Which is the anode and which is the cathode in this cell b Write the overall

Which is the anode and which is the cathode in this

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Which is the anode and which is the cathode in this cell? b. Write the overall balanced equation and calculate the cell potential.
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Example 3 A voltaic cell is constructed using electrodes based on the following half reactions: Pb 2+ (aq) + 2e -  Pb(s) E o = -0.13V Au 3+ (aq) +3e -  Au(s) E o = +1.50 V a. Which is the anode and which is the cathode in this cell? The anode will involve the oxidation reaction: Pb(s)  Pb 2+ (aq) + 2e - The cathode will involve the reduction reaction: Au 3+ (aq) + 3e - Au(s) This is because the cell potential of the cell will be the most positive (+0.13 V + 1.50 V)
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Example 3 A voltaic cell is constructed using electrodes based on the following half reactions: Pb 2+ (aq) + 2e -  Pb(s) E o = -0.13V Au 3+ (aq) +3e -  Au(s) E o = +1.50 V b. Write the overall balanced equation and calculate the cell potential. The overall balanced reaction is: 3Pb(s) + 2Au 3+ (aq) 3Pb 2+ (aq) + 2Au and the cell potential is : = + = +0.13V + 1.50 V = +1.63V
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Important notes 1. When you multiply an equation by a number to balance the number of electrons, the cell potential value DOES NOT change proportionately. (Look at the formula : E = E o – nRT ln Q (when the concentrations of the products and reactants change, they cancel out in the formula) 2. The cell potential = E anode + E cathode (It is the same as saying E cell = E red + E ox )
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Applications Electrochemistry can be applied as follows: Batteries: a portable, self-contained electrochemical power source that consists of one or more voltaic cells. Batteries can be primary cells (cannot be recharged when “dead”—the reaction is complete) or secondary cells (can be recharged). Prevention of corrosion (“rust-proofing”) Electrolysis
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Some Examples of Batteries Lead–acid battery: reactants and products are solids, so Q is 1 and the potential is independent of concentrations; however, made with lead and sulfuric acid (hazards).
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Some Examples of Batteries Alkaline battery: most common primary battery. They are a type of primary battery dependent upon the reaction between zinc and manganese(IV) oxide.
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Some Examples of Batteries Lithium-ion batteries: rechargeable, light; produce more voltage than Ni-based batteries.
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Some Examples of Batteries Ni–Cd and Ni–metal hydride batteries: lightweight, rechargeable; Cd is toxic and heavy, so hydrides are replacing it.
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Fuel cells When a fuel is burned, the energy created can be converted to electrical energy. Usually, this conversion is only 40% efficient, with the remainder lost as heat. The direct conversion of chemical to electrical energy is expected to be more efficient and is the basis for fuel cells. Fuel cells are NOT batteries; the source of energy must be continuously provided.
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Fuel cells - hydrogen In this cell, hydrogen and oxygen form water. The cells are twice as efficient as combustion.
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