Electrochemistry Redox Oxidation States Voltaic Cells Balancing Equations Concentration Cells Nernst Equation CorrosionBatteriesElectrolysis Equilibrium.

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 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).


The Periodic TableThe Periodic TableAlkali metals

Alkaline Earth metals

Halogens

Noble Gases

Chalcogens

Transition Metals

Inner Transition Metals

Ions and Ionic CompoundsIons and Ionic Compounds

Predicting Ionic Charge

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


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 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)


Balance: MnO4- (aq) + C2O4

2-(aq) Mn2+(aq) + CO2(g) (Acidic)

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 ]


CN-(aq) + MnO4-(aq) CNO-(aq) + MnO2(s) [ Basic Solution ]

Solution Key

• 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


Standard Reduction (Half-Cell) Potentials

• Convenient tabulation of electrochemical data.

• Standard reduction potentials, Ered are measured relative to the standard hydrogen electrode (SHE).


) V 2.37- ( Mg(s) - 2e )(Mg -2 aq

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



) V 2.37- ( Mg(s) - 2e )(Mg -2 aq

• 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


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)










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


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):


0592.0log

log0592.0

0

nEK

Kn

E

eq

eq

0592.0log

nEKeq

• 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-


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.


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.


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-


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.


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


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

ElectrochemistryElectrochemistry

RedoxOxidation States

Voltaic CellsBalancingEquations

Cell EMF

ConcentrationCells

EquilibriumNernst

Equation

Corrosion Batteries Electrolysis

oo EFnG


Oxidation is the loss of electrons

Electrochemistry Redox Oxidation States Voltaic Cells Balancing Equations Concentration Cells Nernst Equation CorrosionBatteriesElectrolysis Equilibrium.

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