Chapter 20. Electrochemistry. Chapter 20 Problems. Electrochemistry 7/3/2012. Problems 15, 17, 19, 23, 27, 29, 33, 39, 59

Transcription

1 Chemistry, The Central Science, 11th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten Chapter 20 John D. Bookstaver St. Charles Community College Cottleville, MO Chapter 20 Problems Problems 15, 17, 19, 23, 27, 29, 33, 39, 59 In the broadest sense, electrochemistry is the study of chemical reactions that produce electrical effects and of the chemical phenomena that are caused by the action of currents or voltages. 1

2 Oxidation-Reduction Reactions Basic Terms Oxidation - loss of electrons Reduction - gain of electrons Redox reaction oxidizing agent - substance that causes oxidation by being reduced reducing agent - substance that causes reduction by being oxidized Electrochemical Reactions In electrochemical reactions, electrons are transferred from one species to another. Oxidation Numbers (State) In order to keep track of what loses electrons and what gains them, we assign oxidation numbers. 2

3 Oxidation and Reduction A species is oxidized when it loses electrons. Here, zinc loses two electrons to go from neutral zinc metal to the Zn 2+ ion. Oxidation and Reduction A species is reduced when it gains electrons. Here, each of the H + gains an electron, and they combine to form H 2. Oxidation and Reduction What is reduced is the oxidizing agent. H + oxidizes Zn by taking electrons from it. What is oxidized is the reducing agent. Zn reduces H + by giving it electrons. 3

4 Oxidation and Reduction Oxidation O.S. of some element increases in the reaction. Electrons are on the right of the equation Reduction O.S. of some element decreases in the reaction. Electrons are on the left of the equation. Assigning Oxidation Numbers 1. Elements in their elemental form have an oxidation number of The oxidation number of a monatomic ion is the same as its charge. Assigning Oxidation Numbers 3. Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions. Oxygen has an oxidation number of 2, except in the peroxide ion, which has an oxidation number of 1. Hydrogen is 1 when bonded to a metal and +1 when bonded to a nonmetal. 4

5 Assigning Oxidation Numbers 3. Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions. Fluorine always has an oxidation number of 1. The other halogens have an oxidation number of 1 when they are negative; they can have positive oxidation numbers, however, most notably in oxyanions. Assigning Oxidation Numbers 4. The sum of the oxidation numbers in a neutral compound is The sum of the oxidation numbers in a polyatomic ion is the charge on the ion. Oxidation State What is the oxidation state of S in H 2 SO 4? H => +1 O => -2 neutral compound, thus sum equals zero 4O => 4*-2 = -8 2H => 2*+1 = +2 0 = +2 + (x) + (-8) x = +6 5

6 Oxidation State What is the oxidation state of Cl in HClO 4? H => +1 O => -2 neutral compound, thus sum equals zero 4O => 4*-2 = -8 H => 1*+1 = +1 0 = +1 + (y) + (-8) y = +7 Oxidation and Reduction Half- Reactions A reaction represented by two halfreactions. Oxidation: Zn(s) Zn 2+ (aq) + 2 e - Reduction: Cu 2+ (aq) + 2 e- Cu(s) Overall: Cu 2+ (aq) + Zn(s) Cu(s) + Zn 2+ (aq) Balancing Oxidation-Reduction Equations Few can be balanced by inspection. Systematic approach required. The Half-Reaction (Ion-Electron) Method 6

7 Balancing Oxidation-Reduction Equations Perhaps the easiest way to balance the equation of an oxidation-reduction reaction is via the half-reaction method. Balancing Oxidation-Reduction Equations This involves treating (on paper only) the oxidation and reduction as two separate processes, balancing these half reactions, and then combining them to attain the balanced equation for the overall reaction. The Half-Reaction Method 1. Assign oxidation numbers to determine what is oxidized and what is reduced. 2. Write the oxidation and reduction halfreactions. 7

8 The Half-Reaction Method 3. Balance each half-reaction. a. Balance elements other than H and O. b. Balance O by adding H 2 O. c. Balance H by adding H +. d. Balance charge by adding electrons. 4. Multiply the half-reactions by integers so that the electrons gained and lost are the same. The Half-Reaction Method 5. Add the half-reactions, subtracting things that appear on both sides. 6. Make sure the equation is balanced according to mass. 7. Make sure the equation is balanced according to charge. The Half-Reaction Method Consider the reaction between MnO 4 and C 2 O 4 2 : MnO 4 (aq) + C 2 O 4 2 (aq) Mn 2+ (aq) + CO 2 (aq) 8

9 The Half-Reaction Method First, we assign oxidation numbers MnO 4 + C 2 O 4 2- Mn 2+ + CO 2 Since the manganese goes from +7 to +2, it is reduced. Since the carbon goes from +3 to +4, it is oxidized. Oxidation Half-Reaction C 2 O 4 2 CO 2 To balance the carbon, we add a coefficient of 2: C 2 O CO 2 Oxidation Half-Reaction C 2 O CO 2 The oxygen is now balanced as well. To balance the charge, we must add 2 electrons to the right side. C 2 O CO e 9

10 Reduction Half-Reaction MnO 4 Mn 2+ The manganese is balanced; to balance the oxygen, we must add 4 waters to the right side. MnO 4 Mn H 2 O Reduction Half-Reaction MnO 4 Mn H 2 O To balance the hydrogen, we add 8 H + to the left side. 8 H + + MnO 4 Mn H 2 O Reduction Half-Reaction 8 H + + MnO 4 Mn H 2 O To balance the charge, we add 5 e to the left side. 5 e + 8 H + + MnO 4 Mn H 2 O 10

11 Combining the Half-Reactions Now we evaluate the two half-reactions together: C 2 O CO e 5 e + 8 H + + MnO 4 Mn H 2 O To attain the same number of electrons on each side, we will multiply the first reaction by 5 and the second by 2. Combining the Half-Reactions 5 C 2 O CO e 10 e + 16 H MnO 4 2 Mn H 2 O When we add these together, we get: 10 e + 16 H MnO C 2 O Mn H 2 O + 10 CO e Combining the Half-Reactions 10 e + 16 H MnO C 2 O Mn H 2 O + 10 CO e The only thing that appears on both sides are the electrons. Subtracting them, we are left with: 16 H MnO C 2 O Mn H 2 O + 10 CO 2 11

12 Balancing in Basic Solution If a reaction occurs in basic solution, one can balance it as if it occurred in acid. Once the equation is balanced, add OH to each side to neutralize the H + in the equation and create water in its place. If this produces water on both sides, you might have to subtract water from each side. Important Electrochemical Terms An electrochemical cell is a device that combines two half-cells with the appropriate connections between electrodes and solutions A voltaic cell is an electrochemical cell in which electric current is generated from a spontaneous redox reaction The anode is the electrode at which oxidation occurs and the cathode is the electrode at which reduction occurs The cell potential (E cell ) is the potential difference that propels electrons from the anode to the cathode EOS Voltaic Cells In spontaneous oxidation-reduction (redox) reactions, electrons are transferred and energy is released. 12

13 Cell Diagrams Conventions Place the anode on the left side of the diagram Place the cathode on the right side of the diagram Use a single vertical line ( ) to represent the boundary between different phases, such as between an electrode and a solution Use a double vertical line ( ) to represent a salt bridge or porous barrier separating two halfcells EOS An Example Cell Diagram EOS Terminology Galvanic (Voltaic) cells. Produce electricity as a result of spontaneous reactions. Electrolytic cells. Non-spontaneous chemical change driven by electricity. Couple, M M n+ A pair of species related by a change in number of e -. 13

14 Voltaic Cells harnessed chemical reaction which produces an electric current Terminology Zn(s) Zn 2+ (aq) Cu 2+ (aq) Cu(s) E cell = V Voltaic Cells We can use that energy to do work if we make the electrons flow through an external device. We call such a setup a voltaic cell. 14

15 Voltaic Cells A typical cell looks like this. The oxidation occurs at the anode. The reduction occurs at the cathode. Voltaic Cells Once even one electron flows from the anode to the cathode, the charges in each beaker would not be balanced and the flow of electrons would stop. Voltaic Cells Therefore, we use a salt bridge, usually a U-shaped tube that contains a salt solution, to keep the charges balanced. Cations move toward the cathode. Anions move toward the anode. 15

16 Voltaic Cells In the cell, then, electrons leave the anode and flow through the wire to the cathode. As the electrons leave the anode, the cations formed dissolve into the solution in the anode compartment. Voltaic Cells As the electrons reach the cathode, cations in the cathode are attracted to the now negative cathode. The electrons are taken by the cation, and the neutral metal is deposited on the cathode. Electromotive Force (emf) Water only spontaneously flows one way in a waterfall. Likewise, electrons only spontaneously flow one way in a redox reaction from higher to lower potential energy. 16

17 Electromotive Force (emf) The potential difference between the anode and cathode in a cell is called the electromotive force (emf). It is also called the cell potential and is designated E cell. Cell Potential Cell potential is measured in volts (V). 1 V = 1 J C Cell Potential (Ecell) The cell potential (Ecell) is the potential difference that propels electrons from the anode to the cathode E cell = E o cathode - E o anode 17

18 Standard Electrode Potentials Cell voltages, the potential differences between electrodes, are among the most precise scientific measurements. The potential of an individual electrode is difficult to establish. Arbitrary zero is chosen for the Standard Hydrogen Electrode (SHE) Standard Hydrogen Electrode Half-cell values are referenced to a standard hydrogen electrode (SHE). By definition, the reduction potential for hydrogen is 0 V: 2 H + (aq, 1M) + 2 e H 2 (g, 1 atm) Standard Hydrogen Electrode In the standard hydrogen electrode, hydrogen gas at exactly 1 bar pressure is bubbled over an inert platinum electrode and into an aqueous solution with the concentration adjusted so that the activity of H 3 O + is exactly equal to one EOS 18

19 Standard Hydrogen Electrode 2 H + (a = 1) + 2 e - H 2 (g, 1 bar) E = 0 V Pt H 2 (g, 1 bar) H + (a = 1) Standard Electrode Potential, E E defined by international agreement. The tendency for a reduction process to occur at an electrode. All ionic species present at a=1 (approximately 1 M). All gases are at 1 bar (approximately 1 atm). Where no metallic substance is indicated, the potential is established on an inert metallic electrode (ex. Pt). Standard Electrode Potentials A standard electrode potential, E o, is based on the tendency for reduction to occur at the electrode The cell voltage, called the standard cell potential (E o cell), is the difference between the standard potential of the cathode and that of the anode Voltaic cells can produce electrical work w = n F E cell Two illustrations Cu and Zn VideoClip EOS 19

20 Standard Cell Potential Pt H 2 (g, 1 bar) H + (a = 1) Cu 2+ (1 M) Cu(s) E cell = V E cell = E cathode - E anode E cell = E Cu 2+ /Cu - E H + /H V = E Cu 2+ /Cu - 0 V E Cu 2+ /Cu = V H 2 (g, 1 atm) + Cu 2+ (1 M) H + (1 M) + Cu(s) E cell = V Measuring Standard Reduction Potential anode cathode cathode anode Cu 2+ /Cu Electrode EOS 20

21 Zn 2+ /Zn Electrode EOS Observed Voltages EOS Standard Reduction Potentials Reduction potentials for many electrodes have been measured and tabulated. 21

22 Important Points About Electrode and Cell Potentials Electrode potentials and cell voltages are intensive properties independent of the amount of matter Cell voltages can be ascribed to oxidation reduction reactions without regard to voltaic cells E cell = E cathode E anode EOS Cell Potentials For the oxidation in this cell, E red = 0.76 V For the reduction, E red = V Cell Potentials E cell = E red (cathode) E red (anode) = V ( 0.76 V) = V 22

23 Oxidizing and Reducing Agents The strongest oxidizers have the most positive reduction potentials. The strongest reducers have the most negative reduction potentials. Oxidizing and Reducing Agents The greater the difference between the two, the greater the voltage of the cell. Criteria for Spontaneous Change If E cell is positive, the reaction in the forward direction (from left to right) is spontaneous If E cell is negative, the reaction is nonspontaneous If E cell = 0, the system is at equilibrium When a cell reaction is reversed, E cell and DG change signs EOS 23

24 Free Energy DG for a redox reaction can be found by using the equation DG = nfe where n is the number of moles of electrons transferred, and F is a constant, the Faraday. 1 F = 96,485 C/mol = 96,485 J/V-mol Free Energy Under standard conditions, DG = nfe Nernst Equation Remember that DG = DG + RT ln Q This means nfe = nfe + RT ln Q 24

25 Nernst Equation Dividing both sides by nf, we get the Nernst equation: E = E RT ln Q nf or, using base-10 logarithms, E = E RT nf log Q Nernst Equation At room temperature (298 K), RT F = V Thus the equation becomes E = E n log Q Summary of Important Relationships EOS 25

26 Concentration and Cell Voltage If the discussion is limited to 25 o C, the equation is E o o cell E cell RT lnq Ecell lnq nf n The Nernst equation relates a cell voltage for nonstandard conditions (E cell ) to a standard cell voltage, E o cell, and the concentrations of reactants and products The Nernst equation is useful for determining the concentration of a species in a voltaic cell through a measurement of E cell EOS A Sample Problem E o cell E cell lnq n EOS EXAMPLE Applying the Nernst Equation for Determining E cell. What is the value of E cell for the voltaic cell pictured below and diagrammed as follows? Pt Fe 2+ (0.10 M),Fe 3+ (0.20 M) Ag + (1.0 M) Ag(s) 26

27 EXAMPLE V E cell = E cell - log Q n V [Fe E cell = E cell - log 3+ ] n [Fe 2+ ] [Ag + ] E cell = V V = V Pt Fe 2+ (0.10 M),Fe 3+ (0.20 M) Ag + (1.0 M) Ag(s) Fe 2+ (aq) + Ag + (aq) Fe 3+ (aq) + Ag (s) Concentration Cells If the cell potential is determined solely by a difference in the concentration of solutes in equilibrium with identical electrodes, that cell is called a concentration cell EOS Concentration Cells Notice that the Nernst equation implies that a cell could be created that has the same substance at both electrodes. For such a cell, E cell would be 0, but Q would not. Therefore, as long as the concentrations are different, E will not be 0. 27

28 ph Measurement One useful concentration cell is the hydrogen cell it is directly related to ph The Nernst equation can be rewritten in terms of concentration cells using hydrogen electrodes E cell = (0.0592)(pH) EOS Applications of Oxidation-Reduction Reactions Batteries 28

29 Alkaline Batteries Hydrogen Fuel Cells Corrosion and 29

30 Corrosion Prevention 30