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ElectrochemistryElectrochemistry
RedoxOxidation States
Voltaic CellsBalancingEquations
Cell EMF
ConcentrationCells
EquilibriumNernst
Equation
Corrosion Batteries Electrolysis
04/19/23
Oxidation and Reduction• Metal undergoes corrosion, it loses electrons to form cations:
Ca(s) +2H+(aq) Ca2+(aq) + H2(g)• Voltaic Cell:Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)
• Oxidized: atom, molecule, or ion becomes more positively charged. – Oxidation is the loss of electrons.
• Reduced: atom, molecule, or ion becomes less positively charged.– Reduction is the gain of electrons.
Oxidation-Reduction ReactionsOxidation-Reduction ReactionsOxidation-Reduction ReactionsOxidation-Reduction Reactions
• Metal undergoes corrosion, it loses electrons to form cations:
Ca(s) +2H+(aq) Ca2+(aq) + H2(g)
• Voltaic Cell: Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)
Oxidation and Reduction
Oxidation-Reduction ReactionsOxidation-Reduction ReactionsOxidation-Reduction ReactionsOxidation-Reduction Reactions
Oxidation Numbers• Oxidation numbers are assigned by a series of rules:
1. If the atom is in its elemental form, the oxidation number is zero. e.g., Cl2, H2, P4.
2. For a monatomic ion, the charge on the ion is the oxidation state.3. Nonmetal usually have negative oxidation numbers:
(a) Oxidation number of O is usually –2. The peroxide ion, O2
2-, has oxygen with an oxidation number of –1.(b) Oxidation number of H is +1 when bonded to nonmetals and –1 when bonded to metals.(c) Halogens generally -1, oxid. # of F is always –1.
4. The sum of the oxidation numbers for the atom is the charge on the molecule (zero for a neutral molecule).
Oxidation-Reduction ReactionsOxidation-Reduction ReactionsOxidation-Reduction ReactionsOxidation-Reduction Reactions
The Periodic TableThe Periodic TableAlkali metals
Alkaline Earth metals
Halogens
Noble Gases
Chalcogens
Transition Metals
Inner Transition Metals
Common (Type II) Cations
Ion Systemic Name
Older Name
Fe3+ Iron(III) Ferric
Fe2+ Iron(II) Ferrous
Cu2+ Copper(II) Cupric
Cu+ Copper(I) Cuprous
Co3+ Cobalt(III) Cobaltic
Co2+ Cobalt(II) Cobaltous
Sn4+ Tin(IV) Stannic
Sn2+ Tin(II) Stannous
Pb4+ Lead(IV) Plumbic
Pb2+ Lead(II) Plumbous
Oxidation Numbers
Oxidation-Reduction ReactionsOxidation-Reduction ReactionsOxidation-Reduction ReactionsOxidation-Reduction Reactions
1. Determine the oxidation numbers (oxidation states) on the Sulfur atom for: (a) H2S (b) S8 (c) SCl2
(d) Na2SO3 (e) SO42-
2. Determine o.s. for the Redox reaction in Ni-Cd Batteries:
Cd(s) + NiO2(s) + 2H2O(l) Cd(OH)2(s) + Ni(OH)2(s)
3. Determine o.s. for the following Redox reaction:
Al(s) + MnO4-(aq) + 2H2O(l) Al(OH)4
-(aq) + MnO2(s)
HW-Answers
• Law of conservation of mass: the amount of each element present at the beginning of the reaction must be present at the end.
• Conservation of charge: electrons are not lost in a chemical reaction.
Half Reactions• Half-reactions are a convenient way of separating
oxidation and reduction reactions.
Balancing Redox ReactionsBalancing Redox ReactionsBalancing Redox ReactionsBalancing Redox Reactions
Half Reactions• The half-reactions for
Sn2+(aq) + 2Fe3+(aq) Sn4+(aq) + 2Fe2+(aq)
are ?
• Oxidation: electrons are products.• Reduction: electrons are reagents.
Balancing Redox ReactionsBalancing Redox ReactionsBalancing Redox ReactionsBalancing Redox Reactions
Balancing Equations by the Method of Half Reactions1. Write down the two half reactions.
2. Balance each half reaction:
a. First with elements other than H and O.
b. Then balance O by adding water.
c. Then balance H by adding H+. d. Finish by balancing charge by adding electrons. 3. Multiply each half reaction to make the number of electrons equal.
4. Add the reactions and simplify.5. Check for balanced atoms and charges! Balance: MnO4
-(aq) + C2O42-(aq) Mn2+(aq) + CO2(g) (Acidic)
Balancing Redox ReactionsBalancing Redox ReactionsBalancing Redox ReactionsBalancing Redox Reactions
Balancing Equations for Reactions Occurring in Basic Solution
• Use OH- and H2O rather than H+ and H2O.
• Follow same method as in Acidic Solution, but OH- is added to “neutralize” the H+ used.
• Consider:
CN-(aq) + MnO4-(aq) CNO-(aq) + MnO2(s) [ Basic Solution ]
Balancing Redox ReactionsBalancing Redox ReactionsBalancing Redox ReactionsBalancing Redox Reactions
• Energy released in a spontaneous redox reaction is used to perform electrical work.
• Voltaic or galvanic cells are devices in which electron transfer occurs via an external circuit.
• Voltaic cells are spontaneous.
• If a strip of Zn is placed in a solution of CuSO4, Cu is deposited on the Zn and the Zn dissolves by forming Zn2+.
Voltaic CellsVoltaic CellsVoltaic CellsVoltaic Cells
Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)
• Flow of electrons from anode to cathode is spontaneous.• Electrons flow from anode to cathode because the cathode
has a lower electrical potential energy than the anode.• Potential difference: difference in electrical potential.
Measured in volts.• One volt is the potential difference required to impart one
joule of energy to a charge of one coulomb:
Cell EMFCell EMFCell EMFCell EMF
• Electromotive force (emf) is the force required to push electrons through the external circuit.
• Cell potential: Ecell is the emf of a cell.
• For 1M solutions at 25 C (standard conditions), the standard emf (standard cell potential) is called Ecell.
C 1J 1
V 1
Cell EMFCell EMFCell EMFCell EMF
Standard Reduction (Half-Cell) Potentials
• Convenient tabulation of electrochemical data.
• Standard reduction potentials, Ered are measured relative to the standard hydrogen electrode (SHE).
Cell EMFCell EMFCell EMFCell EMF
Standard Reduction (Half-Cell) Potentials
• Reactions with Ered < 0 are spontaneous oxidations relative to the SHE.
• The larger the difference between Ered values, the larger Ecell.
• In a voltaic (galvanic) cell (spontaneous) Ered(cathode) is more positive than Ered(anode).
•
Cell EMFCell EMF
anodecathode redredcell EEE
Calculate Eocell for: 2 Al(s) + 3 I2(s) 2 Al3+(aq) + 6 I-(aq)
)(oxidationreduction)(cell EEE
• In a voltaic (galvanic) cell (spontaneous) Ered(cathode) is more positive than Ered(anode) since
• More generally, for any electrochemical process
• A positive E indicates a spontaneous process (galvanic cell).• A negative E indicates a nonspontaneous process.
Spontaneity of Redox ReactionsSpontaneity of Redox ReactionsSpontaneity of Redox ReactionsSpontaneity of Redox Reactions
anodecathode redredcell EEE
processoxidation processreduction redredcell EEE
)(oxidationreduction)(cell EEE
EMF and Free-Energy Change• Can show that under Standard Conditions:
G is the change in free-energy, n is the number of moles of electrons transferred, F is Faraday’s constant, and Eo is the emf of the cell.
• Since n and F are positive, if G > 0 then E < 0.
Spontaneity of Redox ReactionsSpontaneity of Redox ReactionsSpontaneity of Redox ReactionsSpontaneity of Redox Reactions
oo EFnG
J/(V·mol) 96,500C/mol 500,961 F
Calculate Eo and ΔGo for:
(a)4Ag(s) + O2(g) + 4H+(aq) 4Ag+(aq) + 2H2O(l)
(b)2Ag(s) + ½O2(g) + 2H+(aq) 2Ag+(aq) + H2O(l)
Calculate Eo and ΔGo for:
(a)4Ag(s) + O2(g) + 4H+(aq) 4Ag+(aq) + 2H2O(l)
(b)2Ag(s) + ½O2(g) + 2H+(aq) 2Ag+(aq) + H2O(l)
Calculate Eo and ΔGo for:
(a)4Ag(s) + O2(g) + 4H+(aq) 4Ag+(aq) + 2H2O(l)
(b)2Ag(s) + ½O2(g) + 2H+(aq) 2Ag+(aq) + H2O(l)
Calculate Eo and ΔGo for:
(a)4Ag(s) + O2(g) + 4H+(aq) 4Ag+(aq) + 2H2O(l)
(b)2Ag(s) + ½O2(g) + 2H+(aq) 2Ag+(aq) + H2O(l)
The Nernst Equation
• The Nernst equation can be simplified by collecting all constants together using a temperature of 298 K:
• n is number of moles of electrons.
Effect of Concentration on Cell EMFEffect of Concentration on Cell EMFEffect of Concentration on Cell EMFEffect of Concentration on Cell EMF
QnFRT
EE ln
QlogV0592.0
nEE
Calculate the emf for the Zn-Cu Voltaic cell with [Cu2+] = 1.50 M and [Zn2+] = 0.050 M.
Calculate the emf for the Zn-Cu Voltaic cell with [Cu2+] = 1.50 M and [Zn2+] = 0.050 M.
Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)
QlogV0592.0
nEE
Concentration Cells
Effect of Concentration on Cell EMFEffect of Concentration on Cell EMFEffect of Concentration on Cell EMFEffect of Concentration on Cell EMF
Cell EMF and Chemical Equilibrium• A system is at equilibrium when G = 0.• From the Nernst equation, at equilibrium and 298 K (E =
0 V and Q = Keq):
Effect of Concentration on Cell EMFEffect of Concentration on Cell EMFEffect of Concentration on Cell EMFEffect of Concentration on Cell EMF
0592.0log
log0592.0
0
nEK
Kn
E
eq
eq
• A battery is a self-contained electrochemical power source with one or more voltaic cell.
• When the cells are connected in series, greater emfs can be achieved.
BatteriesBatteriesBatteriesBatteries
Lead-Acid Battery• A 12 V car battery consists of 6 cathode/anode pairs each
producing 2 V.
• Cathode: PbO2 on a metal grid in sulfuric acid:
PbO2(s) + SO42-(aq) + 4H+(aq) + 2e- PbSO4(s) + 2H2O(l)
• Anode: Pb:
Pb(s) + SO42-(aq) PbSO4(s) + 2e-
BatteriesBatteriesBatteriesBatteries
Lead-Acid Battery• The overall electrochemical reaction is
PbO2(s) + Pb(s) + 2SO42-(aq) + 4H+(aq) 2PbSO4(s) + 2H2O(l)
for which
Ecell = Ered(cathode) - Ered(anode)
= (+1.685 V) - (-0.356 V)
= +2.041 V.• Wood or glass-fiber spacers are used to prevent the electrodes
from touching.
BatteriesBatteriesBatteriesBatteries
Alkaline Battery• Anode: Zn cap:
Zn(s) Zn2+(aq) + 2e-
• Cathode: MnO2, NH4Cl and C paste:
2NH4+(aq) + 2MnO2(s) + 2e- Mn2O3(s) + 2NH3(aq) + 2H2O(l)
• The graphite rod in the center is an inert cathode.
• For an alkaline battery, NH4Cl is replaced with KOH.
BatteriesBatteriesBatteriesBatteries
Fuel Cells• Direct production of electricity from fuels occurs in a fuel cell.
• On Apollo moon flights, the H2-O2 fuel cell was the primary source of electricity.
• Cathode: reduction of oxygen:
2H2O(l) + O2(g) + 4e- 4OH-(aq)
• Anode:
2H2(g) + 4OH-(aq) 4H2O(l) + 4e-
BatteriesBatteriesBatteriesBatteries
Corrosion of Iron
• Since Ered(Fe2+) < Ered(O2) iron can be oxidized by oxygen.
• Cathode: O2(g) + 4H+(aq) + 4e- 2H2O(l).
• Anode: Fe(s) Fe2+(aq) + 2e-.• Dissolved oxygen in water usually causes the oxidation of
iron.• Fe2+ initially formed can be further oxidized to Fe3+ which
forms rust, Fe2O3.xH2O(s).
CorrosionCorrosionCorrosionCorrosion
Corrosion of Iron• Oxidation occurs at the site with the greatest
concentration of O2.
Preventing Corrosion of Iron• Corrosion can be prevented by coating the iron with paint
or another metal.• Galvanized iron is coated with a thin layer of zinc.
CorrosionCorrosionCorrosionCorrosion
Preventing Corrosion of Iron• To protect underground pipelines, a sacrificial anode is
added.• The water pipe is turned into the cathode and an active
metal is used as the anode.• Often, Mg is used as the sacrificial anode:
Mg2+(aq) +2e- Mg(s), Ered = -2.37 V
Fe2+(aq) + 2e- Fe(s), Ered = -0.44 V
CorrosionCorrosionCorrosionCorrosion
Electrolysis of Aqueous Solutions– In electrolytic cells the anode is positive and the cathode is
negative. (In galvanic cells the anode is negative and the cathode is positive.)
ElectrolysisElectrolysisElectrolysisElectrolysis
Electroplating• Active electrodes: electrodes that take part in electrolysis.• Example: electrolytic plating.
ElectrolysisElectrolysisElectrolysisElectrolysis