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Electrochemistry Chapter 17. Electrochemistry The branch of chemistry that links chemical reactions to the production or consumption of electrical energy.

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Presentation on theme: "Electrochemistry Chapter 17. Electrochemistry The branch of chemistry that links chemical reactions to the production or consumption of electrical energy."— Presentation transcript:

1 Electrochemistry Chapter 17

2 Electrochemistry The branch of chemistry that links chemical reactions to the production or consumption of electrical energy. In chemistry, electrical energy is stored in electrons. – In other words, electrochemistry is based upon the principles oxidation-reduction reactions.

3 Redox Revisited Earlier this year, we examined Redox reactions and how to write acid-base half reactions To review principles of redox reactions: – A redox reaction is the sum of two half reactions, the reduction and oxidation reactions. – Reduction and oxidation reactions happen simultaneously, so the number of electrons gained and lost must match exactly – Oxidation=loss of electrons – Reduction=gain of electrons.

4 Redox revisited Lets look at a common reaction of Zinc metal immersed in a copper sulfate solution. The entire reaction is represented as a single replacement reaction where the blue CuSO 4 solution becomes clear as the Zinc replaces the copper: Zn + CuSO 4 → Cu + ZnSO 4

5 Redox Revisited Zn + CuSO 4 → Cu + ZnSO 4 The Redox half reactions are then represented as: Zn (s) → Zn +2 (aq) + 2e - Cu +2 (aq) + 2e - → Cu (s) Zn (s) + Cu +2 (aq) + 2e - → Cu (s) + Zn +2 (aq) + 2e - So: Zn (s) + Cu +2 (aq) → Cu (s) + Zn +2 (aq)

6 Electrochemical (Galvanic) cells An apparatus that converts chemical energy into electrical work, or vice versa – Contains two compartments – Bridge that allows flow of energy (electrons) “salt bridge” Usually a piece of tubing filled with an electrolyte

7 Electrochemical (Galvanic) cells Anode-compartment in which the oxidation half reaction takes place Cathode-compartment which the reduction half reaction takes place. – We represent the reactions that take place using cell diagrams Cell diagrams are symbols that show how the components of an electrochemical cell are connected.

8 Electrochemical (Galvanic) cells Salt bridge replaced with porous disk to allow ion flow and minimum mixing of solutions. Oxidizing agent “pulls” electrons through wire from reducing agent. Let there be light.

9 Homework Pg #

10 Cell Potential The “pull” on the electrons is called the cell potential ( E ° cell ), or “electromotive force” (emf), is measured in volts. – Volt: 1 joule of work per coulomb of charge transferred. 1 J/C – coulomb: defined as the charge transported by a constant current of one ampere in one second:amperesecond

11 Standard Reduction Potentials E ° --standard reduction potentials in Volts. – E ° cell = E ° cathode - E ° anode Pay close attention to sign of E for certain reactions If the reaction is in reverse, change the sign for the reduction potential. See Table 17.1 on page 843 in your book. All reduction potentials are given with all solutes at 1M and all gases at 1 atm pressure. E ° cell is always positive

12 Standard Reduction Potentials Zn + CuSO 4 → Cu + ZnSO 4 The Redox half reactions are then represented as: Zn (s) → Zn +2 (aq) + 2e - - E ° anode =.76 Cu +2 (aq) + 2e - → Cu (s) E ° cathode =.34 Zn (s) + Cu +2 (aq) → Cu (s) + Zn +2 (aq) E ° cell = 1.1V This cell produces 1.1 volts of electrical energy (work).

13 Standard Reduction Potentials Another example, consider the galvanic cell based on the reaction: Al +3 (aq) + Mg (s) → Al (s) + Mg 2+ (aq) Give the redox half reactions, make sure to balance the reactions (equal # of electrons), and calculate E ° cell ( E ° for half reactions on table 17.1)

14 Standard Reduction Potentials Al +3 (aq) + Mg (s) → Al (s) + Mg 2+ (aq) The Redox half reactions are then represented as: 3(Mg (s) → Mg +2 (aq) + 2e) - - E ° anode =2.37V 2(Al +3 (aq) + 3e - → Al (s) ) E ° cathode =-1.66V 2Al +3 (aq) + 3Mg (s) → 2Al (s) + 3Mg +2 (aq) E ° cell =.71V This cell produces.71 volts of electrical energy (work).

15 Standard Reduction Potentials Sometimes there are multiple possibilities for redox potentials, in this case, pay close attention to what the equation they give states the cell is based on. Example: a galvanic cell is based on the reaction: – MnO 4 - (aq) + H + (aq) + ClO 3 - (aq) → ClO 4 - (aq) + Mn +2 (aq) + H 2 O (l) There are multiple oxidation reactions for MnO 4 - (aq), so you must consult table 17.1 and match the reactants and products.

16 Representing Cells with Line Notations Consider key components of this galvanic cell: Zinc solid electrode Zinc ions in solution Spectator ions Copper ions in solution Copper solid electrode Salt bridge

17 Steps to Representing Cells with Line Notations Rule #1…list everything Separate the cathode/anode with the double line notation (II) that represents the salt bridge Separate substances in different states of matter in the same compartment with a single line (I), separate substances in the same state in the same compartment with a comma. Dispense spectator ions (usually water) When there is no solid electrode listed, assume there is Platinum (Pt) in both.

18 Representing Cells with Line Notations Consider key components of this galvanic cell: Zinc solid electrode Zinc ions in solution Spectator ions Copper ions in solution Copper solid electrode Salt bridge

19 homework Pg 880, #’s odd

20 Determining Spontaneity in Galvanic cells 1. Any reduction reaction is spontaneous when paired with the reverse of any reaction below it on Table If the Cell potential calculated is negative, the reaction is not spontaneous (yes, I know I told you that the cell potential is always positive…in a cell that works).

21 Cell Potential, Electrical Work, And Free Energy (∆G) Potential Difference (V)= work (J)/charge© – When a cell produces a current (V), the cell potential is positive, and the current can be used to do work. – Since the work does not stay in the system (the cell), the sign for work is negative. So: q=charge in coulombs transferred from anode to cathode

22 Cell Potential, Electrical Work, And Free Energy (∆G) Work measured in Joules Coulombs based on Faraday Constant Cell potential difference in V or J/C The Faraday Constant: the charge on one mole of electrons is 96,485 Coulombs of charge. When 1.33 moles is transferred: = 1.33 mol e - x 96,485 C/mol e - This equation calculates amount of work done. HOWEVER, since some work and energy is always lost to the surroundings as heat, there is a way to calculate Maximum work

23 Cell Potential, Electrical Work, And Free Energy (∆G) All galvanic cells have a maximum potential that they never reach because of energy lost as heat. To calculate maximum work, use the maximum potential in your equation. And then the Maximum work equals ∆G.

24 Dependence of a Cell on Concentration Simply put…if the product concentration is raised above 1.0M, E ° cell will be less than what is listed 17.1 If the reactant concentration is above 1.0M, E ° cell will be greater than listed in table 17.1

25 Dependence of a Cell on Concentration The dependence of cell potential on concentration relies directly on the dependence of free energy on concentration. Remember that And since And Then: “Nerst Equation”

26 Nerst Equation At 25°C: n= moles of electrons

27 homework Pg 881, #’s odd.


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