Example 2:Example 2:  Calculate the values of Δ G° and K eq at 25°C for the following reaction:  3Mg (s) + 2Al +3 (1M)  3Mg +2 (1M) + 2Al (s)

Batteries  Connects objects  Converts chemical---electrical energy  Two or more voltaic cells connected to each other

Types of BatteriesTypes of Batteries 1)Dry Cells  Alkaline batteries 2)Lead Storage Batteries 3)Fuel Cells

Dry Cells—GeneralDry Cells—General  Composed of “primary cells”  Irreversible redox reactions, not capable of being recharged  Fairly expensive and maximum voltage of 1.55V  “Typical batteries”---seen with flashlights, other electronics

Dry Cells—In DetailDry Cells—In Detail  Anode:  Zn (s)  Zn +2 (aq) + 2e -  Cathode:  Mixture of carbon rod and MnO 2(s)  Electrolyte mixture of NH 4 Cl and ZnCl 2  2MnO 2(s) +H + (aq) + 1e -  MnO(OH) (s) THEN---- 2MnO(OH) (s)  Mn 2 O 3(s) + H 2 O (l)

Dry Cells—Alkaline CellsDry Cells—Alkaline Cells  Longer shelf-life, more current generated over time, more expensive  Different electrolyte—KOH  Same half-reactions but occur in basic solution.  Reduction: 2MnO 2(s) + H 2 O (l) + 2e -  Mn 2 O 3(s) + 2OH - (aq)  Oxidation: Zn (s) + 2OH -  Zn(OH) 2(s) + 2e -  No decrease in voltage as current is generated.

Lead Storage BatteryLead Storage Battery  Made by several lead plates connected together and all in a H 2 SO 4 solution—composed of “secondary cells”  Reversible  Rechargeable

Lead Storage Battery—In Detail  Many voltaic cells—increase current capacity  Each voltaic cell has approximately 2V capacity, 6 cells connected together and results in a 12V battery  PbSO 4(s) produced at both electrodes

Lead Storage Battery—In Detail  Anode:  Pb (s) + HSO 4 - (aq)  PbSO 4(s) + H + (aq) + 2e -  Cathode:  PbO 2(s) + HSO 4 - (aq) + 3H + + 2e -  PbSO 4(s) + 2H 2 O (l)  Electrolyte solution is sulfuric acid (H 2 SO 4 )

Lead Storage Battery— Discharging/Recharging  Discharging  PbSO 4 collects at electrodes  Water dilutes sulfuric acid solution  Recharging  Requires external energy source  Forces electrons to move in the direction of the reverse reaction  Produces negative cell potential, nonspontaneous

Fuel CellsFuel Cells  Electrochemical cell that uses a reaction with oxygen for electrical energy  Components exist outside typical battery  Fuel + Oxygen  Oxidation products

Example  Hydrogen—Oxygen Fuel Cell

Iron Corrosion—in generalIron Corrosion—in general  A redox reaction in a makeshift voltaic cell  Processes are separate on metal, but often occur at same areas  Spontaneous process  Electrons move through metal and electrolyte solution is air

Iron Corrosion—now in detail  Cracked/dented iron more susceptible to corrosion  Higher energy state---oxidation likely  Oxygen from air oxidizes iron  Occurs at “anodic areas”  Fe (s)  Fe e -  Electrons travel along the iron to “cathodic areas” where reduction occurs  Oxygen is reduced  O 2(g) + 2H 2 O (l) + 4e -  4OH - (aq) **Iron goes through 2 oxidations before “rust” forms.

Corrosion ProtectionCorrosion Protection  Multiple ways  2 main ways 1) Galvanized Iron 2) Cathodic Protection

Corrosion Protection— Galvanized Iron  Iron coated with more reactive metal (Zn)  Outer layer of zinc placed around iron  Zinc reacts with oxygen in the place of iron  Zinc sacrifices itself—goes through corrosion (Ex. Galvanized iron nails)

Corrosion Protection— Cathodic Protection  Iron/steel connected directly or indirectly to an active metal (Mg, Al, or Zn)  Active metal sacrifices itself in place of the iron/steel  “Sacrificial anode”  Iron acts as the cathode and reduction occurs there  Ex. Ships, plumbing, pipes

Homework  Electrochemistry Review Worksheet