9.2 Electrochemical Cells IB SL Chemistry Mrs. Page

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

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.

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

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

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

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)

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

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

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

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 = +0.76 V Cu2+(aq) + 2e-  Cu(s) Eo = +0.34 V --------------------------------------------- Net Redox Equation: Cu2 +(aq) + Zn(s)  Zn2+(aq) + Cu(s) Eo = +1.10 V

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

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

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 https://mindtouch.oneonta.edu/User:thomaste/ELECTROCHEMISTRY-SPRING_2011 14 14 14

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)

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

VOLTAIC (GALVANIC) CELLS https://mindtouch.oneonta.edu/User:thomaste/ELECTROCHEMISTRY-SPRING_2011 17 17 17

Voltaic Cells in Action http://www.youtube.com/watch?v=A0VUsoeT9aM

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

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

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.

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

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

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

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

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

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

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

Electrolytic Cells in Action http://www.youtube.com/watch?v=Jxzs0HKRzEo

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

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