Electrochemistry Chapter 20

oxidation: lose e- -increase oxidation number reduction: gain e- -reduces oxidation number LEO goes GER Oxidation-Reduction Review

Oxidation Numbers Review Review rules p HNO 3 6. Ag 10. CO 2 1.CuCl 2 7. PbSO (NH 4 ) 2 Ce(SO 4 ) 3 1.O 2 8. PbO Cr 2 O 3 1.H 2 O 2 9. Na 2 C 2 O 4 1.MgSO 4

Oxidation-Reduction Review 1.CH 4 (g) + H 2 O(g) → CO(g) + 3H 2 (g) 1.2AgNO 3 (aq) + Cu(s) → Cu(NO 3 ) 2 (aq) + 2Ag(s) 1.Zn(s) + 2HCl(aq) → ZnCl 2 (aq) + H 2 (g) 1.2H + (aq) + 2CrO 4 2- (aq) → Cr 2 O 7 2- (aq) + H 2 O(l)

Balancing Redox Reactions Half Reactions: show only the oxidation or only the reduction part of the reaction Sn 2+ (aq) + 2Fe 3+ (aq) → Sn 4+ (aq) + 2Fe 2+ (aq) ox.: Sn 2+ (aq) → Sn 4+ (aq) + 2e - red.: 2e - + 2Fe 3+ (aq) → 2Fe 2+ (aq)

Balancing Redox Reactions 1.split into the half-reactions 2.balance each half-reactions a.first balance elements other than H and O b.balance O atoms by adding H 2 O as needed c.balance H atoms by adding H + as needed d.balance charge by adding e - as needed 3.multiply half-reactions by as needed to make the # of e - lost in ox. equal to # of e - gained in red. 4.add half reactions, simplify by canceling 5.double check everything

Balancing Redox Reactions What is reduced? What is oxidized? Reducing agent? Oxidizing agent?

Balancing Redox Reactions MnO 4 - (aq) + C 2 O 4 2- (aq) → Mn 2+ (aq) + CO 2 (aq)

Balancing Redox Reactions Do free electrons appear anywhere in the balanced equation for a redox reaction? Explain

Balancing Redox Reactions Cr 2 O 7 2- (aq) + Cl - (aq) → Cr 3+ (aq) + Cl 2 (g) (acidic solution)

Balancing Redox Reactions Cu(s) + NO 3 - (aq) → Cu 2+ (aq) + NO 2 (aq) (acidic)

Balancing Redox Reactions Mn 2+ (aq) + NaBiO 3 (s) → Bi 3+ (aq) + MnO 4 - (aq) (acidic)

Balancing Redox Reactions CN - (aq) + MnO 4 - (aq) → CNO - (aq) + MnO 2 (s) (basic)

Balancing Redox Reactions NO 2 - (aq) + Al(s) → NH 3 (aq) + Al(OH) 4 - (aq) (basic)

Balancing Redox Reactions Cr(OH) 3 (s) + ClO - (aq) → CrO 4 2- (aq) + Cl 2 (g)

Electrochemistry study of the interchange of chemical and electrical energy

Galvanic or Voltaic Cells -Device in which chemical energy is changed to electrical energy -Uses a spontaneous redox reaction to produce a current that can be used to do work -oxidizing agent is separated from the reducing agent and the electrons are forced to transfer through a wire -the current produced through the wire can then be used to do work

Galvanic or Voltaic Cells

Salt Bridge: allows ions to flow so the net charge is zero -without the salt bridge, the current would stop flowing due to a charge buildup -negative at red. side and positive at ox. side

Galvanic or Voltaic Cells Anode: where oxidation occurs Cathode: where reduction occurs AN OX, Red Cat

Why does Na + migrate into the cathode half-cell as the cell reaction proceeds?

Cell Potential Cell Potential: (E° cell ) -the oxidizing agent pulls electrons through the wire from the reducing agent. aka. electromotive force -unit is the volt: 1V = 1J of work per coulomb of charge (J/coulomb) -measured with a voltmeter: -drawing of current through a known resistance -or a potentiometer: measures opposition to current (compares to a known emf)

Standard Reduction Potentials All galvanic cell reactions are redox reactions that can be broken down into 2 half reactions -can assign a potential to each half reaction and calculate the cell potential by adding the potentials of the two half reactions The standard: -platinum electrode in contact with 1M H+ and bathed by H2(g) at 1atm has a reduction potential of 0V → called a standard hydrogen electrode 2H + + 2e - → H 2 E° = 0v

Standard Hydrogen Electrode

Standard Reduction Potentials -Compare all half reactions to the standard hydrogen electrode to obtain the reduction potential of the half-reaction E° means standard state: 1M, 1atm -Potentials of half reactions are given as reduction potentials - pink sheet -oxidation potentials are the reverse E° is an intensive property: does not change when multiplied E° cell = E° cathode - E° anode

SRP Calculation

Calculating Cell Potential Voltaic cell based on the reaction: Fe 3+ (aq) + Cu(s) → Cu 2+ + Fe 2+

Calculating Cell Potential Mg anode and Sn cathode

Calculating Cell Potential Cu cathode and Zn anode

Calculating Cell Potential Cr anode and Ag cathode

Strength of Oxidizing & Reducing Agents more negative E = most likely to be reversed and run as an oxidation reaction -Li + : E = -3.05: poor oxidizing agent -difficult to reduce (gain e-) - it is a very good reducing agent: will easily lose e- Would the halogens be good reducing agents or good oxidizing agents? Why

SRP Table In the reduction half reaction: the reactants are oxidizing agents the products are reducing agents

SRP Table

Relative Strengths of Oxidizing Agents Rank the following ions in order of increasing strength as oxidizing agents: NO 3 -, Ag +, Cr 2 O 7 2- Rank the following ions in order of decreasing strength as reducing agents: I - (aq), Fe(s), Al(s)

Determining Spontaneity +E = spontaneous -E = nonspontaneous Are the following reactions spontaneous? 1.Cu(s) + 2H + (aq) → Cu 2+ (aq) + H 2 (g) 2. Cl 2 (g) + 2I - (aq) → 2Cl - (aq) + I 2 (s)

Free Energy and Redox Gibbs Free Energy: measure of the spontaneity of a process at standard conditions -EMF (electromotive force) also indicates spontaneity ΔG = -nFE n = moles of electrons transferred in balanced equation F = Faraday’s constant = 96485C/mol or J/Vmol Units of ΔG: J/mol -means per mole of reaction

Free Energy and Redox does it make sense? ΔG = -nFE What sign of ΔG indicates a spontaneous reaction? What about E?

Free Energy and Redox Calculate ΔG and K at 298K for: 4Ag(s) + O 2 (g) + 4H + (aq) → 4Ag + (aq) + 2H 2 O

Free Energy and Redox Would the answer be different if the reaction was written: 2Ag(s) + ½ O 2 (g) + 2H + (aq) → 2Ag + (aq) + H 2 O(l)

Nonstandard Conditions Use LeChatelier’s Principle to predict the direction the reaction will shift (there is something called the Nernst equation, but that has been removed from the AP curriculum - you will probably come across it at some point)

Cell Potential and Concentration Standard conditions: 1M concentration Given: 2Al(s) + 3Mn 2+ (aq) 2Al 3+ (aq) + 3Mn(s); E° cell = 0.48V Change concentrations: 1.[Al 3+ ] = 2.0M and [Mn 2+ ] = 1.0M 2. [Al 3+ ] = 1.0M and [Mn 2+ ] = 3.0M

Electrolysis non spontaneous redox reaction -requires an outside source of electrical energy electrolytic cell - 2 electrodes in molten salt or a solution -the electrodes are inert (don’t react, just allow the reaction to occur

Electrolysis Electroplating -the cathode is active -metal deposits on the cathode Ni e - → Ni(s) -2 moles of e- needed to plate 1 mole of Ni from Ni 2+

Electrolysis 1 mole of electrons has a charge of 96,485C (C for coulombs) coulomb is the amount of charge Current is the charge per second = ampere I = q/t I = current (ampere) q = charge (coulombs) t = time (seconds)

Electrolysis Calculate the number of grams of aluminum produced in 1.00h by the electrolysis of molten AlCl 3 if the electrical current is 10.0A

Electrolysis The half-reaction for formation of magnesium metal upon electrolysis of molten MgCl 2 is Mg e - → Mg. 1.Calculate the mass of magnesium formed upon passage of a current of 60.0A for a period of 4.00x10 3 s. b. How many seconds would be required to produce 50.0g of Mg from MgCl 2 if the current is 100.0A?