1. 9.2 Electrochemical Cells
IB SL Chemistry
Mrs. Page

2. Understandings Voltaic (Galvanic) cells:
Voltaic cells convert energy from spontaneous, exothermic chemical processes to electrical energy.
Oxidation occurs at the anode (negative electrode) and reduction occurs at the cathode (positive electrode) in a voltaic cell.
Electrolytic cells:
Electrolytic cells convert electrical energy to chemical energy, by bringing about non-spontaneous processes.
Oxidation occurs at the anode (positive electrode) and reduction occurs at the cathode (negative electrode) in an electrolytic cell. 2 2

3. Application & Skills
Construction and annotation of both types of electrochemical cells.
Explanation of how a redox reaction is used to produce electricity in a voltaic cell and how current is conducted in an electrolytic cell.
Distinction between electron and ion flow in both electrochemical cells.
Performance of laboratory experiments involving a typical voltaic cell using two metal/metal-ion half-cells.
Deduction of the products of the electrolysis of a molten salt.

4. Law of Conservation Of Energy
Energy cannot be created or destroyed, it can only be transformed from one type to another.
Energy is the capacity to do work.
SI Unit: Joules
Types of Energy: kinetic, potential, heat, light, sound, chemical, mechanical

5. Electrochemical CELLS
Electrochemical cells convert between chemical and electrical energy through the transfer of electrons during redox reactions
All electrochemical cell have:
Electrodes: metallic electric conductor through which an electric current enters or leaves an electrolytic cell.
Electrolyte: A substance that dissociates into ions in solution and gains the capacity to conduct electricity.
Two types of Electrodes:
Anode: where oxidation takes place
Cathode: where reduction takes place

6. Electrochemical CELLS
Exchange between chemical energy and electrical energy
2 Types:
Voltaic (Galvanic) Cells: convert spontaneous, exothermic chemical reactions to electrical energy
Electrolytic Cells: convert electrical energy to chemical energy resulting in non-spontaneous reactions occurring

7. Voltaic (Galvanic) Cells:
Two different half-cells are connected together by a salt bridge to allow electron transfer during the redox reaction.
Produces electrical energy.
The electrons are produced at the half-cell that is most easily oxidized.
Cathode is positive electrode, where reduction occurs (always drawn on the right)
Anode is negative electrode, where oxidation occurs. (on left)

8. how Is a redox reaction is used to produce electricity in a voltaic cell

9. how Is a redox reaction is used to produce electricity in a voltaic cell
The reaction can be used to perform electrical work.
The transfer of electrons takes place through an external pathway.
Metal strips are placed in solutions of their ions. The metal strips are connected by a wire for flow of electrons.
The solutions are connected by a salt bridge or separated by a porous glass barrier. This maintains electrical neutrality.
Electrons flow from the anode to the cathode. 9 9 9

10. EXAMPLE: Daniell Cell (Zn & Cu)
how Is a redox reaction is used to produce electricity in a voltaic cell
EXAMPLE: Daniell Cell (Zn & Cu)
A strip of zinc is placed in a copper solution.
Write the oxidation reaction
Write the reduction reaction
Write the overall reaction
Describe two observable changes 10 10

11. Daniell Cell Zn/Cu Voltaic Cell
Date Book Table 24 Standard Electrode Potentials
If reaction is being reduced, change sign of electrode potential!
Half Reactions: Zn(s)  Zn2+(aq) + 2e- Eo = V Cu2+(aq) + 2e-  Cu(s) Eo = V Net Redox Equation: Cu2 +(aq) + Zn(s)  Zn2+(aq) + Cu(s) Eo = V

12. how Is a redox reaction is used to produce electricity in a voltaic cell
At the Zn Electrode (Anode):
Oxidation occurs
Electrons are produced and flow through the external circuit toward the cathode
Zn2+ ions produced and migrate away from the electrode
Negative ions (anions) from the salt bridge migrate into the solution to balance the increase in positive charges. 12 12 12

13. how Is a redox reaction is used to produce electricity in a voltaic cell
At the Cu electrode (Cathode):
Reduction occurs
Electrons move from the anode and move into the electrode
Cu2+ ions migrate to the electrode and gain electrons producing Cu
Positive ions (cations) from the salt bridge migrate into the solution to balance the decrease in positive charges. 13 13 13

14. VOLTAIC (GALVANIC) CELLS
Salt bridge: creates a circuit by allowing ions to flow through to balance the ionic charges of the solutions…
Typical Salt Bridges (Na2SO4, KCl, KNO3)
Electrons DO NOT flow through the salt bridge 14 14 14

15. Cell Diagrams Simple way to represent voltaic cells.
Anode is always on left and cathode on right (alphabetical as your read left to right)
Salt bridge is represented as parallel vertical lines
For Daniell Cell the Diagram would be:
Zn(s)|Zn2+(aq) || Cu2+(aq)|Cu(s)

16. Cell Diagrams
Show the cell diagram for the voltaic cell below:

17. VOLTAIC (GALVANIC) CELLS 17 17 17

18. Voltaic Cells in Action

19. VOLTAIC CELLS
A voltaic cell similar to that shown in the previous slide is constructed. One electrode compartment consists of a cadmium strip placed in a solution of Cd(NO3)2 and the other has a nickel strip placed in a solution of NiSO4. Cadmium is a more reactive metal than nickel.
Write the half-reactions that occur in the two electrode compartments. Write the overall reaction.
Which electrode is the anode and which is the cathode?
Indicate the signs of the electrodes.
Which way do electrons flow?
In which directions do the cations and anions migrate through the solution? 19 19 19

20. VOLTAIC CELLS
A voltaic cell is constructed. One ½ cell consists of a silver strip placed in a solution of AgNO3 and the other has a nickel strip placed in a solution of Ni2NO3. Nickel is a more reactive metal than silver.
Write the half-reactions that occur in the two electrode compartments. Write the overall reaction.
Which electrode is the anode and which is the cathode?
Indicate the signs of the electrodes.
Which way do electrons flow?
In which directions do the cations and anions migrate through the solution? 20 20

21. Electrolytic Cells
Electrolysis: a process in which a chemical reaction, is brought about by passing an electric current through a solution of electrolytes causing the electrolyte's ions to move toward the negative and positive electrodes.
Electrical energy is used to cause a non- spontaneous chemical reaction to occur.

22. Electrolytic Cells
Electricity is supplied from an external source (battery) and is used to make a non-spontaneous reaction take place.
The substance that conducts electricity in the cell is an electrolyte (substance containing ions).
Electrolytes do not conduct when solid because ions are not free to move and they have no delocalized electrons.
Electrolytes do conduct when molten or dissolved in water because the ions are free to move toward opposite charged electrodes 22

23. Electrolytic Cells
The electrolyte conducts electricity by the movement of ions within it.
2 inert electrodes are placed in the solution and attached to the battery.
Chemical reactions occur at each electrode so that the electrolyte is decomposed in the process.
2NaCl(l)  2Na(l) + Cl2(g) 23

24. Electrolytic Cells
Oxidation occurs at the positive electrode (anode) because negative ions (anions) are attracted to it,
2Cl-(l)  Cl2(g) + 2e-
Reduction occurs at the negative electrode (cathode) because positive ions (cations) are attracted to it.
Na+ + e-  Na(l)
Overall Cell Reaction:
NaCl(l)  2Na(l) + Cl2(g) 24

25. Electrolytic Cells Solving Electrolysis Problems:
Identify species present
Identify which species are attracted to anode (neg. electrode) and cathode (pos. electrode)
Deduce ½ equations at each electrode
Deduce overall cell reaction
Draw and annotate electrolytic cell (show direction of e- and ion flow)
State observations that will take place at each electrode 25

26. Electrolytic Cells
Describe the electrolysis of a molten calcium chloride?
Identify species present
CaCl2  Ca2+ + 2Cl-
Ca2+(l) and 2Cl-(l) are present
Identify which species are attracted to anode (ps. electrode) and cathode (neg. electrode)
Cathode (neg. electrode): Ca2+
Anode (pos. electrode): Cl- 26

27. Electrolytic Cells
Describe the electrolysis of a molten calcium chloride?
Deduce ½ equations at each electrode
Cathode (neg.) Ca2+(l) + 2e-  Ca(l)
Anode (pos.) 2Cl-(l)  Cl2(g) + 2e-
Deduce overall cell reaction
Ca2+(l) + 2Cl-(l)  Ca(l) + Cl2(g) 27

28. Electrolytic Cells
Describe the electrolysis of a molten calcium chloride?
Draw and annotate electrolytic cell (show direction of e- and ion flow)
State observations that will take place at each electrode
At anode there will be bubbles of chlorine gas
At the cathode a pool of liquid calcium will form. 28

29. Electrolytic Cells in Action

30. Comparing Voltaic & Electrolytic Cells
Create a chart, differentiating between:
Charge of anode and cathode
Site of oxidation and reduction
Direction of electron flow
Electrode that gains mass
Electrode that loses mass
Spontaneous or not

31. Homework
Read pp. 226 – 232
Quick Questions pp. 229 & 232