Electrochemistry Dr. Susan Lagrone

ASSIGNING OXIDATION NUMBERS 1. F is always -1 2. Elements in standard state = zero 3. Monatomic ions = charge 4. H = +1 with nonmetals (H= -1 with metals) 5. O = -2 usually, but in peroxide = -1, superoxides it can even be a fraction or with F it is positive! 6. Memorize polyatomic ions

OXIDATION – REDUCTION (REDOX) DEFINITIONS: Oxidation is loss of eletrons oxidation number increases (more positive) Reduction is gain of eletrons oxidation number decreases (more negative)

OXIDATION – REDUCTION (REDOX) Given the reaction below, identify Species oxidized Species reduced 2 H2O2 (aq) → O2 (g) + 2 H2O(l), OIL RIG ASSIGN OX # H = +1 H = +1 O = 0 O = -1 O = -2

OXIDATION – REDUCTION (REDOX) Given the reaction below, identify Species oxidized Species reduced 2 H2O2 (aq) → O2 (g) + 2 H2O(l), H2O2 oxygen -1 to zero H2O2 oxygen -1 to -2 H = +1 O = 0 H = +1 ASSIGN OX # O = -1 O = -2

OXIDATION – REDUCTION (REDOX) FUN FACT: Disproportionation - the process by which the same substance is both OXIDIZED AND REDUCED 2 H2O2 (aq) → O2 (g) + 2 H2O(l), O = 0 O = -1 O = -2

OXIDATION – REDUCTION (REDOX) DEFINITIONS: Oxidizer is the substance that causes oxidation to occur, i.e. is reduced (sometimes called oxidizing agent) Reducer is the substance that causes reduction to occur, i.e. is oxidized (sometimes called reducing agent)

OXIDATION – REDUCTION (REDOX) Given the reaction below, identify Oxidizer Reducer Al(NO3)3 (aq) + Mg (s) → Al (s) + Mg(NO3)2(aq) Reduced, Al3+ (aq) Oxidized, Mg (s) Al = +3 Al = 0 ASSIGN OX # O = -2 N = +5 Mg = 0 Mg = +2 O = -2 N = +5

BALANCING REDOX RXNS BY HALF RXNS Assign oxidation numbers to split reaction into oxidation half-reaction and reduction half-reaction Balance the atom oxidized /reduced and do not change these coefficients Balance oxygen with water Balance hydrogen with H+ ions Add electrons to the more positive side to balance the charge All reactions are balanced by atoms AND charge

BALANCING REDOX RXNS BY HALF RXNS Assign oxidation numbers to split reaction into oxidation half- reaction and reduction half-reaction Balance the atom oxidized /reduced and do not change these coefficients Balance oxygen with water Balance hydrogen with H+ ions Add electrons to the more positive side to balance the charge The result is balanced in ACID CHECK: All reactions are balanced by atoms AND charge

BALANCING in BASE – ADD THESE THREE STEPS 6. Add OH- for each H+ to BOTH SIDES OF EQUATION 7. Combine H+ and OH- appearing on same side to make WATER. 8. If possible SUBTRACT WATER FROM BOTH SIDES (cancel waters so it appears only on one side) CHECK: All reactions are balanced by atoms AND charge

BALANCING REDOX RXNS BY HALF RXNS YOU TRY in acid: Cr(s) + NO3- (aq)  Cr3+ (aq) + NO (g) CH3OH(aq) + Ce4+ (aq)  CO2 (aq) + Ce3+ (aq) Fe2+ (aq) + MnO4- (aq)  Mn2+ (aq) + Fe3+ (aq) YOU TRY in base: 4. MnO4- (aq) + ClO4- (aq)  Mn2+ (aq) + ClO3- (aq)

Voltaic Cell (Galvanic) – it is a battery Ecell > 0 (positive) SPONTANEOUS CHEMICAL RXN Thermodynamically favorable ∆G < 0 (negative) Keq > 1 (rxn is product favored)

Voltaic Cell (Galvanic) – it is a battery Ecell is positive Anode = oxidation ( an ox); lose eletrons Cathode = reduction (redc at); gain electrons e- flow from anode to cathode (FATCAT) Salt bridge- maintains electrical neutrality (keeps charge from building up in the cell) Voltmeter – measures cell potential Ammeter – measures current NOTICE: No battery or power source used Line notation – anode is written first Zn(s)  Zn2+ (1.0 M)  Cu2+ (1.0M) Cu Bar denotes change in phase of matter Double bar denotes salt bridge (porous disk) Standard conitions are 1.0 M , 1 atm , 25°C https://www.google.com/search?q=galvanic+cell&espv=2&source=lnms&tbm=isch&sa=X&ved=0ahUKEwirnoWW3Y3TAhXD5CYKHR4_DfcQ_AUIBigB&biw=911&bih=425#imgrc=iDbbso21MeXu1M:

Voltaic Cell (Galvanic) – it is a battery Ecell is positive Question: 1. Which electrode will decrease in mass as the cell operats? 2. Write the cell rxn which occurs at the cathode. 3. Calculate standard cell potential. Reduction potentials: Zn2+ Zn(s) = -0.76V Cu2+ Cu = 0.34V 4. Which ions flow into the cell containing the copper electroe? NOTICE: No battery or power source used https://www.google.com/search?q=galvanic+cell&espv=2&source=lnms&tbm=isch&sa=X&ved=0ahUKEwirnoWW3Y3TAhXD5CYKHR4_DfcQ_AUIBigB&biw=911&bih=425#imgrc=iDbbso21MeXu1M:

The hydrogen electrode 2H+ (aq) + 2 e-  H2 (g) Defined: E0 = 0.00V All reduction potentials are measured relative to this electrode. http://chemistry.stackexchange.com/questions/4686/gas-electrode-working

Problem 5: Sketch the galvanic cell given the half cell reactions below. Mn2+ + 2 e-  Mn(s) E0 = -1.18V Fe3+ + 3 e-  Fe(s) E0 = -0.036V

Indicate electron flow in cell Show ion movement through bridge Problem 5: continued Label the anode Indicate electron flow in cell Show ion movement through bridge Write the line notation for this cell Calculate ∆G0 for the cell

2 Al(s) + 3 Zn2+ (aq) → 2 Al3+ (aq) + 3 Zn(s) Problem 6: 2 Al(s) + 3 Zn2+ (aq) → 2 Al3+ (aq) + 3 Zn(s) Respond to the following statements and questions that relate to the species and the reaction represented above. Write the complete electron configuration (e.g., 1s2 2s2 . . .) for Zn2+. Which species, Zn or Zn2+, has the greater ionization energy? Justify your answer. Identify the species that is oxidized in the reaction

2 Al(s) + 3 Zn2+ (aq) → 2 Al3+ (aq) + 3 Zn(s) Problem 7: 2 Al(s) + 3 Zn2+ (aq) → 2 Al3+ (aq) + 3 Zn(s) Respond to the following statements and questions that relate to the species and the reaction represented above. Indicate the movement of ions in the salt bridge. Label the cathode. Indicate electron flow in the cell. Calulate standard cell potential, E0. Which electrode gains mass during cell operation If the concentration of zinc nitrate changes to 0.50M will the cell potential increase, decrease, or remain the same? Justify your answer.

Problem 8: Place the following in order of increasing oxidizing strength. Use table on next slide. (all under standard conditions) Cadmium(II), iodate ion, potassium ion, water, iodine, and AuCl4-

Problem 9: Answer the following questions using the table on previous slide. Is H+ (aq) capable of oxidizing Cu(s) to Cu2+ (aq)? Is Fe+3 (aq) capable of oxidizing I- (aq)? Is H2 (g) capable of reducing Ag+ (aq)? Is Fe+2 (aq) capable of reducing Cr3+ (aq) to Cr2+ (aq)?

Problem 10: Consider the species at standard conditions Problem 10: Consider the species at standard conditions. Na+, Cl-, Ag+, Ag, Zn2+, Zn, Pb Which is the strongest oxidizing agent? (oxidizer) Which is th strongrest reducing agent? Which species can be oxidized by sulfate ion, SO42-, in acid? Which species can be reduced by aluminum solid?

Problem 11: Use the half reactions below to explain why household bleach ( a highly alkaline solution of hypochlorite) should not be mixed with household ammonia or glass cleansers containing ammonia. ClO- + H2O + 2 e-  2 OH- + Cl- E0 = 0.90V N2H4 + 2 H2O + 2 e-  2 OH- + 2 NH3 E0 = -0.10V