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CONCEPT: OXIDATION-REDUCTION REACTIONS Chemists use some important terminology to describe the movement of electrons. In reactions we have the movement of electrons from one reactant to another. L E O Agent G E R Agent Rules for Assigning an Oxidation Number (O.N.) A. General Rules 1. For an atom in its elemental form (Na, O2, S8, etc.): O.N. = 0 2. For an ion the O.N. equals the charge: Na +, Ca 2+, NO3 B. Specific Rules 1. Group 1A: O.N. = +1 2. Group 2A: O.N. = +2 3. For hydrogen: O.N. = +1 with nonmetals O.N. = -1 with metals and boron 4. For Fluorine: O.N. = -1 5. For oxygen: O.N. = -1 in peroxides (X2O2, X = Group 1(A) element) O.N. = 1 in superoxides (XO2, X = Group 1(A) element) 2 O.N. = - 2 in all other compounds 6. Group 7A O.N. = -1 (except when connected to O) Page 2
CONCEPT: OXIDATION-REDUCTION REACTIONS (PRACTICE) EXAMPLE: In the following reaction identify the oxidizing agent and the reducing agent: a. 2 C6H6 (l) + 15 O2 (g) 12 CO2 (g) + 6 H2O (g) PRACTICE: What is the oxidation number of each underlined element? a. P4 b. BO3 3- c. AsO4 2- d. HSO4 PRACTICE: In the following reaction identify the oxidizing agent and the reducing agent: a. Cr2O7 2- + 6 Fe 2+ + 14 H + 2 Cr 3+ + 6 Fe 3+ + 7 H2O Page 3
} } CHEMISTRY - CLUTCH CONCEPT: BASIC REDOX CONCEPTS OXIDATION-REDUCTION (REDOX) reactions deal with the transfer of electrons from one reactant to another. Lose Electrons Oxidation } Element becomes more positive Li (s) + Cl 2 (g) Li + (aq) + 2 Cl aq) Oxidation Gain } Element Number Electrons becomes Increases more negative Reduction } } Oxidation Number Decreases Reducing Agent (Reductant) Oxidizing Agent (Oxidant) Li (s) Li + (aq) + e Cl 2 (g) + 2 e 2 Cl (aq) Electrical Charge The units for electrical charge are measured in (C). (1.602 10 19 C) (6.022 10 23 mol 1 ) = 9.647 104 C 1 mole e charge mole e Charge of 1 electron Faraday Constant q = n F Faraday Constant Electrical Current The units for electrical current are in (A). I = q t Current Charge Time Electrical Voltage The relationship between work and voltage can be expressed as: w = E q Work Voltage Charge The relationship between Gibbs Free Energy and electric potential can be expressed as: ΔG = n F E Gibbs Free Energy mole e Faraday Constant Voltage Ohm's Law The units for resistance are in (Ω). I = E R Current Voltage Resistance Power Power represents work done per unit of time. The units for power are in (W). P = E I Power Voltage Current Page 4
CONCEPT: BALANCING REDOX REACTIONS Now that you have refreshed your memory in terms of redox reactions you now have to master balancing them. Generally, you will need to balance a redox reaction in an acidic or basic solution. Balancing A Redox Reaction in Acidic Reactions: STEP 1: Write the equation into 2 half-reactions. STEP 2: Balance elements that are not oxygen or hydrogen. STEP 3: Balance Oxygens by adding H2O. STEP 4: Balance Hydrogens by adding H +. STEP 5: Balance overall charge by adding electrons (e ) to the more side. STEP 6: Combine the half-reactions and cross out reaction intermediates. Balancing A Redox Reaction in Basic Reactions: Follow Steps 1-6 from above. STEP 7: Balance H + by adding OH ions to both sides of the chemical reaction. Page 5
PRACTICE: BALANCING ACIDIC REDOX REACTIONS EXAMPLE: Balance the following reaction in an acidic solution. NO2 + MnO4 NO3 + Mn 2+ PRACTICE: Balance the following reaction in an acidic solution. Cl2(g) + S2O3 2- (aq) Cl - (aq) + SO4 2- (aq) Page 6
PRACTICE: BALANCING BASIC REDOX REACTIONS EXAMPLE: Balance the following reaction in a basic solution. Mo3O9 3- (aq) + Br (aq) Mo (aq) + BrO4 2- (aq) PRACTICE: Balance the following reaction in a basic solution. XeO2 (aq) H2Xe (aq) + XeO4 Page 7
CONCEPT: NERNST EQUATION The Nernst Equation reveals the quantitative connection between the concentrations of compounds and cell potential. E o = Standard Cell Potential E = E o RT nf ln A b B a A A R = Gas Constant = n = Number of electrons transferred F = Faraday's constant = A = Activity represents the cell potential under non-standard conditions, while represents it under standard conditions. At 25 o C, RT F = J 8.314 ( 298.15 K) K mol = 0.0257 J 9.649 10 4 C C = 0.0257 V mol The Nernst Equation then becomes, o E Cell = E Cell By multiplying ln by 2.303 we can obtain the log function. o E Cell = E Cell 0.0257 V n 0.05916 V n lnq logq The cell potential calculated from Nernst equation is the maximum potential at the instant the cell circuit is connected. As the cell discharges and current flows, the electrolyte concentrations will change, Q increases and Ecell deceases. Over time the reaction will reach equilibrium at then Q = K and cell potential will equal zero. o E Cell = E Cell RT ln K = 0 ΔG = ΔG o RT ln K = 0 nf Page 8
CONCEPT: ELECTRICAL CURRENT EXAMPLE 1: Gold can be plated out of a solution containing Au 3+ based on the following half reaction: Au 3+ (aq) + 3 e - Au (s) a) What mass of gold is plated by a 41 minute flow of 6.8 A current? EXAMPLE 2: A solution of Mn +5 is used to plate out Mn in an electrochemical cell. If a total of 1.13 g of Mn is plated out in a total time of 1600 seconds, what was the electrical current used? (MW of Mn is 54.94 g/mol) Page 9
PRACTICE: ELECTRIC CURRENT (CALCULATIONS 1) EXAMPLE: If steady current of 15 amperes is provided by a stable voltage of 12 Volts for 600 seconds, answer each of the following questions. a) Calculate the total charge that passes through the circuit in this time. b) Calculate the total number of moles of electrons that pass through the circuit in this time. c) Calculate the total amount of energy that passes through the circuit in this time. d) Calculate the power that the battery provides during this process. Page 10
CONCEPT: GALVANIC VS. ELECTROLYTIC CELLS With REDOX reactions we will now deal with a new variable: a reaction s cell potential, which uses the variable. The greater this variable then the more likely will occur. The smaller this variable then the more likely will occur. Galvanic/Voltaic Cell: or electricity so it s a. Ionization Energy Electron Affinity Anode Cathode Producing ñ Voltage [Anode] [Cathode] ( ) ( + ) Galvanic/Voltaic Cell Cd 2+ (aq) + 2 e Cd (s) E o = - 0.40 V Ni 2+ (aq) + 2 e Ni (s) E o = - 0.25 V For all redox reactions under standard conditions it is possible to determine if a reaction is spontaneous or not. SUniverse G o K E o Reaction Under Standard State Conditions Cell Type Positive Negative >1 Positive Negative Positive <1 Negative 0 0 1 0 Page 11
CONCEPT: ELECTROLYTIC CELLS In terms of spontaneity the following correlations between the following variables can be made: ΔG o K E o ΔS o Q vs. K Reaction Classification Cell Type < 0 > 1 > 0 > 0 Q < K > 0 < 1 < 0 < 0 Q > K = 0 = 1 = 0 = 0 Q = K Electrolytic Cell: A non-spontaneous electrochemical cell that electricity and so requires a battery. Electron Affinity Ionization Energy ( ) ( + ) Electrolytic Cell Page 12
PRACTICE: GALVANIC VS. ELECTROLYTIC CELLS (CALCULATIONS) EXAMPLE: A certain electrochemical cell involves a five electron change and has a value of Keq = 3.0 x 10 16 at 298 K. The value of H o for the reaction is -68.3 kj/mol. Calculate the values of G o, E o, for a standard electrochemical cell constructed based on this reaction and also S o for the reaction. PRACTICE: Given the following standard reduction potentials, Hg2 2+ (aq) + 2 e 2 Hg(l) E = +0.789 V Hg2Cl2(s) + 2 e 2 Hg(l) + 2 Cl - (aq) E = +0.271 V determine Ksp for Hg2Cl2(s) at 25 C. Page 13
PRACTICE: CELL NOTATION (CALCULATIONS 1) EXAMPLE: The cell notation for a redox reaction is given as the following at (T = 298 K): Zn (s) Ι Zn 2+ (aq, 0.37 M) ΙΙ Ni 2+ (aq, 0.059 M) Ι Ni (s) a) Write the balanced half-reactions occurring at the anode and the cathode. b) Write out the complete balanced redox reaction. c) Determine the E o cell. d) Calculate the maximum electrical work that can be produced by this cell. e) Calculate the reactant quotient, Q, for this cell and the cell potential under non-standard conditions. Page 14
PRACTICE: GALVANIC VS. ELECTROLYTIC CELLS (CALCULATIONS 2) EXAMPLE: Answer each of the following questions based on the following half reactions: HALF REACTIONS E o (V) Cl2 (g) + 2 e 2 Cl (aq) + 1.36 l2 (g) + 2 e 2 l (aq) + 0.535 Pb 2+ (aq) + 2 e Pb (s) - 0.126 V 2+ (aq) + 2 e V (s) - 1.18 a) Which is the strongest oxidizing agent? b) Which is the strongest reducing agent? c) Will I (aq) reduce Cl2 (g) to Cl (g)? Page 15
Electric current, or flow of electrons, is measured in Amperes (A). One Ampere is the delivery of one coulomb (C) of charge per second. What mass of Zinc (in g) is oxidized (to Zn 2+ ) by a dry cell battery that supplies 125 ma of current for two hours (recall that Faraday s constant is the charge in coulombs on a mole of electrons)? a. 0.03 b. 0.1 c. 0.3 d. 1.0 e. 3.0 Page 16
Consider the combustion of formaldehyde CH2O and the data below for the following five questions. I. CH2O(g) + O2(g) CO2(g) + H2O(l) II. 4H + + 4e - + O2(g) 2H2O(l) ΔE = 1.23V Which compound is reduced in reaction I? a. CH2O b. H2O c. O2 d. CO2 e. None of these What is the change in oxidation number of the carbon in reaction I? a. -4 b. -1 c. 0 d. 1 e. 4 What is the ΔG for the combustion of formaldehyde (kj/mol)? a. 224 b. -521 c. 96 d. -150 e. more date is required How many electrons are required to balance the half reaction CH2O CO2 in acidic solution? a. 1 b. 2 c. 3 d. 4 e. 5 What is the standard half cell potential (in volts) for the CH2O oxidation? a. 1.7 b. 0.76 c. 0.12 d. 3.33 e. 1.69 Page 17
The purpose of a galvanic cell is to: a. Transduce chemical energy to electrical energy. b. Purify solids. c. Allow for oxidation without reduction. d. To consume electricity. Page 18
During the process for electrolysis of water a current is passed through water and produces hydrogen gas and oxygen gas. Which of the following statements is true? a. O2 gas is produced at the anode. b. H2 gas is produced at the cathode. c. In the reaction, H2 moles are twice the O2 moles. d. All of the following are correct. Page 19
Which statement is false? a. Reduction occurs at the cathode. b. A reducing agent will lose electrons. c. Cations migrate to the cathode in both electrolytic and electrochemical cells. d. Li (s) is the strongest oxidizing agent; F2 is the strongest reducing agent. Page 20
Which of the following reactions may occur at the anode? a. Ga 3+ (aq) + 3 e Ga (s) b. Cu 2+ (aq) + 2 e Cu (s) c. 2 Cl (aq) Cl2 (g) + 3 e d. Co (s) + e Co + (aq) Page 21
Define a salt bridge. A) A pathway, composed of salt water, that ions pass through. B) A pathway in which no ions flow. C) A pathway between the cathode and anode in which ions are reduced. D) A pathway between the cathode and anode in which ions are oxidized. E) A pathway by which counterions can flow between the half-cells without the solutions in the half-cell totally mixing. Page 22
What statement is NOT true about standard electrode potentials? A) E cell is positive for spontaneous reactions. B) Electrons will flow from more negative electrode to more positive electrode. C) The electrode potential of the standard hydrogen electrode is exactly zero. D) E cell is the difference in voltage between the anode and the cathode. E) The electrode in any half-cell with a greater tendency to undergo reduction is positively charged relative to the standard hydrogen electrode and therefore has a positive E. Page 23
Use the standard reduction potentials below to determine which element or ion is the best reducing agent. Pd 2+ (aq) + 2 e Pd (s) E = + 0.90 V 2 H + (aq) + 2 e H2 (g) E = 0.00 V Mn 2+ (aq) + 2 e Mn (s) E = 1.18 V a) Pd (s) b) H + (aq) c) Mn 2+ (aq) d) H2 (g) Page 24
Consider an electrochemical cell where the following reaction takes place: Na2O (aq) + Ba (s) 2 Na (s) + BaO (aq) What is the cell notation for this cell? What is the ratio of oxidizing agent to reducing agent? Page 25
CONCEPT: BATTERIES Batteries can be thought of as portable galvanic or voltaic cells that generate electricity. Dry-Cell Batteries These types of batteries are called dry-cell because they lack large amounts of liquid water. In a typical dry-cell battery the zinc sleeve acts like the : Zn (s) Zn 2+ (aq) + 2 e E o = - 0.76 V The portion has a carbon tube immersed in a damp paste of MnO2 that also possesses NH4Cl. 2 MnO2 (s) + 2 NH4 + (aq) + 2 e Mn2O3 (s) + 2 NH3 (g) + H2O (l) E o = 0.74 V The counterpart to the more acidic dry-cell battery is the more commonly used battery. Zn (s) + 2 OH (aq) Zn(OH)2 (s) + 2 e 2 MnO2 (s) + 2 H2O (l) + 2 e 2 MnO(OH) (s) + 2 OH (aq) Page 26