A Molecular Approach, 4e - Notes for Chapter (19).doc

Chapter 19. Electrochemistry Chapter 19. Electrochemistry Chapter 19. Electrochemistry Student Objectives 19.1 Pulling the Plug on the Power Grid Know that electrical power comes from a power grid, a distribution system for electricity. Know that fuel cell technology can replace the power source for an automobile or a home. Know that the oxidation and reduction reactions for the combination of hydrogen and oxygen to make water can power a fuel cell. 19.2 Balancing Oxidation–Reduction Equations Define oxidation and reduction. Balance aqueous redox reactions in acidic solution by using the half-reaction method. Balance aqueous redox reactions in basic solution by using the half-reaction method. 19.3 Voltaic (or Galvanic) Cells: Generating Electricity from Spontaneous Chemical Reactions Understand that the components in redox reactions can be separated into two compartments or cells and that the energy released by these reactions can be used to do electrical work. Define electrochemical cell, voltaic (galvanic) cell, electrolytic cell, half-cell, and electrode. Understand the rate of flowing electrons as current, and know the unit for current. Define potential difference and know its unit. Define electromotive force, cell potential, and standard cell potential. Know and understand that electrochemical reactions are a combination of half-cell reactions, oxidation occurring at the anode and reduction occurring at the cathode. Know and understand the components of an electrochemical cell, including the half-cells, electrodes, and salt bridge. Understand the direction of electron flow and the potential of a voltaic cell. Use cell notation to represent electrochemical cells. 19.4 Standard Electrode Potentials Know that the standard hydrogen electrode has an electrochemical potential of 0.00 V. Understand standard potentials and use them to generate potentials for electrochemical cells. Calculate standard potentials for electrochemical cells from standard electrode potentials of the half reactions. Calculate standard potentials and predict the direction of spontaneity for redox reactions. Use standard potentials to predict whether metals will dissolve in acid. 19.5 Cell Potential, Free Energy, and the Equilibrium Constant Calculate and interconvert between the standard change in free energy, G°, the standard potential for an electrochemical cell, E°, and the equilibrium constant, K. 19.6 Cell Potential and Concentration Use the Nernst equation to calculate the cell potential under nonstandard conditions. Define concentration cell and understand how a concentration cell works. Know and understand nerve cells as a biological example of a concentration cell. 19.7 Batteries: Using Chemistry to Generate Electricity Know some characteristics of dry-cell, lead-storage, nickel–cadmium, nickel–metal hydride, and lithium ion batteries. Know the primary difference between batteries and fuel cells. Know and understand how the hydrogen–oxygen fuel-cell and the fuel-cell breathalyzer work. 19.8 Electrolysis: Driving Nonspontaneous Chemical Reactions with Electricity Know and understand how an electrolytic cell works. Predict the products of some common electrolysis reactions. Calculate and interconvert between current, time, and amount for electrolysis reactions. 19.9 Corrosion: Undesirable Redox Reactions Know that corrosion generally refers to the oxidation reactions of metals. Know the steps by which iron rusts. Know some strategies for preventing corrosion. Section Summaries Lecture Outline Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples Teaching Tips Suggestions and Examples Misconceptions and Pitfalls Lecture Outline Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples 19.1 Pulling the Plug on the Power Grid Fuel cells definition uses reaction Intro figure: illustration of a smart phone powered by a fuel cell unnumbered figure: photo of a fuel cell capable of providing power for a house unnumbered figure: photo of a Toyota Mirai, a hydrogen-powered car 19.2 Balancing Oxidation–Reduction Equations Redox reactions oxidation: loss of electrons, increase in oxidation number reduction: gain of electrons, decrease in oxidation number Balancing redox reactions oxidation numbers separation of half-reactions mass balance charge balance electron balance recombination of half-reactions Examples 19.1 and 19.2 Half-Reaction Method of Balancing Aqueous Redox Equations in Acidic Solution Example 19.3 Balancing Redox Reactions Occurring in Basic Solution 19.3 Voltaic (or Galvanic) Cells: Generating Electricity from Spontaneous Chemical Reactions Properties of voltaic cells electrical current: ampere (A) 1 A = 1 C/s potential difference: volt (V) 1 V = 1 J/C Components of voltaic cells half-cells electrodes anode: oxidation cathode: reduction inert electrodes salt bridge Cell notation Figure 19.1 A Spontaneous Oxidation–Reduction Reaction Figure 19.2 A Voltaic Cell Figure 19.3 An Analogy for Electrical Current Figure 19.4 Inert Platinum Electrode Teaching Tips Suggestions and Examples Misconceptions and Pitfalls 19.1 Pulling the Plug on the Power Grid This example of a fuel cell is based on the reaction of hydrogen with oxygen. This reaction can be demonstrated as an explosion; the same reaction in a fuel cell releases the same amount of energy but in a controlled way. Batteries are ubiquitous devices and should motivate interest in this subject. Ask for applications of batteries in everyone’s daily lives. The combustion of hydrogen with oxygen occurs rapidly and often explosively. The hydrogen-oxygen fuel cell controls the kinetics but harnesses the same amount of energy. 19.2 Balancing Oxidation–Reduction Equations Identifying the individual half-reactions is generally straightforward. Particularly in redox reactions, both mass and charge must be balanced. There are other slightly different work-flows for balancing redox reactions. Working through a few examples in class helps students get the hang of balancing these reactions. Half-reactions that are not balanced initially with respect to mass (e.g., oxygen on one side but not the other side) are initially confusing to some students. Students may have difficulty counting the electrons in the half-reactions. 19.3 Voltaic (or Galvanic) Cells: Generating Electricity from Spontaneous Chemical Reactions The voltage or potential between the two half-cells drives the current. Demonstrations of redox reactions and voltaic cells are simple and quite satisfying to the students. Conceptual Connection 19.1 Voltaic Cells Electrochemical cell notation summarizes the critical components of an electrochemical cell in a compact form. An oxidation or reduction cannot occur in the absence of its counterpart. It can appear that they are independent since each half-cell reaction appears to occur in an individual compartment. The presence of inert electrodes in cell notation is initially confusing to some students. Lecture Outline Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples 19.4 Standard Electrode Potentials Standard hydrogen electrode reaction half-cell potential Standard electrode potentials measured versus reference like SHE table of values sign and value relative oxidizing strength predicting spontaneous direction metals dissolving in acid Figure 19.5 An Analogy for Electrode Potential Figure 19.6 The Standard Hydrogen Electrode Figure 19.7 Measuring Electrode Potential unnumbered figure: potential energy of anode and cathode for a spontaneous reaction unnumbered figure: potential energy of anode and cathode for a nonspontaneous reaction Table 19.1 Standard Electrode Potentials at 25 °C Example 19.4 Calculating Standard Potentials for Electrochemical Cells from Standard Electrode Potentials of the Half-Reactions Figure 19.8 Mn/Ni2+ Electrochemical Cell Example 19.5 Predicting Spontaneous Redox Reactions and Sketching Electrochemical Cells Figure 19.9 Mg/Fe2+ Electrochemical Cell Figure 19.10 Fe/Pb2+ Electrochemical Cell unnumbered figure: photo of reaction of Zn with HCl(aq) 19.5 Cell Potential, Free Energy, and the Equilibrium Constant Free energy change, G° Faraday’s constant, F = 96,485 C/mol e? G° = nFE° Equilibrium constant Example 19.6 Relating G° and Example 19.7 Relating and K Teaching Tips Suggestions and Examples Misconceptions and Pitfalls 19.4 Standard Electrode Potentials The electrode potential of SHE has been set at zero. Other potentials are measured relatively. Conceptual Connection 19.2 Standard Electrode Potentials The standard potential for an electrochemical cell determines the direction of spontaneity. Conceptual Connection 19.3 Selective Oxidation Conceptual Connection 19.4 Metals Dissolving in Acids Half-reactions with negative electrode potentials can still be spontaneous depending on the potential of the other half-reaction. Stronger oxidizing agents have higher electrode potentials; stronger reducing agents have more negative electrode potentials. Students sometimes forget that a strong oxidizing agent itself is reduced. When balancing redox reactions, one must often multiply a half-reaction by a factor such as 2 or 3 in order to balance the electrons. Students are tempted to multiply the standard half-cell potential by that same factor. 19.5 Cell Potential, Free Energy, and the Equilibrium Constant The standard cell potential, the standard free energy change, and the equilibrium constant can be used to predict the spontaneity of a reaction under standard conditions. Conceptual Connection 19.5 Periodic Trends and the Direction of Spontaneity for Redox Reactions Conceptual Connection 19.6 Relating K, , and Students often try to memorize all these equations and their variations, but they all can be derived easily using relationships the students already know from thermodynamics. Lecture Outline Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples 19.6 Cell Potential and Concentration Nonstandard conditions reaction quotient, Q Concentration cells identical electrodes, different concentrations concentration equilibration Human nerve cells ion channels in membranes sodium–potassium pump membrane potential Figure 19.11 Cell Potential and Concentration Example 19.8 Calculating Ecell Under Nonstandard Conditions Figure 19.12 Cu/Cu2+ Concentration Cell Chemistry and Medicine: Concentration Cells in Human Nerve Cells Figure 19.13 Concentration Changes in Nerve Cells Figure 19.14 Potential Changes across the Nerve Cell 19.7 Batteries: Using Chemistry to Generate Electricity Battery types dry cell standard: graphite and zinc electrodes; moist MnO2 and NH4Cl alkaline: graphite and zinc electrodes; moist MnO2 and KOH lead-acid storage rechargeable 6 cells for total of 12 V other rechargeable nickel–cadmium nickel–metal hydride lithium ion Fuel cell definition example: hydrogen–oxygen cell example: fuel-cell breathalyzer Figure 19.15 Dry-Cell Batteries Figure 19.16 Lead–Acid Storage Battery unnumbered figure: photo of rechargeable batteries and charger Table 19.2 Energy Density and Overcharge Tolerance of Several Rechargeable Batteries Figure 19.17 Lithium Ion Battery Figure 19.18 Hydrogen–Oxygen Fuel Cell Chemistry in Your Day: The Fuel-Cell Breathalyzer Teaching Tips Suggestions and Examples Misconceptions and Pitfalls 19.6 Cell Potential and Concentration The potential of a cell can be determined from the standard cell potential and the concentration ratios at nonstandard conditions, i.e. the Q expression. Concentration cells continue to operate until the concentrations become equal. This is the same concept as thermal equilibration and heat flow. A nerve cell is an excellent example of a concentration cell with biological relevance. Conceptual Connection 19.7 Relating Q, K, Ecell, and The Nernst equation is easily derived using the equation for G from thermodynamics. An electrochemical cell can be established using two identical electrodes; if the concentrations are different, the cell potential will not be zero. 19.7 Batteries: Using Chemistry to Generate Electricity Batteries are necessary for keeping electronics operating, so there may be considerable interest. A type of battery comes in a specific voltage; the size determines the amount of current available. Larger potentials represent multiple cells connected in series. Fuel cells also generate a specific voltage. Some devices are capable of high throughput of the chemical reactants and can generate considerable power. Lecture Outline Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples 19.8 Electrolysis: Driving Nonspontaneous Chemical Reactions with Electricity Reactivity spontaneous: voltaic cell nonspontaneous: electrolytic cell Design Examples electrolysis of water silver plating Predicting products pure molten salts mixtures of cations and anions aqueous solutions aqueous NaCl and overvoltage Stoichiometry Figure 19.20 Electrolysis of Water Figure 19.21 Silver Plating Figure 19.22 Voltaic versus Electrolytic Cells Figure 19.23 Electrolysis of Molten NaCl unnumbered figure: photo of battery-powered electrolysis of water Figure 19.24 Electrolysis of Aqueous NaI Figure 19.25 Electrolysis of Aqueous NaCl: The Effect of Overvoltage Example 19.9 Predicting the Products of Electrolysis Reactions Figure 19.26 Electrolytic Cell for Copper Plating Example 19.10 Stoichiometry of Electrolysis 19.9 Corrosion: Undesirable Redox Reactions Corrosion definition oxidation of metals reduction of oxygen examples of rust Preventing corrosion paint and covering sacrificial electrodes galvanizing unnumbered figure: photo of rusted ship unnumbered figure: photo of aluminum pipe Figure 19.27 Corrosion of Iron: Rusting unnumbered figure: photo of rust unnumbered figure: illustration of sacrificial anode unnumbered figure: photo of galvanized nails Teaching Tips Suggestions and Examples Misconceptions and Pitfalls 19.8 Electrolysis: Driving Nonspontaneous Chemical Reactions with Electricity Electrolytic cells use electrical input from an external source to drive a nonspontaneous electrochemical reaction. The applied potential determines what kinds of reactions can occur; the current and time determine the moles of material converted. Students sometimes have trouble predicting the products of electrolysis reactions in aqueous solution because they neglect the possibility of decomposing the water itself. Students also do not include the proper number of moles of electrons from the balanced equation when performing the dimensional analyses. 19.9 Corrosion: Undesirable Redox Reactions Corrosion typically occurs when metals are converted to their metal oxides. Rust, a metal oxide of iron, does not have the desired structural properties that iron and steel have. Corrosion is prevented by physically blocking the oxidation reaction or by providing a sacrificial electrode. Conceptual Connection 19.8 Sacrificial Electrodes Additional Problem for Calculating Standard Potentials for Electrochemical Cells from Standard Electrode Potentials of the Half-Reactions (Example 19.4) Use tabulated standard half-cell potentials to calculate the standard cell potential for the reaction in an electrochemical cell at 25 °C: Zn2+(aq) + H2O2(aq) Zn(s) + O2(g) + 2 H+(aq) Separate into half-cells Begin by separating the reaction into oxidation and reduction half-reactions. [In this case, the zinc cations form zinc metal, a reduction reaction.] Oxidation: H2O2(aq) + 2 e? O2(g) + 2 H+(aq) Reduction: Zn2+(aq) Zn(s) + 2 e? Assign Potentials Look up the standard half-cell potentials for each half-reaction. For the oxidation, remember to change the sign of to obtain . Add the half-cell reactions together. Ox: 2 e? + H2O2(aq) O2(g) + 2 H+(aq) Red: Zn2+(aq) Zn(s) + 2 e? _________________________________________ Zn2+(aq) + H2O2(aq) Zn(s) + O2(g) + 2 H+(aq) = = 0.70 V = 0.76 V Solve Add the half-cell potentials together to obtain the overall standard potential. = + = 0.70 V + (0.76 V) = 1.46 V Additional Problem for Relating and K (Example 19.7) Use the tabulated half-cell potentials to calculate K for the oxidation of nickel by chlorine: Cl2(g) + Ni(s) 2 Cl?(aq) + Ni2+(aq) Sort You are given the redox reaction and asked to find K. Given Cl2(g) + Ni(s) 2 Cl?(aq) + Ni2+(aq) Find K Strategize Use the tabulated values of the half-cell potentials to calculate . Then use Equation 19.6 to calculate K from . Conceptual Plan , = + K Solve Break the reaction up into oxidation and reduction half-reactions and find the standard half-cell potentials for each. Sum the values to find . Compute K from . The value of n (the number of moles of electrons) corresponds to the number of electrons that are canceled in the half reactions. Solution Ox: Ni(s) 2 e + Ni2+(aq) Red: Cl2(g) + 2 e 2 Cl(aq) _________________________________ Cl2(g) + Ni(s) 2 Cl(aq) + Ni2+(aq) = = +0.23 V = 1.36 V = + 0 = +0.23 V + 1.36 = +1.59 V Check The reaction is spontaneous and strongly favors the products. Chlorine is a powerful oxidizing agent. Additional Problem for Calculating Ecell under Nonstandard Conditions (Example 19.8) An electrochemical cell is based on two half-reactions: Oxidation: Fe(s) Fe2+(aq, 0.010 M) + 2 e? Reduction: Br2(l) + 2 e? 2 Br?(aq, 1.0 M) Compute the cell potential. Sort You are given the half-reactions for a redox reaction and the concentrations of the aqueous reactants and products. You are asked to find the cell potential. Given [Fe2+] = 0.010 M; [Br-] = 1.0 M Find Ecell Strategize Use the tabulated values of the half-cell potentials to calculate . Then use Equation 19.9 to calculate Ecell. Conceptual Plan , , [Fe2+], [Br-] Ecell Solve Write the oxidation and reduction half-reactions and sum them. Find the standard half-cell potentials of each and sum them to find . Compute Ecell from . The value of n is the number of moles of electrons. Determine Q based on the overall balanced equation and the given concentrations of reactants and products. Solution Ox: Fe(s) 2 e + Fe2+(aq) Red: Br2(l) + 2 e 2 Br(aq) _____________________________________ Br2(l) + Fe(s) 2 Br(aq) + Fe2+(aq) = = +0.45 V = 1.09 V = + = +0.45 V + 1.09 = +1.54 V Check The reaction is spontaneous. Bromine is a powerful oxidizing agent. Additional Problem for Stoichiometry of Electrolysis (Example 19.10) Aluminum can be plated out of a solution containing Al3+ according to the half-reaction: Al3+(aq) + 3 e Al(s) What mass of aluminum (in grams) will be plated by the flow of 50 A of current 10 hours? Sort You are given the half-reaction for the plating of aluminum, which shows the stoichiometry between the moles of aluminum and the moles of electrons. You are also given the current and time. You are asked to find the mass of gold that will be deposited in that time. Given 3 mol e = 1 mol Al 50 A 10 hours Find g Al Strategize You need to find the amount of aluminum, which is related stoichiometrically to the number of electrons that have flowed through the cell. Begin with time in hours and convert to seconds. Given the current and time, find the number of coulombs. Then use Faraday’s constant to calculate the number of moles of electrons and the stoichiometry of reaction to find the number of moles of gold. Finally, use the molar mass of aluminum to convert to the mass of aluminum. Conceptual Plan hr s C mol e mol Al g Al Solve Follow the conceptual plan, cancelling units to arrive at the mass of aluminum. Solution Check The units (g Al) are correct. The magnitude of the answer seems to make sense. (This amount is about 14 drink cans.) It takes a considerable amount of electricity to make the aluminum. In fact, aluminum is made electrolytically and is energy-intensive and expensive for this reason. This is why there is incentive to recycle aluminum cans. 258 Copyright © 2017 by Education, Inc. 257 Copyright © 2017 by Education, Inc.