1. REDOX Part 2 - Electrochemistry Text Ch. 9 and 10

2. Electrochemistry  Electrochemistry is the relation the flow of electric current to chemical changes. Chemical changes can produce electrical energy  Electrolysis/ Electrochemical Cells electrical energy converted to chemical energy (breaking of water, electroplating) [Non-spontaneous]  Galvanic cells chemical energy converted to electrical energy (battery) [Spontaneous]

3. Electrolysis / Electrochemical Cells  Terms Cation – positive ions Cathode – where reduction takes place (electrons are added), the negative electrode Anion – negative ions (move towards the anode) Anode – where oxidation occurs (electrons are lost), the positive electrode

4. Electrolysis / Electrochemical Cells Volt (V) – electrical potential difference (electrical pressure) Ampere (A) – rate of flow of electrons or current Coulomb (c) – the quantity of electricity produced by a current of one ampere flowing for one second 96 500 c = mol of electrons

5. Electrolysis / Electrochemical Cells  At times we may wish to calculate the mass of substance produced during a reaction. Using Faraday’s Law, we can do this.  Faraday’s Law For every mol of electrons, there will be a certain amount of product based upon the current applied to the system.  The amount of metal deposited at the cathode is not related to the voltage used but to three factors related to the reduction half-reaction

6. Electrolysis / Electrochemical Cells Example: Na+ + e -  Na  From this reaction we can see that we require 1 mole of electrons to reduce 1 mole of sodium ions. So, to determine the mass of sodium that is deposited on the cathode, we need to know: a) the number of electrons that flowed to the cathode (based on the current and the time the cell operates) b) the charge on the ions (which determines how many electrons are needed per ion) c) the molar mass of the metal

7. Electrolysis / Electrochemical Cells  Ex. If we ran 7.89 A for 1200 s, in a solution of sodium chloride, what mass of sodium metal would be produced?  Ex. 2 What mass of copper will be produced at the cathode from a solution of copper (II) sulphate (CuSO 4 ) by a current of 5.0 amps operating for 20.0 minutes?

8. Electrolysis / Electrochemical Cells  Electroplating  Electrolysis  Lab

9. Electrolytic Cells / Galvanic Cells  It is possible to use oxidation-reduction reactions as a source of electrical energy.  Note that in the method we used for balancing redox reactions, we separated the equation into two half-cell reactions.

10. Electrolytic Cells / Galvanic Cells  For example Cu 2+ + 2e -  Cu  This is the reduction half-cell reaction and will always take place in conjunction with another half- cell reaction which must be an oxidation reaction Zn  Zn 2+ + 2e -  An electrochemical cell is two half-cell reactions arranged so that we can take advantage of the movement of electrons.

11. Electrolytic Cells / Galvanic Cells  Terms Cathode – positive electrode, undergoes reduction, strongest oxidizing agent, gains mass Anode – negative electrode, undergoes oxidation, strongest reducing agent, loses mass Voltage – a measure of the ability to do electrical work. The units of measure are volts. 1 volt = 1.6 x 10-19 Joules/electron = 1 Joule/Coulomb

12. Electrolytic Cells / Galvanic Cells  Voltage of a Battery Electrons will flow from the oxidized substance to the reduced substance. The voltage of the battery is based upon the Reduction Potential (ability to attract electrons to itself in a half reaction) of the substances used. The difference between the two substances is called the Potential Difference (flow of electrons), and is the theoretical voltage of the battery.  Ex. Lithium and Fluorine

13. Electrolytic Cells / Galvanic Cells  Finding Standard Electrical Potential ****Atoms close to the top of the chart will gain electrons, be reduced (Strong Oxidizing Agents) ****Atoms close to the bottom of the chart will lose electrons, be oxidized (Strong Reducing Agents) Notice that all the half reactions are written at reduction. We need to recognize which substance is oxidized and which is reduced.

14. Electrolytic Cells / Galvanic Cells  Ex. Calculate the potential difference between Copper and Tin. What is the expected voltage and which way is the flow of the electrons? Draw the cell, labeling the cathode, anode and all parts.

15. Electrolytic Cells / Galvanic Cells  Ex. Give the voltage for a battery made of Fe (II) and Ni. Identify the anode and the cathode.