indicated. l For the last reaction l Cu(s)½Cu+2(aq)½½Fe+2(aq), Fe+3(aq)½Pt(s) l

Galvanic Cell l l 1) 2) 3) 4) The reaction always runs spontaneously in the direction that produced a positive cell potential. Four things for a complete description. Cell Potential Direction of flow Designation of anode and cathode Nature of all the componentselectrodes and ions

Practice l Completely describe the galvanic cell based on the following half-reactions under standard conditions. l Mn. O 4 - + 8 H+ +5 e- ® Mn+2 + 4 H 2 O Eº=1. 51 l Fe+3 +3 e- ® Fe(s) Eº=0. 036 V

Potential, Work and DG l emf = potential (V) = work (J) / Charge(C) l E = work done by system / charge l E = -w/q l Charge l -w is measured in coulombs. = q. E = 96, 485 C/mol el q = n. F = moles of e- x charge/mole el w = -q. E = -n. FE = DG l Faraday

Potential, Work and DG DGº = -n. FE º l if E º > 0, then DGº < 0 spontaneous l if E º < 0, then DGº > 0 nonspontaneous l l In fact, reverse is spontaneous. l Calculate DGº for the following reaction: l Cu+2(aq)+ Fe(s) ® Cu(s)+ Fe+2(aq) l Fe+2(aq) + e-® Fe(s) l Cu+2(aq)+2 e- ® Cu(s) Eº = 0. 44 V Eº = 0. 34 V

Cell Potential and Concentration l Qualitatively - Can predict direction of change in E from Le. Châtelier. l 2 Al(s) + 3 Mn+2(aq) ® 2 Al+3(aq) + 3 Mn(s) if Ecell will be greater or less than Eºcell if [Al+3] = 1. 5 M and [Mn+2] = 1. 0 M l Predict [Al+3] = 1. 0 M and [Mn+2] = 1. 5 M l if [Al+3] = 1. 5 M and [Mn+2] = 1. 5 M l if

The Nernst Equation l DG = DGº +RTln(Q) l -n. FE = -n. FEº + RTln(Q) l E = Eº - RTln(Q) n. F l 2 Al(s) + 3 Mn+2(aq) ® 2 Al+3(aq) + 3 Mn(s) Eº = 0. 48 V l Always have to figure out n by balancing. l If concentration can gives voltage, then from voltage we can tell concentration.

The Nernst Equation l As reactions proceed concentrations of products increase and reactants decrease. l Reach equilibrium where Q = K and Ecell = 0 l 0 = Eº - RTln(K) n. F l Eº = RTln(K) n. F l n. FEº = ln(K) RT

Batteries are Galvanic Cells l Car batteries are lead storage batteries. l Pb +Pb. O 2 +H 2 SO 4 ®Pb. SO 4(s) +H 2 O l Dry Cell Zn + NH 4+ +Mn. O 2 ® Zn+2 + NH 3 + H 2 O l Alkaline Zn +Mn. O 2 ® Zn. O+ Mn 2 O 3 (in base) l Ni. Cad l Ni. O 2 + Cd + 2 H 2 O ® Cd(OH)2 +Ni(OH)2

Corrosion l Rusting - spontaneous oxidation. l Most structural metals have reduction potentials that are less positive than O 2. l Fe ® Fe+2 +2 e. Eº= 0. 44 V l O 2 + 2 H 2 O + 4 e- ® 4 OHEº= 0. 40 V l Fe+2 + O 2 + H 2 O ® Fe 2 O 3 + H+ l Reaction happens in two places.

Salt speeds up process by increasing conductivity Water Rust e. Iron Dissolves- Fe ® Fe+2

Preventing Corrosion l Coating to keep out air and water. l Galvanizing - Putting on a zinc coat l Has a lower reduction potential, so it is more. easily oxidized. l Alloying with metals that form oxide coats. l Cathodic Protection - Attaching large pieces of an active metal like magnesium that get oxidized instead.

Electrolysis l Running a galvanic cell backwards. l Put a voltage bigger than the potential and reverse the direction of the redox reaction. l Used for electroplating.

1. 10 e- e- Zn 1. 0 M Zn+2 Anode 1. 0 M Cu+2 Cathode Cu

e- A battery >1. 10 V Zn 1. 0 M Zn+2 Cathode 1. 0 M Cu+2 e- Cu Anode

Calculating plating l Have to count charge. l Measure current I (in amperes) l 1 amp = 1 coulomb of charge per second lq = I x t l q/n. F = moles of metal l Mass of plated metal l How long must 5. 00 amp current be applied to produce 15. 5 g of Ag from Ag+

Other uses l Electroysis of water. l Seperating mixtures of ions. l More positive reduction potential means the reaction proceeds forward. l We want the reverse. l Most negative reduction potential is easiest to plate out of solution.