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Chapter 20: Electrochemistry

CHEMISTRY Matter and Change. Chapter 20: Electrochemistry. Table Of Contents. CHAPTER 20. Section 20.1 Voltaic Cells Section 20.2 Batteries Section 20.3 Electrolysis. Click a hyperlink to view the corresponding slides. Exit. Voltaic Cells. SECTION 20.1.

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Chapter 20: Electrochemistry

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  1. CHEMISTRY Matter and Change Chapter 20: Electrochemistry

  2. Table Of Contents CHAPTER20 Section 20.1 Voltaic Cells Section 20.2 Batteries Section 20.3 Electrolysis Click a hyperlink to view the corresponding slides. Exit

  3. Voltaic Cells SECTION20.1 • Describe a way to obtain electrical energy from a redox reaction. • Identify the parts of a voltaic cell, and explain how each part operates. • Calculate cell potentials, and determine the spontaneity of redox reactions. oxidation: the loss of electrons from the atoms of a substance; increases an atom’s oxidation number reduction: the gain of electrons by the atoms of a substance; decreases the atom’s oxidation number

  4. Voltaic Cells SECTION20.1 salt bridge electrochemical cell voltaic cell half-cell anode cathode reduction potential standard hydrogen electrode In voltaic cells, oxidation takes place at the anode, yielding electrons that flow to the cathode, where reduction occurs.

  5. Voltaic Cells SECTION20.1 Redox in Electrochemistry • Electrochemistry is the study of the redox processes by which chemical energy is converted to electrical energy and vice versa. • Redox reactions involve a transfer of electrons from the species that is oxidized to the species that is reduced.

  6. Voltaic Cells SECTION20.1 Redox in Electrochemistry (cont.)

  7. Voltaic Cells SECTION20.1 Redox in Electrochemistry (cont.) • 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. • An electrochemical cellis an apparatus that uses a redox reaction to produce electrical energy or uses electrical energy to cause a chemical reaction.

  8. Voltaic Cells SECTION20.1 Redox in Electrochemistry (cont.) • A voltaic cell is a type of electrochemical cell that converts chemical energy to electrical energy by a spontaneous redox reaction.

  9. Voltaic Cells SECTION20.1 Chemistry of Voltaic Cells • An electrochemical cell consists of two parts called half-cells, in which the separate oxidation and reduction reactions take place. • The electrode where oxidation takes place is called the anode. • The cathode is the electrode where reduction occurs.

  10. Voltaic Cells SECTION20.1 Chemistry of Voltaic Cells(cont.) • Electric potential energy is a measure of the amount of current that can be generated from a voltaic cell to do work. • Electric charge can flow between two points only when a difference in electric potential energy exists between the two points. • A volt is a unit used to measure cell potential—the force from the difference in electric potential energy between two electrodes.

  11. Voltaic Cells SECTION20.1 Calculating Electrochemical Cell Potentials • The tendency of a substance to gain electrons is its reduction potential. • When two half-reactions are coupled, the voltage generated corresponds to the difference in potential between the half-reactions.

  12. Voltaic Cells SECTION20.1 Calculating Electrochemical Cell Potentials (cont.) • The standard hydrogen electrode consists of a small sheet of platinum immersed in a hydrochloric acid solution thathas a hydrogen ion concentration of 1 M. Hydrogen gas (H2), at a pressure of 1 atm, is bubbled in and the temperature ismaintained at 25°C.

  13. Voltaic Cells SECTION20.1 Calculating Electrochemical Cell Potentials (cont.) • The standard hydrogen electrode is the standard against which all other reduction potentials are measured.

  14. Voltaic Cells SECTION20.1 Calculating Electrochemical Cell Potentials

  15. Voltaic Cells SECTION20.1 Calculating Electrochemical Cell Potentials

  16. Voltaic Cells SECTION20.1 Use Standard Reduction Potentials • Cell potentials can be used to determine if a proposed reaction under standard conditions will be spontaneous. • If the calculated potential is positive, the reaction is spontaneous. • If the calculated potential is negative, the reaction is not spontaneous.

  17. Section Check SECTION20.1 In electrochemistry, the site where oxidation occurs is called ____. A.electrode B.anode C.cathode D.ion

  18. Section Check SECTION20.1 The standard potential of a voltaic cell is the difference between the: A.electrode voltage B.standard reduction potential of the cell and hydrogen C.standard reduction potentials of the half- cell reactions D.half-cell reactions and the salt bridge

  19. Batteries SECTION20.2 • Describe the structure, composition, and operation of the typical carbon-zinc dry-cell battery. • Distinguish between primary and secondary batteries, and give two examples of each type. • Explain the structure and operation of the hydrogen-oxygen fuel cell. • Describe the process of corrosion of iron and methods to prevent corrosion. reversible reaction: a reaction that can take place in both the forward and reverse directions

  20. Batteries SECTION20.2 battery dry cell primary battery secondary battery fuel cell corrosion galvanization Batteries are voltaic cells that use spontaneous reactions to provide energy for a variety of purposes.

  21. Batteries SECTION20.2 Dry Cells • A battery is one or more voltaic cells in a single package that generates electric current.

  22. Batteries SECTION20.2 Dry Cells (cont.) • A dry-cellis an electrochemical cell in which the electrolyte is a moist paste. The paste in a zinc-carbon cell consists of zinc chloride, manganese(IV) oxide, ammonium chloride, and a small amount of water. • The anode is the zinc shell. • The cathode is a carbon rod, but reduction occurs in the paste.

  23. Batteries SECTION20.2 Dry Cells (cont.) • In the alkaline cell, zinc is in a powdered form and mixed with potassium hydroxide contained in a steel case. • Alkaline batteries are small and more useful in small devices.

  24. Batteries SECTION20.2 Dry Cells (cont.) • Silver batteries are similar to alkaline but smaller.

  25. Batteries SECTION20.2 Dry Cells (cont.) • Primary batteries produce electric energy by means of redox reaction that are not easily reversed. • Secondary batteriesdepend on reversible redox reactions and are rechargeable.

  26. Batteries SECTION20.2 Lead-Acid Storage Battery • Lead-acid storage batteries are common in automobiles. • The electrolyte solution is sulfuric acid, hence the name. • The anode consists of grids of porous lead. • The cathode consists of lead grids filled with lead(IV) oxide.

  27. Batteries SECTION20.2 Lithium Batteries • Lithium is the lightest known metal and has the lowest standard reduction potential of the metallic elements. • Lithium batteries can be either primary or secondary.

  28. Batteries SECTION20.2 Fuel Cells • A fuel cell is a voltaic cell in which the oxidation of a fuel is used to produce electric energy.

  29. Batteries SECTION20.2 Fuel Cells (cont.) • How a fuel cell works • Potassium hydroxide is often the electrolyte • The oxidation half reaction 2H2(g) + 4OH– → 4H2O + 4e– • The reduction half reaction O2(g) + 2H2O(l) + 4e– → 4OH–(aq) • When combined, the equation is the same as burning hydrogen in oxygen to form water.

  30. Batteries SECTION20.2 Corrosion • Corrosionis the loss of metal resulting from an oxidation-reduction reaction of the metal with substances in the environment. • Rusting begins in a chip or pit in the iron surface, which become the anode. Fe(s) → Fe2+(aq) + 2e– • Iron(II) becomes part of the water solution.

  31. Batteries SECTION20.2 Corrosion (cont.) • The cathode is usually at the edge of the water drop where water, iron, and air come into contact. • The reduction reaction is O2(g) + 4H+(aq) + 4e– → 2H2O(l). • Next the oxidation 4Fe2+(aq) + 2O2(g) + 2H2O(l) + 4e– → 2Fe2O3(s) + 4H+. • Rusting is slow, but salts speed the process.

  32. Batteries SECTION20.2 Corrosion (cont.)

  33. Batteries SECTION20.2 Corrosion (cont.) • Paint and other covers seal out moisture to prevent corrosion. • Blocks of metal that are more easily oxidized than steel, such as magnesium, aluminum, or titanium, are often attached to the hulls of ships—they corrode while the iron in the hull isprotected. They are called sacrificial anodes.

  34. Batteries SECTION20.2 Corrosion (cont.) • Galvanization is the process of coating iron with a layer or zinc. • Zinc oxidizes at the surface, creating a layer of metal-oxide that protects from further corrosion.

  35. Section Check SECTION20.2 Which type of battery has a reversible spontaneous reaction? A.alkaline battery B.secondary battery C.primary battery D.zinc-carbon battery

  36. Section Check SECTION20.2 Which is NOT a method of preventing corrosion? A.painting B.galvanization C.coating with electrolytes D.sacrificial anode

  37. Electrolysis SECTION20.3 • Describe how it is possible to reverse a spontaneous redox reaction in an electrochemical cell. redox reaction: an oxidation-reduction reaction electrolysis electrolytic cell • Compare the reactions involved in the electrolysis of molten sodium chloride with those in the electrolysis of brine. • Discuss the importance of electrolysis in the smelting and purification of metals. In electrolysis, a power source causes nonspontaneous reactions to occur in electrochemical cells.

  38. Electrolysis SECTION20.3 Reversing Redox Reactions • The use of electrical energy to bring about a chemical reaction is called electrolysis. • An electrochemical cell in which electrolysis occurs is called an electrolytic cell.

  39. Electrolysis SECTION20.3 Applications of Electrolysis • Electrolysis of water is one method of obtaining hydrogen gas for commercial use.

  40. Electrolysis SECTION20.3 Applications of Electrolysis (cont.) • Electrolysis can separate molten sodium chloride into sodium metal and chlorine gas in a chamber called a Down’s cell.

  41. Electrolysis SECTION20.3 Applications of Electrolysis (cont.) • In the decomposition of brine, an aqueous solution of NaCl, electrolysis is used to produce hydrogen gas, chlorine gas, and sodium hydroxide.

  42. Electrolysis SECTION20.3 Applications of Electrolysis (cont.) • The Hall-Héroult process requires large amounts of energy, which is the primary reason for recycling aluminum.

  43. Electrolysis SECTION20.3 Applications of Electrolysis (cont.) • Objects can be electroplated with a metal such as silver. • The cathode is the object to be electroplated, where reduction occurs. • The anode is a bar of silver, where silver is oxidized and silver ions are transferred to the cathode.

  44. Section Check SECTION20.3 Which of the following is NOT a product of electrolysis of brine? A.Chlorine gas B.Hydrogen gas C.Sodium hydroxide solution D.Sodium metal

  45. Section Check SECTION20.3 What is required to drive a nonspontaneous reaction in an electrolytic cell? A.electrodes B.additional ions C.an energy source D.an electrolyte

  46. Electrochemistry CHAPTER20 Resources Chemistry Online Study Guide Chapter Assessment Standardized Test Practice

  47. Voltaic Cells SECTION20.1 Study Guide Key Concepts • In a voltaic cell, oxidation and reduction take place at electrodes separated from each other. • The standard potential of a half-cell reaction is its voltage when paired with a standard hydrogen electrode under standard conditions. • The reduction potential of a half-cell is negative if it undergoes oxidation when connected to a standard hydrogen electrode. The reduction potential of a half-cell is positive if it undergoes reduction when connected to a standard hydrogen electrode.

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