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) +NH 4 + (aq) + 2e -  Mn 2 O 3(s) + 2NH 3(g) + 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 -  ZnO (s) + H 2 O (l) + 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) + SO 4 -2 (aq)  PbSO 4(s) + 2e -  Cathode:  PbO 2(s) + SO 4 -2 (aq) + 4H + + 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

Electric PotentialElectric Potential  “driving force” moving electrons through the connecting wire of a voltaic cell.  Units: volts (V)

Electrode Potential (E cell )Electrode Potential (E cell )  Also called “cell voltage”  Difference between the electric potential between an electrode and the solution it is submerged in  Voltaic cell’s potential to do work on the environment through the generation of an electric current  Magnitude indicates amount of current generated through redox reaction Measured by voltmeter  V = J/C

Reduction PotentialReduction Potential  A half-reaction’s likelihood to act as a reduction reaction within a voltaic cell

How to calculate the Electrode Potential (E cell ) for an entire reaction?  E°rxn. = E°oxidation + E°reduction  Look up these values using your reference table on p. 664 of your textbook.  If equation reversed, change the sign (+ or - ) of standard electrode potential

Example 1:Example 1:  Calculate the cell potential for a voltaic cell with the following cell notation.  Fe Fe +3 Ag + Ag

Example 2:Example 2:  Calculate the cell potential for the following voltaic cell.  K K + Na + Na

Homework  Cell Potential Worksheet