G12A Chapter 5 Section 5.1 Voltaic cell (1).pptx

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1. G12A Chemistry: Chapter 5: Electrochemistry Lesson 1 : Voltaic Cell

2. Chapter 5: Electrochemistry Lesson 1: Voltaic Cell

3. Starter Activity

4. Electrochemistry A camera trap captured this image of a mountain lion. Camera traps are a noninvasive way to study animals by using a sensor that triggers a camera’s shutter when the animal approaches. Batteries power both the camera and the sensor. Chemical energy can be converted to electric energy and electrical energy to chemical energy.

5. Redox in Electrochemistry Electrochemistry is the study of the red ox processes by which chemical energy is converted to electrical energy and vice versa. RED OX Reduction: Gaining electron Oxidation: losing electron Recall: Because any loss of electrons by one substance must be accompanied by a gain in electrons by something else,  oxidation  and  reduction always occur together.

6. Redox in Electrochemistry

7. Electrochemical cells Apparatus that uses a red ox reaction to produce electric energy Oxidation : Zn(s) → Zn 2+ ( aq ) + 2e Reduction : Cu 2+ ( aq ) + 2e - → Cu Oxidation half-reaction: electron lost Reduction half-reaction: electrons gained Bulb doesn’t light up What is missing here?

8. Redox in Electrochemistry What do you think would happen if you separated the oxidation half-reaction from the reduction half-reaction? Can redox occur? 1) There is no way for zinc atoms to transfer electrons to copper(II) ion. 2) A positive charge builds up in one solution and a negative charge builds up in the other. The buildup of positive zinc ions on the left prohibits the oxidation of zinc atoms. On the other side, the buildup of negative ions prohibits the reduction of copper ions.

9. Redox in Electrochemistry

10. Salt-Bridge A salt bridge is a pathway to allow the passage of ions from one side to another, so that ions do not build up around the electrodes. Oxidation : Zn(s) → Zn 2+ ( aq ) + 2e Reduction : Cu 2+ ( aq ) + 2e - → Cu

11. Salt-Bridge the purpose of a salt bridge is not to move electrons from the electrolyte; rather it's to maintain charge  balance   because the electrons are moving from one-half cell to the other. Salt bridge prevents the diffusion or mechanical flow of solution from one-half cell to another. It maintains electrical neutrality within the internal circuit. If no salt bridge were present, the solution in one half cell would accumulate negative charge and the solution in the other half cell would accumulate positive charge as the reaction proceeded, quickly preventing further reaction, and hence production of electricity.

12. Electrochemical Cell

13. Starter: Voltaic Cells Main Idea : In voltaic cells, oxidation takes place at the anode, yielding electrons that flow to the cathode, where reduction occurs. can you clap with one hand? You cannot, you need two Similarly, in Voltic cells there are two half-cells , and both are required to produce energy .

14. Invention of voltaic cell Did you know, the anatomy of a frog has led to this revolutionary invention!!

15. Voltaic Cells

16. Voltaic cell

17. Voltaic cells Building the simplest Voltaic Cell Let’s try to build a voltaic cell. We place a copper plate and a zinc plate in solution of sodium chloride

18. Chemistry of voltaic cells Half-cells: The cell in which either oxidation or reduction takes place Zn (s) / Zn 2+ ( aq ) Oxidation Cu (s) / Cu 2+ ( aq ) Reduction The electrode where oxidation takes place is called the anode. ( An-Ox ) The electrode where reduction takes place is called the cathode ( red-cat ) Anode Cathode Electron flow from anode to cathode and current flow from cathode to anode

19. Voltaic cells and energy The roller coaster at the top of the track has high potential energy relative to track below because of the difference in height. Similarly, an electrochemical cell has potential energy to produce a current because there is a difference in the ability of the electrodes to move electrons from the anode to the cathode .

20. Calculating Electrochemical Cell Potentials The tendency of a substance to gain electrons is its reduction potential The reduction potential of an electrode cannot be determined directly , because the reduction half-reaction must be coupled with an oxidation half-reaction. Chemists decided to measure the reduction potential of all electrodes against one electrode:  the standard hydrogen electrode (SHE) Electric charge can flow between two points only when a difference in electrical potential energy exists between the two points.

21. Calculating Electrochemical Cell Potentials The measure of the ability of a species to gain or lose electrons Electrode Potential (E). The measure of the ability of a species to gain or lose electrons at its standard state (1M concentration, 1atm pressure and and temperature 25°C)is called Standard Electrode Potential (E°).

22. Standard hydrogen electrode [SHE] SHE can act as an oxidation half-reaction or a reduction half reaction, depending on the half-cell to which it is connected NOTE: Electrode potential of SHE is always Zero.

23. Standard hydrogen electrode [SHE] Salt bridge 1M acid solution , H + ( aq ) H 2 (g) bubbles Platinum electrode H 2 (g) (at 1 atm) Oxidation : H 2 (g) → 2 H + ( aq ) + 2e Reduction : 2 H + ( aq ) + 2e - → H 2 (g)

24. Electromotive Force (EMF):

25. Electromotive Force (EMF): Electrons generated at the anode, the site of oxidation, are thought to be pushed or driven toward the cathode by the electromotive force (EMF). This force is due to the difference in electric potential energy between the two electrodes and is referred to as the cell potential.

26. What is a Volt? A Volt is a unit used to measure cell potential. The electric potential difference of a voltaic cell is an indication of the energy that is available to move electrons from the anode to the cathode.

27. Voltaic cell Anode Zinc electrode Copper electrode AnOx Anode Oxidation Negative charge (-) Cathode Positive charge (+) RedCat Reduction at cathode Summary- Concept map standard potential of a voltaic cell is the difference between the standard reduction potentials of the half-cell reactions Oxidation half-cell Reduction half-cell

28. Cu 2+ ( aq ) + 2e - → Cu E (v)=+0.3419 2 H + ( aq ) + 2e - → H 2 (g) E (v)=+0.000V Half- Cell Potentials

29. Half- Cell Potentials

30. Determining Electrochemical Cell Potentials E = Overall standard cell potential   E   E  

31. Activity- Label the parts

32. Activity- Label the parts

33. Cu 2+ ( aq ) + 2e - → Cu E (v)=+0.3419 2 H + ( aq ) + 2e - → H 2 (g) E (v)=+0.000V Half- Cell Potentials

34. Determining Electrochemical Cell Potentials E = Overall standard cell potential   E   E   Recall

35. Electrochemical Cell Potentials Cu 2+ ( aq ) + 2e - → Cu (s) (reduction half-cell reaction) H 2 (g) + Cu 2+ ( aq ) → 2 H + ( aq ) + Cu (s) H 2 (g) | H + ( aq ) || Cu 2+ ( aq ) | Cu(s) E = +0.3419 V   This reaction can be written in a form called Cell notation H 2 (g) → 2 H + ( aq ) + 2e - (oxidation half-cell reaction) Reactant Product Oxidation half-cell Reduction half-cell Reactant Product

36. 2 H + ( aq ) ) + Zn (s) → Zn 2+ ( aq ) + H 2 (g) Zn (s) | Zn 2+ ( aq ) || H 2 (g) | 2 H + ( aq ) E = -0.762 V   This reaction can be written in a form called Cell notation 2 H + ( aq ) + 2e - → H 2 (g) (reduction half-cell reaction) Reactant Product Oxidation half-cell Reduction half-cell Reactant Product Zn (s) → Zn 2+ ( aq ) + 2e - (oxidation half-cell reaction) Electrochemical Cell Potentials

37. The Notation for Voltaic Cells

38. Determining Electrochemical Cell Potentials E = E E   E = +0.3419 − (-0.762V) = +1.104 V   0.000 V Standard hydrogen electrode +0.342 V Cu 2+ | Cu electrode -0.762 V Zn 2+ | Zn electrode Zn(s) | Zn 2+ ( aq ) || Cu 2+ ( aq ) | Cu(s) E = 1.04 V   E = E E   Note: Low reduction potential value for oxidation and high reduction potential value for Reduction

39. Cell potential in class Example-1 Analyze the Problem Known Standard reduction potentials for the half-cells E = E E   Unknown overall cell reaction = ? E Cell = ? cell notation = ? The following reduction half-reactions represent the half-cells of a voltaic cell. I 2 (s) + 2e – → 2I – ( aq ) and Fe 2+ ( aq ) + 2e – → Fe(s) Determine the overall cell reaction and the standard cell potential. Describe the cell using cell notation I 2 (s) + 2e – →2I – (aq) (reduction half-cell reaction) Fe(s) +2e – → Fe 2+ (aq) (oxidation half-cell reaction) E = +0.536 V   E = -0.447 V   Solve the Problem E = E E   E = +0.536 − (-0.447) = +0.983V  

40. Cell potential in class Example-1 Rewrite the iron half-reaction in the correct direction. I 2 (s) + 2e – →2I – (aq) (reduction half-cell reaction) Fe(s) → Fe 2+ (aq)+2e – (oxidation half-cell reaction) I 2 (s) + Fe(s) → 2I – (aq) + Fe 2+ (aq) Fe | Fe 2+ || I 2 | I – Cell notation Evaluate the answer E = +0.536 − (-0.447) = +0.983V E positive so the reaction is spontaneous   Solve the Problem Salt bridge

41. Cell potential Practice problems

42. Cell potential Practice problems Answers 1. 2. 3. 4.

43. Using Standard Reduction Potentials Another important use of standard reduction potential is to determine if a proposed reaction under standard conditions will be spontaneous . Electrons in a voltaic cell always flow from the half-cell with the lower standard reduction potential to the half-cell with the higher reduction potential, giving a positive cell voltage. How can standard reduction potentials indicate spontaneity?

44. Using Standard Reduction Potentials To predict whether any proposed redox reaction will occur spontaneously : Step 1 : write the process in the form of half-reactions and look up the reduction potential of each Step 2 : Use the values to calculate the potential of a voltaic cell operating with these two half-cell reactions. Step 3 : Evaluate if the reaction spontaneous or nonspontaneous  if the calculated potential is positive, the reaction is spontaneous.  If the value is negative, the reaction is not spontaneous.

45. Using Standard Reduction Potentials To predict whether any proposed redox reaction will occur spontaneously : Step 1 : write the process in the form of half-reactions and look up the reduction potential of each Step 2 : Use the values to calculate the potential of a voltaic cell operating with these two half-cell reactions. Step 3 : Evaluate if the reaction spontaneous or nonspontaneous  if the calculated potential is positive, the reaction is spontaneous.  If the value is negative, the reaction is not spontaneous. The reverse of a nonspontaneous reaction will occur because it will have a positive cell voltage, which means that the reverse reaction is spontaneous.

46. Using Standard Reduction Potentials – Practice problems

47. Using Standard Reduction Potentials – Practice problems Answers 9.

48. Voltaic cell Anode Zinc electrode Copper electrode AnOx Anode Oxidation Negative charge (-) Cathode Positive charge (+) RedCat Reduction at cathode Summary- Concept map standard potential of a voltaic cell is the difference between the standard reduction potentials of the half-cell reactions Oxidation half-cell Reduction half-cell

49. Section Summary- Concept map-2 H 2 (g) 1 atm, Pt Negative oxidation Positive reduction Standard potential is measured by connecting with SHE Half Cell Either oxidation or reduction E (v)=+0.00V Cathode reduction Standard hydrogen electrode (SHE) Anode- Oxidation STD conditions 1 atm 298k 1mol/dm -3