Chapter 18

 Define key terms and concepts.  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 reactions for redox chemical equations.  Explain how an electrochemical cell works.

 Calculate the standard emf of a galvanic cell.  Relate Gibbs Free Energy, the equilibrium constant, and standard emf.  Determine if a reaction will occur spontaneously using the Nernst Equation.

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

 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.

 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

 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.

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

 Identify what is being oxidized and reduced in the following reactions: H 2 + Ag +  Ag + H + Fe + CuSO 4 → FeSO 4 + Cu H 2 + O 2  H 2 O Cu + AgCl  CuCl 2 + Ag

 Cellular Respiration C 6 H 12 O O 2 → 6 CO H 2 O + Energy  Photosynthesis 6 CO H 2 O + Light Energy→ C 6 H 12 O O 2  Fermentation C 6 H 12 O 6 → 2 C 2 H 5 OH + 2 CO 2

 Iron is oxidized to produce Iron (II) hydroxide, then Iron (III) Hydroxide, which is typically written as Fe 2 O 3 x H 2 O  Salt water can act as an electrolyte to facilitate this reaction.

 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.

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

 Most abundant element in the universe, but highly flammable  Often used to release metals from their ores after mining WO 3 + 3H 2  W + 3H 2 O

 Identify the oxidizing and reducing agents in the following reactions: H 2 + O 2  H 2 O Al + 3O 2  Al 2 O 3 Cu + AgNO 3  Cu(NO 3 ) 2 + Ag

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

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

Balance the following redox reaction. Identify what is being oxidized and what is being reduced. Fe + S 8  FeS

Balance the following redox reaction. Identify what is being oxidized and what is being reduced. MgCl 2 + Fe  Mg + FeCl 3

Balance the following redox reaction. Identify what is being oxidized and what is being reduced. Mg + O 2  MgO

Balance the following redox reaction. Identify what is being oxidized and what is being reduced. Fe 3+ + Sn 2+  Fe 2+ + Sn 4+

 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.

Cr 2 O 7 2- (aq) + 6Fe 2+ (aq)  2Cr 3+ (aq) + 6Fe 3+ (aq)

5Fe 2+ (aq) + MnO 4 - (aq)  5Fe 3+ (aq) + Mn 2+ (aq)

In a concentrated solution, zinc metal reduces nitrate ion to ammonium ion, and zinc is oxidized to Zn 2+. Write the balanced net ionic equation for this reaction.

Iodic Acid, HIO 3, can be prepared by reacting I 2 with concentrated nitric acid. Write the balanced equation for this if the skeleton reaction is I 2(s) + NO 3 - (aq)  IO 3 - (aq) + NO 2(g)

Balance the following redox reaction that occurs in acidic solution: H 2 S + NO 3 -  S 8 + NO 2

 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 H 2 O molecules that appear on both sides of the reaction.  Combine the half reactions together to create the final overall reaction.

Permanganate ion oxidizes sulfite ion in basic solution according to the following skeletal question. Write the balanced equation for this redox reaction. MnO 4 - (aq) + SO 3 2- (aq)  MnO 2(s) + SO 4 2-

Balance the following redox reaction that occurs in basic solution: Mn 2+ + ClO 3 -  MnO 2 + ClO 2

Balance the following redox reaction that occurs in basic solution: H 2 O 2 + ClO 2  ClO O 2

 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.

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

 Reaction at Anode (oxidation): Zn (s)  Zn 2+ (aq) + 2e -  Reaction at Cathode (reduction): 2MnO 2(s) + H 2 O + 2e -  Mn 2 O 3(s) + 2OH - (aq)  Overall Reaction: Zn (s) + 2MnO 2(s) + H 2 O  Zn 2+ + Mn 2 O 3(s) + 2OH - (aq)

 Cell Voltage ◦ Voltage across the electrodes of a galvanic cell ◦ Also called the cell potential ◦ Measured using a voltmeter  Electromotive Force, E ◦ Also called emf ◦ A measure of voltage  Cell Diagram Zn (s) ∣Zn 2+ (1M)∥Cu 2+ (1M)∣Cu (s)

 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 (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

 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: Cl 2(g), H 2 O 2(aq), Fe 3+ (aq) Order the following reducing agents by increases strength under standard-state conditions: H 2(g), Al (s), Cu (s)

Using table 19.1, calculate the E° cell for the following reaction: Mg (s) + HCl (aq)  MgCl 2(aq) + H 2(g)

Using table 19.1, calculate the E° cell for the following reaction: Cu (s) + AgNO 3(aq)  Cu(NO 3 ) 2(aq) + Ag (s)

 The standard emf can be related to Gibbs Free Energy ( ∆G) and the equilibrium constant, K.  Relates to standard states ◦ Where R = J/Kmole ◦ F is Faraday’s Constant = 9.647x10 4 C/mole e - ◦ n is the number of moles of e - ◦ K is the equilibrium constant ∆G° = -nFE° cell E° cell = log K RT nF ∆G° = -RTlnK

∆G°KE° cell Reaction Under Standard State Conditions ->1+Favors the Products 0=10At Equilibrium +<1-Favors the Reactants

Using the standard reduction potentials provided in table 19.1, calculate the equilibrium constant for the following reaction: Mg (s) + HCl (aq)  MgCl 2(aq) + H 2(g)

Using the standard reduction potentials provided in table 19.1, calculate the equilibrium constant for the following reaction: Cu (s) + AgNO 3(aq)  Cu(NO 3 ) 2(aq) + Ag (s)

Using the Gibbs Free Energy values provided in Appendix A-2, calculate the equilibrium constant for the following reaction: Cu (s) + Fe 3+ (aq)  Cu 2+ (aq) + Fe (s)

Using the Gibbs Free Energy values provided in Appendix A-2, calculate the equilibrium constant for the following reaction: H 2(g) + O 2(g)  H 2 O (l)

Using the standard reduction potentials provided in table 19.1, calculate the Gibbs Free Energy for the following reaction: N 2(g) + O 2(g)  NH 3(g)

Using the standard reduction potentials provided in table 19.1, calculate the Gibbs Free Energy for the following reaction: Cu (s) + AgNO 3(aq)  Cu(NO 3 ) 2(aq) + Ag (s)

 Relates emf and the concentrations of the reactants in nonstandard states. ◦ Where R = J/Kmole ◦ F is Faraday’s Constant = 9.647x10 4 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 Cu 2+ is 0.25M and Fe 3+ is 0.20M? Cu (s) + Fe 3+ (aq)  Cu 2+ (aq) + Fe (s)

Using the Nernst Equation, determine if the following reaction will proceed spontaneously at 298K if the concentration of Cu(NO 3 ) 2 is 0.015M and AgNO 3 is 0.030M? Cu (s) + AgNO 3(aq)  Cu(NO 3 ) 2(aq) + Ag (s)

Using the Nernst Equation, determine if the following reaction will proceed spontaneously at 298K if the concentration of MgCl 2 is 0.6M and HCl is 0.55M? Mg (s) + HCl (aq)  MgCl 2(aq) + H 2(g)