Galvanic Cells

Exam: Friday, 1230h in Rm. 325. SCH4U0-B.

Galvanic Cells Wednesday, January 14, 2009 1:39 PM

- Based on spontaneous electron transfer reaction. - This type of cell provides useful energy to do work. ○ Battery for you Mp3 players, cell phones, etc. - Uses a chemical reaction to produce electrical energy - Cell potential (Eo cell) ○ Work provided by an electrochemical cell ○ Unit is volts - We can predict the voltage that an electrochemical cell can produce. - To do this we must look at an electrochemical cell. - The diagram shows an electrochemical cell consisting of a copper electrode in a copper II sulphate solution and a zinc electrode in a zinc sulphate solution. ○ The salt bridge (a glass tube referred to as a U tube) allows for the equalization of charge. The salt used should not take part in the reaction. This is filled with a salt slurry. The salt in the salt bridge should be the salts inside the solution (ie. Potassium nitrate, any soluble ionic salt as long as it is not part of the solution) - To predict the voltage produced by this electrochemical cell we must look at a table containing Standard Reduction Potentials. - Standard Reduction Potentials (measured in volts): ○ The tendency of a half cell reaction to occur as a reduction under standard conditions (25oC, 101.325kPa, 1.0M solutions). ○ Metals higher on the table have a greater attraction for electrons  They are oxidizing agent and therefore are reduced.  These reactions will be written as reductions - Metals lower on the table have a lower attraction for electrons ○ They are reducing agent and therefore are oxidized. ○ These reactions will be written as oxidations (as the textbook shows reductions we must write the reverse reaction) - Example 1: Copper/Zinc Electrochemical Cell ○ Copper is higher on the table and is written as a reduction (Cathode)  Cu2+(aq) + 2e- -> Cu(s) Eocell = 0.34 V ○ Zinc is lower on the table and is written as an oxidation (Anode)  Zn(s) -> Zn2+(aq) + 2e- Eocell = -0.76V ○ When we reverse the reaction we change the sign ○ The two half reactions are added together  Cu2+(aq) + 2e- -> Cu(s) Eocell = 0.34 V  Zn(s) -> Zn2+(aq) + 2e- Eocell = 0.76V  Cu2+(aq) + Zn(s) -> Cu(s) + Zn2+ 1.10V  This will produce 1.10 V - Example 2: What voltage will be produced by a cell composed of: ○ Ag(s) | Ag+(aq) | Cu2+(aq) | Cu(s) (Standard notation) ○ Cathode | Electrolyte solution || Electrolyte solution | Anode - Silver is higher than copper on the table and will be written as a reduction (cathode) ○ Ag+(aq) + 1e- -> Ag(s) Eocell = 0.80 V - Copper is lower on the table and is written as an oxidation. (Anode) ○ Cu(s) -> Cu2+(aq) + 2e- Eo cell = -0.34 V - Remember the number of electrons gained and lost must be equal. We must multiply the silver half reaction by 2. Important: When multiplying a half reaction we DO NOT multiply the Eo cell value. - The two half reactions are added together ○ 2Ag+(aq) + 2e- -> 2Ag(s) Eo cell = 0.80 V Eocell = -0.34 V ○ Cu(s) -> Cu2+(aq) + 2e○ 2Ag+(aq) + Cu(s) -> Cu2+ + 2Ag(s) = 0.46 V Electrochemical Cells and Spontaneity - When Eo cell > 0 the reaction is spontaneous - When Eo cell < 0 the reaction is not spontaneous - When Eo cell = 0 the reaction is at equilibrium (no flow of electrons) Electrolytic Cell - The opposite of a galvanic cell. Based of a non-spontaneous electron transfer reaction (Eo cell < 0) - We must supply an external energy source (battery etc.) to force chemical reactions. ○ Example - electroplating

Chemistry Page 1