Redox Reactions and Electrochemistry Chapter 17

Unit Objectives Identify redox reactions that occur in daily life. Identify what is being oxidized (reducing agent) and reduced (oxidizing agent) in a redox reaction. Write the half and overall balanced equations for redox reactions. Explain how an electrochemical cell works. Explain the difference between a galvanic and electrolytic cell.

Unit Objectives Write the cell notation for a galvanic cell. Calculate the standard reduction potential for a galvanic cell. Relate Gibbs Free Energy, the equilibrium constant, and standard reduction potential. Determine if a reaction will occur spontaneously using the Nernst Equation.

Redox Reactions in Everyday Life Reduction-Oxidation Reactions (Redox) occur all around us Burning of fuels Converting food to energy Photosynthesis Batteries Occur Together

Redox Reactions in Everyday Life Reduced items, such as food and fuels, are high in energy Oxidized items – carbon dioxide, water (byproducts) are low in energy The energy released during redox reactions are what power our homes, cars, and bodies.

Redox Reactions Involve the reactions of metals with non-metals There are three ways to view a redox reaction If something is oxidized, something else must be reduced

Redox Reactions - Electrons An increase in the oxidation number means loss of electrons and oxidation has occurred. A decrease in the oxidation number means electrons have been gained and reduction has occurred.

Redox Reactions - Electrons LEO the lion goes GER Lose Electrons – Oxidation Gain Electrons – Reduction H2 + F2 → 2HF Oxidation Reaction: H2 → 2H+ + 2e- Reduction Reaction: F2 + 2e- → 2F-

Identify what is being oxidized and reduced in the following reactions: H2 + Ag+  Ag + H+ Fe + CuSO4 → FeSO4 + Cu H2 + O2  H2O Cu + AgCl  CuCl2 + Ag

Biological Redox Reactions Cellular Respiration C6H12O6 + 6 O2 → 6 CO2 + 6 H2O + Energy Photosynthesis 6 CO2 + 6 H2O + Light Energy→ C6H12O6 + 6 O2 Fermentation C6H12O6 → 2 C2H5OH + 2 CO2

Rust Iron is oxidized to produce Iron (II) hydroxide, then Iron (III) Hydroxide, which is typically written as Fe2O3 x H2O Salt water can act as an electrolyte to facilitate this reaction.

Oxidizing and Reducing Agents

Oxidizing and Reducing Agents

Oxygen as an Oxidizing Agent One of the most common oxidizing agents (undergoes reduction). Oxygen occupies about 50% by mass of the accessible portion of the Earth and almost two-thirds of your body. Found in carbohydrates, fats, sugars, proteins contained in food. Used in combustion of fuels to power our industries, schools, and homes. Also causes corrosion, food spoilage and food decay.

Oxidants Another name for oxidizing agents Used to destroy microorganisms Cleaners such as bleach Antioxidants (such as Vitamin C) can prevent oxidation to living tissue

Hydrogen as a Reducing Agent Most abundant element in the universe, but highly flammable Often used to release metals from their ores after mining WO3 + 3H2  W + 3H2O

Identify the oxidizing and reducing agents in the following reactions: H2 + O2  H2O Al + 3O2  Al2O3 Cu + AgNO3  Cu(NO3)2 + Ag

Redox Reactions

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Writing Half Reactions Cr3+ + Zn  Cr + Zn2+ Step 1: Split reaction into half-reactions (reduction and oxidation) and balance the matter Zn  Zn2+ (oxidation) Cr3+  Cr (reduction) Step 2: Balance the charge or oxidation number with electrons Zn  Zn2+ + 2e (oxidation) 3e + Cr3+  Cr (reduction) Step 3: Check atom balance and charge balance on both sides of the equations.

Combining Half Reactions Step 4: Multiply each reaction so the electrons gained the reduction half-reaction = electrons lost in oxidation half-reaction. 2(Cr3+ + 3e  Cr) 2Cr3+ + 6e  2Cr 3(Zn  Zn2+ + 2e) 3Zn  3Zn2+ + 6e Step 5: Combine the reactions, canceling the electrons. 2Cr3+ + 6e  2Cr 3Zn  3Zn2+ + 6e 2Cr3+ 3Zn  2Cr + 3Zn2+

Balance the following redox reaction Balance the following redox reaction. Identify what is being oxidized and what is being reduced. Fe + S8  FeS

Balance the following redox reaction Balance the following redox reaction. Identify what is being oxidized and what is being reduced. MgCl2 + Fe  Mg + FeCl3  

Balance the following redox reaction Balance the following redox reaction. Identify what is being oxidized and what is being reduced. Mg + O2  MgO

Balance the following redox reaction Balance the following redox reaction. Identify what is being oxidized and what is being reduced. Fe3+ + Sn2+  Fe2+ + Sn4+

Balancing Redox Reactions in Acidic Solution Identify what is being oxidized and reduced and write the half reactions Complete and balance each half reaction Balance everything except O and H Balance the O by adding water to one side of the reaction Balance the H by adding H+ to one side of the reaction. Balance the charges by adding electrons on the more positive side. Combine the half reactions together to create the final overall reaction.

Cr2O72-(aq) + 6Fe2+(aq)  2Cr3+(aq) + 6Fe3+(aq) https://www.youtube.com/watch?v=Hz8dpVTfVlE

5Fe2+(aq) + MnO4-(aq)  5Fe3+(aq) + Mn2+(aq) https://www.youtube.com/watch?v=2lEmCK3stMs

In a concentrated solution, zinc metal reduces nitrate ion to ammonium ion, and zinc is oxidized to Zn2+. Write the balanced net ionic equation for this reaction.

I2(s) + NO3-(aq)  IO3-(aq) + NO2(g) Iodic Acid, HIO3, can be prepared by reacting I2 with concentrated nitric acid. Write the balanced equation for this if the skeleton reaction is I2(s) + NO3-(aq)  IO3-(aq) + NO2(g)

Balance the following redox reaction that occurs in acidic solution: H2S + NO3-  S8 + NO2

Balancing Redox Reactions in Basic Solution Begin by balancing the reaction as if it was in an acid solution Add the same number of OH- ions to each side of the reaction as you have H+ ions Simplify the H+ and OH- ions to be written as water molecules. Cancel any H2O molecules that appear on both sides of the reaction. Combine the half reactions together to create the final overall reaction.

MnO4-(aq) + SO32-(aq)  MnO2(s) + SO42- Permanganate ion oxidizes sulfite ion in basic solution according to the following skeletal question. Write the balanced equation for this redox reaction. MnO4-(aq) + SO32-(aq)  MnO2(s) + SO42-

Balance the following redox reaction that occurs in basic solution: Mn2+ + ClO3- MnO2 + ClO2

Balance the following redox reaction that occurs in basic solution: H2O2 + ClO2  ClO2- + O2

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Electrochemistry Field of chemistry involving the study of chemical reactions that are driven by an electrical current. The reactions that occur are redox reactions. Occur through the use of an electrochemical cell.

Electrochemistry

Electrochemistry Dry Cell Battery Mercury Battery Lead Storage Battery Lithium-Ion Battery Fuel Cells

Electrochemistry Reaction at Anode (oxidation): Zn(s)  Zn2+(aq)+ 2e- Reaction at Cathode (reduction): 2MnO2(s) + H2O + 2e-  Mn2O3(s) + 2OH-(aq) Overall Reaction: Zn(s) + 2MnO2(s) + H2O  Zn2+ + Mn2O3(s) + 2OH-(aq)

Zn(s)∣Zn2+ (1M)∥Cu2+ (1M)∣Cu(s) Electrochemistry Cell Voltage Voltage across the electrodes of a galvanic cell Also called the cell potential Measured using a voltmeter Cell Diagram Zn(s)∣Zn2+ (1M)∥Cu2+ (1M)∣Cu(s)

Electrolysis The process in which electrical energy is used to carry out a nonspontaneous chemical reaction Electrolytic Cell A setup used to carry out electrolysis

Standard Reduction Potentials Standard Reduction Potentials (E°) The voltage associated with a reduction reaction at an electrode when all states are 1M and all gases are at 1atm. Standard emf (E°cell) E°cell = E°cathode - E°anode

Standard Reduction Potentials E° apply to half-reactions Change the sign of E° if the reaction is reversed The more positive E°, the more likely the substance will be reduced. The value of E° is not affected by the amount of solution present or the number of moles in solution. Under standard-state conditions, any species on the left of side of a given half reaction will react spontaneously with any species on the right side of a given half reaction as long as that species has a lower E° value.

Order the following oxidizing agents by increasing strength under standard-state conditions: Cl2(g) , H2O2(aq) , Fe3+(aq) Order the following reducing agents by increases strength under standard-state conditions: H2(g) , Al(s) , Cu(s)

Using table 17.2, calculate the E°cell for the following reaction: Mg(s) + HCl(aq)  MgCl2(aq) + H2(g)

Using table 17.2, calculate the E°cell for the following reaction: Cu(s) + AgNO3(aq)  Cu(NO3)2(aq) + Ag(s)

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Thermodynamics of Redox Reactions The standard emf can be related to Gibbs Free Energy (∆G) and the equilibrium constant, K. Relates to standard states Where R = 8.314 J/K•mole F is Faraday’s Constant = 9.647x104 C/mole e- n is the number of moles of e- K is the equilibrium constant ∆G° = -nFE°cell ∆G° = -RTlnK E°cell = log K RT nF

Relationship between ∆G°, K and E°cell Reaction Under Standard State Conditions - >1 + Favors the Products =1 At Equilibrium <1 Favors the Reactants

Using the standard reduction potentials provided in table 17 Using the standard reduction potentials provided in table 17.2, calculate the equilibrium constant for the following reaction: Mg(s) + HCl(aq)  MgCl2(aq) + H2(g)

Using the standard reduction potentials provided in table 17 Using the standard reduction potentials provided in table 17.2, calculate the equilibrium constant for the following reaction: Cu(s) + AgNO3(aq)  Cu(NO3)2(aq) + Ag(s)

Using the standard reduction potentials provided in table 17 Using the standard reduction potentials provided in table 17.2, calculate the Gibbs Free Energy for the following reaction: N2(g) + O2(g)  NH3(g)

Using the standard reduction potentials provided in table 17 Using the standard reduction potentials provided in table 17.2, calculate the Gibbs Free Energy for the following reaction: Cu(s) + AgNO3(aq)  Cu(NO3)2(aq) + Ag(s)

The Nernst Equation Relates emf and the concentrations of the reactants in nonstandard states. Where R = 8.314 J/K•mole F is Faraday’s Constant = 9.647x104 C/mole e- n is the number of moles of e- Q is the reaction quotient E = E° - InQ RT nF

Using the Nernst Equation, determine if the following reaction will proceed spontaneously at 298K if the concentration of Cu2+ is 0.25M and Fe3+ is 0.20M? Cu(s) + Fe3+(aq)  Cu2+(aq) + Fe(s)

Using the Nernst Equation, determine if the following reaction will proceed spontaneously at 298K if the concentration of Cu(NO3)2 is 0.015M and AgNO3 is 0.030M? Cu(s) + AgNO3(aq)  Cu(NO3)2(aq) + Ag(s)

Using the Nernst Equation, determine if the following reaction will proceed spontaneously at 298K if the concentration of MgCl2 is 0.6M and HCl is 0.55M? Mg(s) + HCl(aq)  MgCl2(aq) + H2(g)

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