Introduction to Electrochemical Cells: Parts and Functions

1. Chapter 20 Preview Lesson Starter Objectives Electrochemical Cells

2. Section 1 Introduction to Electrochemistry Chapter 20 Lesson Starter • In your text book the temperature of the aqueous CuSO4 solution to rise when electrons are transferred directly from Zn atoms to Cu2+ ions. • This increase in temperature indicates that chemical energy in the form of heat is released. • DEMONSTRATION • Energy in an Electrochemical Cell

3. Section 1 Introduction to Electrochemistry Chapter 20 Objectives Identify parts of an electrochemical cell and their functions. Write electrode half reactions for cathodes and anodes.

4. Section 1 Introduction to Electrochemistry Chapter 20 Because oxidation-reduction reactions involve electron transfer, the net release or net absorption of energy can occur in the form of electrical energy rather than as heat. The branch of chemistry that deals with electricity-related applications of oxidation-reduction reactions is calledelectrochemistry.

5. Section 1 Introduction to Electrochemistry Chapter 20 Electrochemical Cells Oxidation-reduction reactions involve a transfer of electrons. If the two substances are in contact with one another, a transfer of energy as heat accompanies the electron transfer. If the substance that is oxidized is separated from the substance that is reduced, the electron transfer is accompanied by a transfer of electrical energy instead of energy as heat.

6. Section 1 Introduction to Electrochemistry Chapter 20 Electrochemical Cells, continued A porous barrier,or salt bridge can be used to separate the oxidation and reduction half-reactions.

7. Ion Movement Through a Porous Barrier Section 1 Introduction to Electrochemistry Chapter 20

8. Section 1 Introduction to Electrochemistry Chapter 20 Electrochemical Cells, continued Electrons can be transferred from one side to the other through an external connecting wire. Electric current moves in a closed loop path, or circuit,so this movement of electrons through the wire is balanced by the movement of ions in solution. An electrodeis a conductor used to establish electrical contact with a nonmetallic part of a circuit, such as an electrolyte.

9. Electron Pathway in an Electrochemical Cell Section 1 Introduction to Electrochemistry Chapter 20 Electron Pathway in an Electrochemical Cell

10. Section 1 Introduction to Electrochemistry Chapter 20 Electrochemical Cell Click below to watch the Visual Concept. Visual Concept

11. Section 1 Introduction to Electrochemistry Chapter 20 Parts of an Electrochemical Cell Click below to watch the Visual Concept. Visual Concept

12. Section 1 Introduction to Electrochemistry Chapter 20 Electrochemical Cells, continued Half-Cell A single electrode immersed in a solution of its ions is a half-cell. The electrode where oxidation occurs is called theanode. • example: Zn(s) Zn2+(aq) + 2e− The electrode where reduction occurs is called the cathode. • example: Cu2+(aq) + 2e− Cu(s)

13. Section 1 Introduction to Electrochemistry Chapter 20 Half-Cell Click below to watch the Visual Concept. Visual Concept

14. Section 1 Introduction to Electrochemistry Chapter 20 Electrochemical Cells, continued Half-Cell, continued • Both oxidation and reduction must occur in an electrochemical reaction. • The two half-cells taken together make an electrochemical cell. The Complete Cell • An electrochemical cell may be represented by the following notation: anode electrode|anode solution||cathode solution|cathode electrode • The double line represents the salt bridge, or the porous barrier.

15. Section 1 Introduction to Electrochemistry Chapter 20 Electrochemical Cells, continued The Complete Cell, continued • The Zn/Cu electrochemical cell, can be written as Zn(s)|Zn2+(aq)||Cu2+(aq)|Cu(s). • The electrochemical reaction can be found by adding the anode half-reaction to the cathode half-reaction. • The overall (or net) reaction for the Zn/Cu cell is Zn2+(aq) + Cu(s). Zn(s) + Cu2+(aq) An electrochemical cell that consists of this Zn and Cu reaction is called the Daniell Cell.

16. Section 1 Introduction to Electrochemistry Chapter 20 Half-Reaction Equation Click below to watch the Visual Concept. Visual Concept

17. Section 2 Voltaic Cells Chapter 20 Preview Lesson Starter Objectives How Voltaic Cells Work Corrosion and Its Prevention Electrical Potential

18. Section 2 Voltaic Cells Chapter 20 Lesson Starter • Make a list of items that we use batteries to run. • What are the different types of batteries that are used? • alkaline, rechargeable, mercury • What are the advantages and disadvantages of the different types of batteries from a consumer point of view?

19. Section 2 Voltaic Cells Chapter 20 Objectives Describe the operation of voltaic cells, including dry cells, lead-acid batteries, and fuel cells. Identify conditions that lead to corrosion and ways to prevent it. Describe the relationship between voltage and the movement of electrons. Calculate cell voltage/potentials from a table of standard electrode potentials.

20. Section 2 Voltaic Cells Chapter 20 Voltaic cellsuse spontaneous oxidation-reduction reactions to convert chemical energy into electrical energy. Voltaic cells are also called galvanic cells. The most common application of voltaic cells is in batteries.

21. Section 2 Voltaic Cells Chapter 20 Voltaic Cell Click below to watch the Visual Concept. Visual Concept

22. Section 2 Voltaic Cells Chapter 20 How Voltaic Cells Work • Electrons given up at the anode pass along the external connecting wire to the cathode. • The movement of electrons through the wire must be balanced by the movement of ions in the solution. • Dry cells are voltaic cells. • The three most common types of dry cells are the zinc-carbon battery, the alkaline battery, and the mercury battery

23. Particle Models for Redox Reactions in Electrochemical Cells Section 2 Voltaic Cells Chapter 20

24. Galvanic Cell Section 2 Voltaic Cells Chapter 20

25. Section 2 Voltaic Cells Chapter 20 Battery Click below to watch the Visual Concept.

26. Section 2 Voltaic Cells Chapter 20 How Voltaic Cells Work, continued Zinc-Carbon Dry Cells Batteries such as those used in flashlights are zinc-carbon dry cells. Zinc atoms are oxidized at the negative electrode, or anode. The carbon rod is the cathode or positive electrode. MnO2 is reduced in the presence of H2O.

27. Dry Cells Section 2 Voltaic Cells Chapter 20

28. Section 2 Voltaic Cells Chapter 20 How Voltaic Cells Work, continued Alkaline Batteries Alkaline batteries do not have a carbon rod cathode, which allows them to be smaller. The half-reaction at the anode is The reduction at the cathode is the same as that for the zinc-carbon dry cell.

29. Model of a Mercury Cell Section 2 Voltaic Cells Chapter 20

30. Section 2 Voltaic Cells Chapter 20 How Voltaic Cells Work, continued Mercury Batteries • The anode half-reaction is identical to that found in the alkaline dry cell. • The cathode half-reaction is

31. Section 2 Voltaic Cells Chapter 20 Parts of an Acidic and Alkaline Battery Click below to watch the Visual Concept. Visual Concept

32. Section 2 Voltaic Cells Chapter 20 How Voltaic Cells Work, continued Fuel Cells • A fuel cell is a voltaic cell in which the reactants are being continuously supplied and the products are being continuously removed. Cathode: O2(g) + 2H2O(l) + 4e− 4OH−(aq) 4e− + 4H2O(l) Anode: 2H2(g) + 4OH−(aq) Net reaction: 2H2 + O2 2H2O • Fuel cells are very efficient and have very low emissions

33. Fuel Cell Section 2 Voltaic Cells Chapter 20

34. Section 2 Voltaic Cells Chapter 20 Corrosion and Its Prevention One of the metals most commonly affected by corrosion is iron. Rust is hydrated iron(III) oxide. 4Fe(s) + 3O2(g) + xH2O(l) 2Fe2O3 •xH2O(s) The anode and cathode reactions occur at different regions of the metal surface. • Anode: Fe(s) Fe2+(aq) + 2e− • Cathode: O2(g) + 2H2O(l) + 4e− 4OH−(aq)

35. Section 2 Voltaic Cells Chapter 20 Corrosion and Its Prevention, continued For corrosion to occur, water and oxygen must be present with the iron.

36. Section 2 Voltaic Cells Chapter 20 Corrosion and Its Prevention, continued • Coating steel with zinc in a process called galvanizingcan prevent corrosion. • Zinc is more easily oxidized than iron • Zinc will react before the iron is oxidized. • This is called cathodic protection. • The more easily oxidized metal used is called a sacrificial anode.

37. Section 2 Voltaic Cells Chapter 20 Electrical Potential In a voltaic cell, the oxidizing agent at the cathode pulls the electrons through the wire away from the reducing agent at the anode. The “pull,” or driving force on the electrons, is called the electric potential. Electric potential, or voltage, is expressed in units of volts (V), which is the potential energy per unit charge. Current is the movement of the electrons and is expressed in units of amperes, or amps (A).

38. Section 2 Voltaic Cells Chapter 20 Electrical Potential, continued Electrode Potentials The tendency for the half-reaction of either copper or zinc to occur as a reduction half-reaction in an electrochemical cell can be quantified as areduction potential. The difference in potential between an electrode and its solution is known aselectrode potential. This potential difference, or voltage, is proportional to the energy required to move a certain electric charge between the electrodes.

39. Section 2 Voltaic Cells Chapter 20 Electrical Potential, continued Electrode Potentials, continued • The potential difference measured across the complete voltaic cell is easily measured. • It equals the sum of the electrode potentials for the two half-reactions. • An individual electrode potential cannot be measured directly. A relative value for the potential of a half-reaction can be determined by connecting it to a standard half-cell as a reference.

40. Section 2 Voltaic Cells Chapter 20 Electrical Potential, continued Electrode Potentials, continued The standard half-cell is called a standard hydrogen electrode, or SHE. It consists of a platinum electrode dipped into a 1.00 M acid solution surrounded by hydrogen gas at 1 atm pressure and 25°C.

41. Section 2 Voltaic Cells Chapter 20 Electrical Potential, continued Electrode Potentials, continued The anodic reaction for the standard hydrogen electrode is The cathodic reaction is An arbitrary potential of 0.00 V is assigned to both of these half-reactions.

42. Section 2 Voltaic Cells Chapter 20 Electrical Potential, continued Electrode Potentials, continued • The potential of a half-cell under standard conditions measured relative to the standard hydrogen electrode is astandard electrode potential,E0. • Electrode potentials are expressed as potentials for reduction. • Effective oxidizing agents have positive E0values. • example: Cu2+ and F2 • Effective reducing agents have negative E0values. • example: Li and Zn

43. Comparing Reduction Potentials of Various Metals Section 2 Voltaic Cells Chapter 20

44. Standard Reduction Potentials Section 2 Voltaic Cells Chapter 20

45. Standard Electrode Potentials Section 2 Voltaic Cells Chapter 20

46. Standard Electrode Potentials Section 2 Voltaic Cells Chapter 20

47. Section 2 Voltaic Cells Chapter 20 Electrode Potential Click below to watch the Visual Concept. Visual Concept

48. Section 2 Voltaic Cells Chapter 20 Electrical Potential, continued Electrode Potentials, continued When a half-reaction is written as an oxidation reaction, the sign of its electrode potential is reversed. oxidation half-reaction: Zn Zn2+ + 2e− E0= +0.76 V reduction half-reaction: Zn2+ + 2e− Zn E0= −0.76 V

49. Section 2 Voltaic Cells Chapter 20 Electrical Potential, continued Electrode Potentials, continued • Standard electrode potentials can be used to predict if a redox reaction will occur spontaneously. • A spontaneous reaction will have a positive value for E0cell. E0cell= E0cathode− E0anode • The half-reaction that has the more negative standard reduction potential will be the anode.

50. Section 2 Voltaic Cells Chapter 20 Electrical Potential, continued • Sample Problem A • Write the overall cell reaction, and calculate the cell potential for a voltaic cell consisting of the following half-cells: • an iron (Fe) electrode in a solution of Fe(NO3)3 • and a silver (Ag) electrode in a solution of AgNO3.