Chapter 13 Electrochemistry PDF

Larry Brown Tom Holme Chapter 13 Electrochemistry Jacqueline Bennett SUNY Oneonta www.cengage.com/chemistry/brown Chapter Objectives Describe at least three types of corrosion and identify chemical reactions responsible for co...

Larry Brown Tom Holme Chapter 13 Electrochemistry Jacqueline Bennett SUNY Oneonta www.cengage.com/chemistry/brown Chapter Objectives Describe at least three types of corrosion and identify chemical reactions responsible for corrosion. Define oxidation and reduction. Write and balance half-reactions for simple redox processes. Describe the differences between galvanic and electrolytic cells. Use standard reduction potentials to calculate cell potentials under both standard and nonstandard conditions. 2 Chapter Objectives Use standard reduction potentials to predict the spontaneous direction of a redox reaction. Calculate the amount of metal plated, the amount of current needed, or the time required for an electrolysis process. Distinguish between primary and secondary batteries. Describe the chemistry of some common battery types and explain why each type of battery is suitable for a particular application. Describe at least three common techniques for preventing corrosion. 3 Corrosion Corrosion is the degradation of metals by chemical reactions with the environment. Uniform corrosion occurs evenly over a large portion of the surface area of a metal. Galvanic corrosion occurs when two different metals contact each other in the presence of an appropriate electrolyte. Crevice corrosion occurs when two pieces of metal touch each other, leaving a small gap or crevice between the metals. 4 Corrosion Different metals corrode differently. Aluminum has a greater tendency to corrode than iron, but corrosion of aluminum is not problematic compared to iron. The aluminum oxide corrosion product forms a protective layer on the surface of aluminum metal. The iron oxide corrosion product flakes off the surface of iron, exposing fresh iron to corrosion. 5 Corrosion Corrosion occurs in a variety of forms. The chain shows uniform corrosion. The grill cover shows crevice corrosion where the handle is attached. 6 Oxidation-Reduction Reactions and Galvanic Cells Special conditions must be present before iron reacts with oxygen to form iron(III) oxide. Rust formation is a slow process, so the basics of electrochemistry must be investigated using more easily observed reactions. Reactions that transfer electrons between reactants are known as oxidation-reduction or redox reactions. Oxidation is the loss of electrons from some chemical species. Reduction is the gain of electrons to some chemical species. 7 Oxidation-Reduction and Half-Reactions For an oxidation-reduction reaction to occur, one reactant must be oxidized and one reactant must be reduced. Oxidation cannot occur without reduction. When copper wire is placed in a silver nitrate solution, a redox reaction occurs. A reaction is observed to occur because the solution changes color and crystals form on the surface of the copper wire. 8 Oxidation-Reduction and Half-Reactions When a clean copper wire is placed into a colorless solution of silver nitrate, it is quickly apparent that a chemical reaction occurs: silver metal forms on the copper wire the solution turns blue. 9 Oxidation-Reduction and Half-Reactions The solution’s blue color is indicative of Cu2+ ions in solution. Cu2+ is formed when a copper atom loses two electrons. The copper metal is oxidized. The crystals forming on the surface of the copper wire are silver metal. Silver is formed when a silver cation gains an electron. The silver cation is reduced. 10 Oxidation-Reduction and Half-Reactions For the reaction between silver cation and copper metal, two half-reactions are written. One for the oxidation of copper and one for the reduction of silver. Neither half-reaction can occur without the other. The half-reactions as written indicate that Ag+ only accepts one electron whereas Cu loses two electrons. The electron transfer must balance, so the reduction half- reaction is multiplied by 2. 11 Oxidation-Reduction and Half-Reactions Add the two half-reactions together, the electrons cancel out, leaving the net ionic equation for the redox reaction. 12 Oxidation-Reduction and Half-Reactions The species undergoing oxidation is referred to as a reducing agent. The Cu was oxidized and is the reducing agent. The Cu facilitated the reduction of Ag+ by losing electrons. The species undergoing reduction is referred to as an oxidizing agent. The Ag+ was reduced and is the oxidizing agent. The Ag+ facilitated the oxidation of Cu by gaining electrons. 13 Building a Galvanic Cell A galvanic cell is any electrochemical cell in which a spontaneous chemical reaction can be used to generate an electric current. The name electrochemistry comes from the observation of electric currents in galvanic cells. To harness electricity from a galvanic cell, each half-reaction is prepared separately in half-cells. Cu metal immersed in Cu2+ solution is one half-cell. Ag metal immersed in Ag+ solution is the second half-cell. 14 Building a Galvanic Cell A salt bridge is crucial to a galvanic cell. The salt bridge allows ions to flow between each half-cell, completing the circuit. 15 Building a Galvanic Cell Current flows by the migration of ions in solution. To transfer current between the half-cells, a salt bridge is used. The salt bridge contains a strong electrolyte that allows either cations or anions to migrate into the solution where they are needed to maintain charge neutrality. A metal wire cannot transport ions and cannot be used. 16 Building a Galvanic Cell For a salt bridge composed of NH4Cl: NH4+ will flow into the Ag+ beaker to offset the removal of Ag+ from solution. Cl– will flow into the Cu2+ beaker to offset the production of Cu2+ in solution. The circuit is completed by connecting wires to each metal strip. A voltage potential of 0.46 V will be measured for the described cell. 17 Terminology for Galvanic Cells Electrodes are the electrically conducting sites at which either oxidation or reduction occurs. The electrode where oxidation occurs is the anode. The electrode where reduction occurs is the cathode. Cell notation - a shorthand notation for the specific chemistry of an electrochemical cell. Cell notation lists the metals and ions involved in the reaction. A vertical line, |, denotes a phase boundary. A double vertical line, ||, denotes a salt bridge. The anode is written on the left, the cathode on the right. 18 Terminology for Galvanic Cells General form of cell notation For the previous example of copper and silver The electrolyte concentration is also given. An electrochemical cell is at its standard state when the electrolyte concentrations are 1 M. For half-cells that generate or consume a gas, a partial pressure of 1 atm is required for the standard state. 19 Atomic Perspective on Galvanic Cells Before half-cells are connected by a salt bridge, a small build up of charge occurs for each half-cell at the interface between the electrode and the electrolyte. At the anode, some oxidation occurs and cations dissolve into solution, leaving a negative charge on the anode. At the cathode, some reduction occurs and cations are removed from solution, leaving a positive charge on the cathode. 20 Atomic Perspective on Galvanic Cells An equilibrium can be described for each half-cell, the half- reaction equilibrium. Not an oxidation-reduction equilibrium. The build up of charge on the electrode means there is potential for electrical work. This potential is the cell potential, or electromotive force (EMF). EMF is related to the maximum work obtainable from an electrochemical cell. w = qE max q is the charge, E is the cell potential. 21 Atomic Perspective on Galvanic Cells Without a salt bridge to close the circuit, local charges build up around both electrodes. Neither electrode reaction can proceed to any significant extent, so no cell voltage can be measured. 22 Galvanic Corrosion and Uniform Corrosion Metals in contact with a solution establish an oxidation half- reaction equilibrium. If the solution contains a substance that can undergo reduction, a redox reaction may occur. For two metals in contact, such as a tin-plated steel can, exposure to air and moisture results in rapid corrosion. The half-reaction equilibrium for the tin facilitates the process by which iron is oxidized. This is an example of galvanic corrosion. 23 Galvanic Corrosion and Uniform Corrosion A “tin can” is usually tin-plated steel. If the tin coating is scratched to expose the underlying steel, iron in the steel will corrode rapidly. 24 Galvanic Corrosion and Uniform Corrosion Metal not in contact with another metal can also corrode. A nonmetal is involved in the second half-cell. For the corrosion of iron, iron is one half-cell and oxygen dissolved in water is the second half-cell. The electrode for the oxygen half-cell is the iron itself. Dissolved salts facilitate the corrosion reaction. This is an example of uniform corrosion. 25 Galvanic Corrosion and Uniform Corrosion Iron is oxidized, and oxygen from the air is reduced. Water is needed for ion mobility between the anodic and cathodic regions, and the presence of ionic salts speeds the reaction considerably. 26 Cell Potentials The relative corrosivities of various plated steels can be expressed as cell potential. A voltmeter measures the size of the electrical potential and also its polarity - the locations of the negative charge (negative pole) and the positive charge (positive pole). An electric potential has a fixed polarity and voltage. Reversing the poles of a battery with respect to a voltmeter changes the sign on the measured voltage but does not influence the electrochemical reaction in the battery. 27 Measuring Cell Potential When a voltmeter is connected to the previously described copper/silver cell, a potential of 0.462 V is measured. Connecting the copper half-cell to a reducing iron(III)/iron(II) half-cell, a cell potential of 0.434 V is measured. Connecting the iron(III)/iron(II) half-cell to the silver half-cell results in a cell potential of 0.028. For the three cell potentials measured, the fact that 0.462 V = 0.434 V + 0.028 V suggests two things: The behavior of cell potentials is akin to state functions. If a specific standard electrode is chosen, comparison to all other electrodes will result in a practical system for determining cell potential. 28 Measuring Cell Potential Measurement of standard cell voltages for various combinations of the same half-reactions suggests that a characteristic potential can be associated with a particular half-reaction. 29 Measuring Cell Potential The standard hydrogen electrode (SHE) is the choice for the standard component in cell potential measurements. The cell is constructed of a platinum wire or foil as the electrode. The electrode is immersed in a 1 M HCl solution through which H2 gas with a pressure of 1 atm is bubbled. The SHE has been chosen as the reference point for the scale of standard reduction potentials, and assigned a potential of exactly zero volts. 30 Measuring Cell Potential For the Standard Hydrogen Electrode The half-reaction for the SHE is: The half-cell notation is: Pt(s) | H2(g, 1 atm) | H+ (1 M). The half-cell is assigned a potential of exactly zero volts. The cell potential is attributed to the other half-reaction. 31 Measuring Cell Potential For some galvanic cells, the SHE acts as the anode and for other galvanic cells, the SHE acts as the cathode. The anode is the site of oxidation, releasing electrons and creating a negatively charged electrode. If the anode is connected to the positive terminal of the voltmeter, a negative potential is measured. The cathode is the side of reduction, consuming electrons and creating a positively charged electrode. If the cathode is connected to the positive terminal of the voltmeter, a positive potential is measured. 32 Measuring Cell Potential When the SHE is always connected to the positive terminal, the sign of the potential tells us the direction of the redox reaction. When the potential is negative, the SHE is the anode, and H2 is oxidized to H+(aq). When the potential is positive, the SHE is the cathode, and H+(aq) is reduced to H2. 33 Measuring Cell Potential Just like a commercial battery, a galvanic cell has a fixed polarity. Electrons flow through the external circuit from the anode to the cathode. Reversing the leads of the voltmeter changes the sign of the reading, but does not influence the flow of current. 34 Standard Reduction Potentials To compare the oxidation-reduction trends of species used in electrochemistry, all half-cell potentials are written as reductions. A table of standard reduction potentials lists the potential of any half-reaction when connected to a SHE. All materials are 1 M for aqueous species and 1 atm partial pressure for gases. 35 Standard Reduction Potentials Standard reduction potentials for several half-reactions involved in the cells discussed in the text. 36 Standard Reduction Potentials Although the half-reactions are listed as reductions in the table, one half-reaction in any electrochemical cell must be an oxidation and, therefore, reversed from what appears in the table. The cell potential sign must be changed when writing the half-reaction as an oxidation. Some half-reactions have positive potentials, whereas others have negative potentials. 37 Standard Reduction Potentials All potentials are measured with a SHE connected to the positive terminal. If the voltage is positive, the SHE is the anode, the oxidation site. A positive standard reduction potential means the half- reaction proceeds as written (reduction occurs). If the voltage is negative, the SHE is the cathode, the reduction site. A negative standard reduction potential means the half- reaction proceeds as an oxidation. 38 Standard Reduction Potentials The tendency for the chemicals involved in a half-reaction to be an oxidation or reduction depends on the value of the reduction potential. A large, positive value for the standard reduction potential implies the substance is reduced readily and a good oxidizing agent. A large, negative value for the standard reduction potential implies the substance is oxidized readily and a good reducing agent. 39 Standard Reduction Potentials For a galvanic cell, the half-reaction with the more positive reduction potential will be the cathode. The half-reaction with the more negative reduction potential will be the anode. The standard reduction potential for any pair of half-reactions, E°cell, is calculated from the standard reduction potentials for the cathode and anode. E° is the standard reduction potential for the cathode and red E°ox is the standard reduction potential for the anode. 40 Standard Reduction Potentials For standard reduction potentials arranged horizontally, the anode and cathode for a galvanic cell can easily be determined. The reduction potential furthest to the left is the anode. 41 Example Problem 13.1 Using standard reduction potentials, identify the anode and the cathode and determine the cell potential for a galvanic cell composed of copper and iron. Assume standard conditions. Confirm that the potential of the following galvanic cell is 0.462 V: 42 Nonstandard Conditions The cell potential at nonstandard conditions is calculated using the Nernst equation. Q is the reaction quotient, F is the Faraday constant, and n is the number of electrons transferred in the reaction. F = 96,485 J V-1 mol-1 or 96,485 C mol-1 43 Example Problem 13.2 Assume that you have a cell that has an iron(II) concentration of 0.015 M and an H+ concentration of 1.0 × 10-3 M. The cell temperature is 38°C, and the pressure of hydrogen gas is maintained at 0.04 atm. What would the cell potential be under these conditions? 44 Cell Potentials and Free Energy Corrosion is a spontaneous process and has a negative Gibbs free energy change. The Gibbs free energy change for an electrochemical reaction can be calculated from the standard reduction potential. n is the number of electrons transferred and F is Faraday’s constant. The minus sign in required because a galvanic cell has a positive cell potential, spontaneously generates electrical work, and thus must have a negative ΔG value. 45 Example Problem 13.3 Suppose that we wish to study the possible galvanic corrosion between zinc and chromium, so we set up the following cell: What is the chemical reaction that takes place and what is the standard free energy change for that reaction? 46 Equilibrium Constants The cell potential can be used to calculate the equilibrium constant for an electrochemical reaction. n is the number of electrons transferred, R is the universal gas law constant, and F is Faraday’s constant. 47 Equilibrium Constants The relationship between the cell potential and the equilibrium constant can be re-written in terms of the common (base 10) log. The equation can be simplified for reaction carried out at standard temperature, 25°C (298 K). 48 Equilibrium Constants The equilibrium constant increases as the cell potential and number of electrons transferred increases. The different lines correspond to reactions involving 1, 2, or 3 electrons transferred. 49 Batteries A battery is a cell or series of cells that generate an electrical current. Batteries are the means by which we harness the electrical work of a galvanic cell and use it productively. 50 Primary Cells Single-use batteries that cannot be recharged are primary cells, or primary batteries. The most prevalent type of primary cell is the alkaline battery. An alkaline battery has a zinc electrode at which oxidation occurs. The cathode is derived from manganese(IV) oxide. 51 Primary Cells The chemistry of an alkaline dry cell battery. The net reaction is shown above. The alkaline battery is termed a dry cell because the KOH electrolyte is in the form of a paste or gel. 52 Primary Cells Lithium batteries are small and long lasting. They are used for medical devices like pacemakers. Lithium is the anode. Manganese(IV) oxide is the cathode as in the alkaline battery, but in this case the MnO2 reacts with lithium ions. The lithium battery provides a stable current and electrical potential for long periods. 53 Primary Cells Zinc-air batteries are also primary cells. Zinc is the anode. Oxygen reacts at the cathode. In a zinc-air battery, one of the reactants is oxygen from the surrounding air. As a result, these batteries can offer a very attractive energy density. 54 Secondary Cells Rechargeable batteries are secondary cell or secondary batteries. Nickel-metal-hydride batteries are an example of secondary cells. The anode reaction is The complex cathode reaction can be represented as 55 Secondary Cells Nickel-metal-hydride batteries have become popular as rechargeable cells. 56 Secondary Cells The lead-acid storage battery found in cars is a secondary cell. The anode for a lead-acid battery is lead metal. The cathode for a lead-acid battery is lead oxide. The lead-acid storage battery consists of Pb anodes alternating with PbO2 cathodes, all immersed in sulfuric acid. 57 Secondary Cells Comparison of battery characteristics 58 Fuel Cells A fuel cell is a voltaic cell in which the reactants can be supplied continuously and the products of the cell reaction are continuously removed. Most common type is based on the reaction of hydrogen and oxygen to produce water. Oxygen is reduced at the cathode. Hydrogen is oxidized at the anode. 59 Limitations of Batteries Corrosion is a major cause for the loss of performance in batteries. Protective plating of materials used in batteries is an attempt to limit the performance-diminishing effects of corrosion on batteries. 60 Electrolysis Electrolysis is the process of passing an electric current through an ionic solution or molten salt to produce a chemical reaction. Electrolytic cells are divided into two categories based on the nature of the electrodes used. Passive electrolysis: the electrodes are chemically inert materials that simply provide a path for electrons. Active electrolysis: the electrodes are part of the electrolytic reaction. 61 Electrolysis and Polarity Electrolysis changes the polarity of the electrodes in a system. For reduction, electrons are forced to the cathode. The cathode becomes the negative electrode. For oxidation, electrons are pulled from the anode. The anode becomes the positive electrode. In electrolysis, an external source of current drives a redox reaction that would otherwise not be spontaneous. The flow of ions through the solution completes the circuit. 62 Passive Electrolysis in Refining Aluminum Electrolysis provides the means to overcome the nonspontaneous reaction to separate aluminum from its oxide. The Hall-Heroult refining process uses carbon electrodes as inert sites for passive electrolysis. The Hall-Heroult process involves the electrolytic refining of aluminum from Al2O3 to produce aluminum metal and oxygen gas. 63 Active Electrolysis and Electroplating The process of depositing a thin coat of metal on another metal by using electrolysis is electroplating. In some cases, the thin coating is cosmetic, or to provide some vital functionality for the coated piece, such as corrosion resistance or desirable conductive properties. Silver is plated onto electrical devices because silver is a good conductor and resistant to corrosion. The solution from which silver is plated contains CN–(aq) ions, which form a complex with Ag+. The need for uniform coatings makes this an important step. 64 Active Electrolysis and Electroplating The object being electroplated is the cathode. Anode Cathode Opposite reactions at the anode and cathode are common for electroplating operations. Silver is transferred from the anode to the cathode, coating the cathode in a thin layer of silver. The zero cell potential is not critical since an external current drives electrolysis. 65 Active Electrolysis and Electroplating Barrel plating is often used to apply coatings to small parts. 66 Electrolysis and Stoichiometry For electroplating, it can be vitally important to use carefully controlled amounts of materials. Controlling the flow of electrons (current) in an electroplating operation provides a method to accurately limit the amount of material deposited. Electroplating is often used to prevent galvanic corrosion in an electrical apparatus in places where different metals come into contact with one another. 67 Current and Charge When current is measured in an electric circuit, the observation is the flow of charge for a period of time. The unit of current, the ampere (A), is defined as one coulomb per second: 1 A = 1 C s-1. If a known current flows through a circuit for a known time, the charge can be easily calculated. 68 Current and Charge Using Faraday’s constant, F = 96,485 C mol-1 and the calculated charge, the number of moles of electrons that pass through the circuit can be calculated. If the number of electrons required to reduce each metal cation is known, the number of moles of material plated can be calculated. Electricity use is often measured in terms of power. The SI unit for power is the watt (1 watt = 1 J s-1) Electrical utilities normally determine consumption in kilowatt- hours, kWh (1 kWh = 3.60 × 106 J) 69 Example Problem 13.4 In a process called flash electroplating, a current of 2.50 x 103 A passes through an electrolytic cell for 5.00 minutes. How many moles of electrons are driven through the cell? 70 Example Problem 13.5 Suppose that a batch of parts is plated with copper in an electrolytic bath running at 0.15 V and 15.0 A for exactly 2 hours. What is the energy cost of this process if the electric utility charges the company $0.0500 per kWh? 71 Calculations Using Masses of Substances in Electrolysis A knowledge of current, how long the current flows, stoichiometry, and the number of electrons required to reduce a metal cation are used to answer the following questions. How much material is plated given a specific current for an allotted time or electrical energy expenditure? How long must a given current to pass through the cell to yield a desired mass of plated material? 72 Example Problem 13.6 An electrolysis cell that deposits gold (from Au+(aq)) operates for 15.0 minutes at a current of 2.30 A. What mass of gold is deposited? 73 Example Problem 13.7 Suppose that you have a part that requires tin coating. You’ve calculated that you need to deposit 3.60 g of tin to achieve an adequate coating. If your electrolysis cell (using Sn2+) runs at 2.00 A, how long must you operate the cell to obtain the desired coating? 74 Batteries in Engineering Design Lithium-ion batteries have recently come under scrutiny because of a high-profile fire during the first commercial flight of the Boeing Dreamliner 787 75 Batteries in Engineering Design What factors led to the Dreamliner’s Problems? The Dreamliner requires higher levels of electrical power than conventional aircraft due in large part to replacing the traditional hydraulic system for controlling the flight of the plane with an all-electronic alternative. Increased fuel efficiency was a major goal, so it was important to minimize weight. Lithium-ion batteries have high energy capacity and low weight. 76 Batteries in Engineering Design Schematic of lithium-ion battery 77 Batteries in Engineering Design The anode is a form of graphite into which lithium atoms have been incorporated, or “intercalated.” The cathode is cobalt oxide that also has lithium incorporated into it. Because the electrodes both consist of solids with lithium embedded in them, it is difficult to write a simple chemical equation for the cell reaction.The chemistry is summarized reasonably by the following reactions: 78 Batteries in Engineering Design Two reasons energy density is high: Li+/Li has one of the largest standard reduction potentials Both lithium and carbon are relatively light High temperatures cause lithium-ion batteries to degrade fairly quickly, so engineering designs that use them generally must account for this. 79