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Chapter 20: Electrochemistry

Chapter 20: Electrochemistry. Section 1: Introduction to Electrochemistry. Electrochemistry: deals with electricity-related applications of oxidation-reduction reactions.

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Chapter 20: Electrochemistry

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  1. Chapter 20: Electrochemistry

  2. Section 1:Introduction to Electrochemistry • Electrochemistry: deals with electricity-related applications of oxidation-reduction reactions. • Oxidation-Reduction Reactions involve a transfer of electrons. If both substances are in contact with each other, there is also a transfer of energy as heat.

  3. Energy as heat is given off when electrons are transferred directly from Zn atoms to Cu2+ ions. This causes the temperature of the aqueous CuSO4 solution to rise.

  4. Electrochemical Cells • If the substances are separated by a porous barrier or salt bridge, the transfer of energy as heat becomes a transfer of electrical energy. • The barrier prevents the metal atoms from getting through but the ions can move through the barrier, which prevents a charge from building up on the electrodes. • Allows for the movement of charge through ions.

  5. 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 the movement of electrons is balanced by the movement of ions in solution.

  6. An electrode is a conductor used to establish electrical contact with a nonmetallic part of a circuits, such as an electrolyte. • In the diagram, the Zn and Cu strips are electrodes. • A single electrode immersed a solution of its own ions is a half-cell.

  7. The Half-Cells • Half-Cells can be represented by half-reactions. • The Half-cell with the Zn electrode in ZnSO4 solution has the half-reaction: Zn(s)  Zn2+(aq)+ 2e- • Oxidation is occurring so this electrode is called the anode. • The Half-cell with the Cu electrode in CuSO4 solution has the half-reaction: Cu2+(aq)+ 2e-  Cu(s) • Reduction is occurring so this electrode is called the cathode.

  8. In Ch 19 we learned that both oxidation and reduction must occur in an electrochemical reaction. The two half-cells together make an electrochemical cell.

  9. The Complete Cell • Notation: Anode | Anode || Cathode | Cathode Electrode Solution Solution Electrode For the Zinc and Copper cell, the cell notation is: Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s) The overall electrochemical reaction is found by adding the anode half reaction to the cathode half reaction: 1st + 3rd 2nd + 4th Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s) *If the change in charge is not equal, you may need to multiply the half-reaction to balance out.

  10. Did You Know? • In 1836, John Daniell developed the Daniell Cell, an electrochemical cell with Zn and Cu. This was the first battery to produce a constant electrical current over a long period of time, and therefore was important for the operation of the electric telegraph.

  11. Homework None! 

  12. Section 2:Voltaic Cells • Voltaic Cells use spontaneous oxidation-reduction reactions to convert chemical energy into electrical energy. • The most common application of voltaic cells is batteries. • The three most common types of dry cells are the zinc-carbon battery, the alkaline battery, and the mercury battery.

  13. Zinc-Carbon Dry Cells • These consist of a zinc container (anode) filled with a moist paste of MnO2, NH4Cl, and ZnCl2. When the external circuit is closed, zinc atoms are oxidized at the negative electrode. Zn(s)  Zn2+(aq) + 2e- • Electrons move across the circuit and reenter the cell through the carbon rod, or cathode. 2MnO2(s) + H2O(l) + 2e-  Mn2O3(s) + 2OH-(aq)

  14. Alkaline Batteries • These do not have a carbon rod which allows them to be smaller. This cell uses a paste of Zn metal and KOH instead of a solid metal anode. Zn(s) + 2OH-(aq)  Zn(OH)2(s) + 2e- • The half-reaction at the cathode is the same as the zinc-carbon dry cell. 2MnO2(s) + H2O(l) + 2e-  Mn2O3(s) + 2OH-(aq)

  15. Mercury Batteries • Mercury Batteries have the same anode half-reaction as in the alkaline dry cell. Zn(s) + 2OH-(aq)  Zn(OH)2(s) + 2e- • However, the cathode half-reaction is different HgO(s) + H2O(l) + 2e-  Hg(l) + 2OH-

  16. Fuel Cells • These are voltaic cells in which the reactants are being continuously supplied and the products are being continuously removed. So in principle, a fuel cell could keep changing chemical energy into electrical energy forever. • The reaction below shows the type of fuel cell used in the United States Space Program. Cathode: O2(g) + 2H2O(l) + 4e-  4OH- Anode: 2H2(g) + 4OH-(aq)  4e- + 4H2O(l) Net Reaction: 2H2 + O2  2H2O

  17. Fuel cells are very efficient and have very low emissions.

  18. Corrosion and Its Prevention • The metal most commonly affected by corrosion is iron. • The formation of rust forms by the following reaction: 4Fe(s) + 3O2(g) + xH2O(l)  2Fe2O3●xH2O(s) • The value of x will vary depending on the amount of water present and will affect the color of rust formed.

  19. Preventing Corrosion • The most common way is to coat steel with zinc in a process called galvanizing. • Zinc is more easily oxidized than iron so it will react first, protecting the iron. This is called cathodic protection. • Ex: The Alaskan Oil Pipeline

  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. • This “pull” is called the electric potential and is expressed in volts (V). • Current is the movement of the electrons and is expressed in units of amperes, or amps (A). • Electrons flow from higher electric potential to lower electric potential.

  21. Electrode Potentials • Standard Electrode Potential: the potential of a half-cell under standard conditions measure relative to the standard hydrogen electrode. Chart on page 664 lists the SEP values. • These are always written as reduction half-reactions. When they are changed to oxidation half-reactions, the sign is reversed. Zn2+ + 2e-  Zn E0 = -0.76 V Zn  Zn2+ + 2e- E0 = +0.76 V

  22. Calculating Cell Potential • Standard electrode potentials can be used to predict if a redox reaction will occur spontaneously. If E0cell is positive, it is spontaneous. E0cell = E0cathode - E0anode • The half-reaction that has the more negative E0 value will be the anode. • You do not need to switch the sign of the anode when doing these calculations because the subtraction in the formula takes care of that.

  23. Example 1 • 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. • Look up E0 for each half-reaction on pg 664. Fe3+(aq) + 3e-  Fe(s) E0 = -0.04 V Ag+(aq) + e-  Ag(s) E0 = +0.80 V • Determine the cathode and anode. Anode is Fe (more negative) and Cathode is Ag

  24. Example 1 Cont. 3. Determine the overall cell reaction. Electrons need to match so, multiple the Ag half-reaction by 3 and reverse the Fe reaction since we determined it is the anode and needs to be oxidation. Fe(s)  Fe3+(aq) + 3e- E0 = -0.04 V 3Ag+(aq) + 3e-  3Ag(s) E0 = +0.80 V *we do not multiply the E0 value by 3 Fe(s) + 3Ag+(aq) 3Ag(s) + Fe3+(aq) *electrons cancel out because they are both 3

  25. Example 1 Cont. • Calculate the cell potential E0cell = E0cathode - E0anode E0cell = +0.80 V – (-0.04 V) = +0.84 V Check your answer: The calculated value for E0cell is positive so that means it is spontaneous and a voltaic cell as the question stated.

  26. Example 2: • Determine the overall electrochemical equation and E0 value for H+/H2 and Fe2+/Fe3+. 2H++ 2e-  H2 E0 = +0.00 V Fe3++ e-  Fe2+E0= +0.77 V Anode is H2(more negative) and Cathode is Fe2+ Multiply Fe half-reaction by 2 and reverse the H reaction. 2Fe3+ + 2e-  2Fe2+E0= +0.77 V H2  2H+ + 2e-E0= +0.00 V H2 + 2Fe3+ 2H+ + 2Fe2+ E0 = 0.77 V – 0.00 V = 0.77 V

  27. Homework • None!

  28. Section 3: Electrolytic Cells • If electrical energy is required to produce a redox reaction and bring about a chemical change in an electrochemical cell, it is an electrolytic cell. • Example: In a voltaic cell consisting of zinc and copper, the electrons move from zinc to copper. When you attach a current, the electrons move from copper to zinc.

  29. Important Differences • The anode and cathode of an electrolytic cell are connected to a battery while a voltaic cell serves as a source of electrical energy. • Electrolytic cells have nonspontaneous redox reactions occurring which is why they require an outside electrical energy source. Voltaic cells have spontaneous redox reactions occurring which produce electricity.

  30. Electroplating • An electrolytic process in which a metal ion is reduced and solid metal is deposited on a surface is called electroplating. • Pennies consist of a zinc core (97.5%) electroplated with a layer of copper. • Video of Electroplating

  31. Rechargeable Cells • A rechargeable cell combines the redox chemistry of both voltaic and electrolytic cells. • When it operates as a battery it is like a voltaic cell, converting chemical energy into electrical energy. • When the battery is being recharged, it is like an electrolytic cell, converting electrical energy into chemical energy.

  32. Electrolysis • Electroplating and recharging a battery are both examples of electrolysis. • In electrolysis, electrical energy is used to force a nonspontaneous chemical reaction to occur. • Electrolysis is used to purify many metals from the ores in which they are found chemically combined in the earth’s crust.

  33. Electrolysis of Water • Hydrogen gas and Oxygen gas will spontaneously combine to form water. • To break apart water into H2 and O2 requires energy and is called the electrolysis of water.

  34. Aluminum Production by Electrolysis • Aluminum is the most abundant metal in the earth’s crust, however it is found as an oxide in an ore called bauxite. • The Hall-Héroult process made the production of aluminum economically reasonable. However, it requires a lot of energy—nearly 5% of the national total. • Recycling aluminum saves almost 95% of the cost of production and is the most economically worthwhile recycling program ever developed.

  35. Homework • Ch 20.3 pg 671 #1, 3 and pg 673 #21, 23, 24

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