Chapter 20 Electrochemistry (modified for our needs)
Chemistry, The Central Science, 10th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten Chapter 20 Electrochemistry Ch. 20 Electrochemistry Oxidation States Balancing Oxidation-Reduction Equations Voltaic cells Cell EMF
Spontaneity of redox reactions The effect of concentration on EMF Electrolysis Electrochemical Reactions In electrochemical reactions, electrons are transferred from one species to another. In order to keep track of what
loses electrons and what gains them, we assign oxidation numbers. Examples of reactions that are redox reactions (write equations)
A piece of solid bismuth is heated strongly in oxygen. A strip or copper metal is added to a concentrated solution of sulfuric acid. Magnesium turnings are added to a solution of iron (III) chloride. A stream of chlorine gas is passed through a solution of cold, dilute sodium hydroxide. A solution of tin ( II ) chloride is added to an acidified
solution of potassium permanganate A solution of potassium iodide is added to an acidified solution of potassium dichromate. 4 Hydrogen peroxide solution is added to a solution of iron (II) sulfate. Propanol is burned completely in air. A piece of lithium metal is dropped into a container of nitrogen gas. Chlorine gas is bubbled into a solution of potassium iodide.
Magnesium metal is burned in nitrogen gas. Lead foil is immersed in silver nitrate solution. Pellets of lead are dropped into hot sulfuric acid Powdered Iron is added to a solution of iron(III) sulfate. 5 Combination: Oxidizing agent of one element will react with the reducing agent of the same element to produce the free element. I- + IO3- + H+ I2 + H2O
Decomposition. a) peroxides to oxides b) Chlorates to chlorides c) Electrolysis into elements. d) carbonates to oxides Oxidation and Reduction A species is oxidized when it loses electrons. Here, zinc loses two electrons to go from neutral zinc metal to the Zn2+ 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 H2. 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. Assigning Oxidation Numbers 1. Elements in their elemental form have an oxidation number of 0. 2. The oxidation number of a monatomic ion is the same as its charge. 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 in which it has an oxidation number of 1. Hydrogen is 1 when bonded to a metal, +1 when bonded to a nonmetal. 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. 4. The sum of the oxidation numbers in a neutral compound is 0. 5. The sum of the oxidation numbers in a polyatomic ion is the charge on the ion. Balancing Oxidation-Reduction Equations
Perhaps the easiest way to balance the equation of an oxidation-reduction reaction is via the half-reaction method. 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. Half-Reaction Method 1. Assign oxidation
numbers to determine what is oxidized and what is reduced. 2. Write the oxidation and reduction halfreactions. 3. Balance each half-reaction. a. Balance elements other than H and O. b. Balance O by adding H2O.
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. 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. Half-Reaction Method Consider the reaction between MnO4 and C2O42 : MnO4(aq) + C2O42(aq) Mn2+(aq) + CO2(aq) Half-Reaction Method First, we assign oxidation numbers. +7
+3 +2 +4 MnO4 + C2O42- Mn2+ + CO2 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
C2O42 CO2 To balance the carbon, we add a coefficient of 2: C2O42 2 CO2 The oxygen is now balanced as well. To balance the charge, we must add 2 electrons to the right side. C2O42 2 CO2 + 2 e Reduction Half-Reaction MnO4 Mn2+ The manganese is balanced; to balance the
oxygen, we must add 4 waters to the right side. MnO4 Mn2+ + 4 H2O To balance the hydrogen, we add 8 H+ to the left side. 8 H+ + MnO4 Mn2+ + 4 H2O Reduction Half-Reaction 8 H+ + MnO4 Mn2+ + 4 H2O To balance the charge, we add 5 e to the left side. 5 e + 8 H+ + MnO4 Mn2+ + 4 H2O
Combining the Half-Reactions Now we evaluate the two half-reactions together: C2O42 2 CO2 + 2 e 5 e + 8 H+ + MnO4 Mn2+ + 4 H2O 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 C2O42 10 CO2 + 10 e 10 e + 16 H+ + 2 MnO4 2 Mn2+ + 8 H2O When we add these together, we get: 10 e + 16 H+ + 2 MnO4 + 5 C2O42 2 Mn2+ + 8 H2O + 10 CO2 +10 e Combining the Half-Reactions 10 e + 16 H+ + 2 MnO4 + 5 C2O42 2 Mn2+ + 8 H2O + 10 CO2 +10 e The only thing that appears on both sides are the electrons. Subtracting them, we are left with:
16 H+ + 2 MnO4 + 5 C2O42 2 Mn2+ + 8 H2O + 10 CO2 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.
(Practice Problems) Voltaic Cells In spontaneous oxidation-reduction (redox) reactions, electrons are transferred and energy is released. 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. See ANIMATIONS 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. 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. 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 Ecell. Cell potential is measured in volts (V). J 1V=1 C Standard Reduction Potentials Reduction
potentials for many electrodes have been measured and tabulated. Standard Hydrogen Electrode Their 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 H2 (g, 1 atm) Standard Cell Potentials The cell potential at standard conditions can be found through this equation: Ecell = E red (cathode) Ered (anode) Because cell potential is based on the
potential energy per unit of charge, it is an intensive property. Cell Potentials For the oxidation in this cell, Ered = 0.76 V For the reduction,
Ered = +0.34 V Ecell = Ered (cathode) Ered (anode) = +0.34 V (0.76 V) = +1.10 V 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. Free Energy G for a redox reaction can be found by using the equation G = 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 Under standard conditions,
G = nFE Nernst Equation Remember that G = G + RT ln Q This means nFE = nFE + RT ln Q Nernst Equation Dividing both sides by nF, we get the Nernst equation:
RT ln Q E = E nF or, using base-10 logarithms, 2.303 RT log Q E = E nF Nernst Equation
At room temperature (298 K), 2.303 RT F = 0.0592 V Thus (when T = 298 K) the equation becomes E = E 0.0592 n
log Q 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, Ecell would be 0, but Q would not.
Therefore, as long as the concentrations are different, E will not be 0. Applications of Oxidation-Reduction Reactions Batteries Hydrogen Fuel Cells Corrosion and
Corrosion Prevention Electrolysis (animations) Using electrical energy to drive a reaction in a non-spontaneous direction Electrolysis Using electrical energy to
drive a reaction in a nonspontaneous direction Used for electroplating, electrolysis of water, separation of a mixture of ions, etc. (Most negative reduction potential is easiest to plate out of solution.) Calculating plating
Have to count charge. Measure current I (in amperes) 1 amp = 1 coulomb of charge per second q=Ixt q/nF = moles of metal Mass of plated metal Faraday Constant (F) (96,480 C/mol e-) gives the amount of charge (in coulombs that exist in 1 mole of electrons passing through a
circuit. 1volt = 1joule/coulomb Calculating plating 1. 2. 3. 4. Current x time = charge Charge Faraday = mole of
eMol of e- to mole of element or compound Mole to grams of compound or the reverse these steps if you want to find the time to plate How many grams of copper are deposited on the cathode of an electrolytic cell if an electric current of 2.00A is run through a
solution of CuSO4 for a period of 20min? How many hours would it take to produce 75.0g of metallic chromium by the electrolytic reduction of Cr3+ with a current of 2.25 A? How many grams of copper are deposited on the cathode of an electrolytic cell if an electric current of 2.00A is run through a solution of CuSO4 for a period of 20min? Answer: Cu2+(aq) + 2e- Cu(s) 2.00A = 2.00C/s and 20min (60s/min) = 1200s
Coulombs of e- = (2.00C/s)(1200s) = 2400C mol e- = (2400C)(1mol/96,480C) = .025mol (.025mol e-)(1mol Cu/2mol e-) = .0125mol Cu g Cu = (.0125mol Cu)(63.55g/mol) = .79g How many hours would it take to produce 75.0g of metallic chromium by the electrolytic reduction of Cr3+ with a current of 2.25 A? Answer: 75.0g Cr/(52.0g/mol) = 1.44mol Cr mol e- = (1.44mol Cr)(3mol e-/1mol Cr) = 4.32mol eCoulombs = (4.32mol e-)(96,480C/mol) = 416,793.6 C
(4.17x105C) Seconds = (4.17x105C)/(2.25C/s) = 1.85x105 s Hours = (1.85x105 s)(1hr/3600 s) = 51.5 hours Problem A student places a copper electrode in a 1 M solution of CuSO4 and in another beaker places a silver electrode in a 1 M solution of AgNO3. A salt bridge composed of Na2SO4 connects the two beakers. The voltage measured across the electrodes is found to be + 0.42 volt.
(a) Draw a diagram of this cell. (b) Describe what is happening at the cathode (Include any equations that may be useful.) Problem A student places a copper electrode in a 1 M solution of CuSO4 and in another beaker places a silver electrode in a 1 M solution of AgNO3. A salt bridge composed of Na2SO4 connects the two beakers. The voltage measured across the electrodes is found to be + 0.42 volt.
(c) Describe what is happening at the anode. (Include any equations that may be useful.) Problem A student places a copper electrode in a 1 M solution of CuSO4 and in another beaker places a silver electrode in a 1 M solution of AgNO3. A salt bridge composed of Na2SO4 connects the two beakers. The voltage measured across the electrodes is found to be + 0.42 volt. (d) Write the balanced overall cell equation.
(e) Write the standard cell notation. Problem A student places a copper electrode in a 1 M solution of CuSO4 and in another beaker places a silver electrode in a 1 M solution of AgNO3. A salt bridge composed of Na2SO4 connects the two beakers. The voltage measured across the electrodes is found to be + 0.42 volt. (f) The student adds 4 M ammonia to the copper sulfate solution, producing the complex ion Cu(NH3)4+2 (aq). The student remeasures the cell potential and discovers the
voltage to be 0.88 volt. What is the Cu2+ (aq) concentration in the cell after the ammonia has been added?
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