Electrochemistry

Electron transfer reactions are oxidation-reduction or redox reactions.

Results in the generation of an electric current (electricity) or be caused by imposing an electric current.

Therefore, this field of chemistry is often called electrochemistry

Electrochemical Cells An apparatus that allows a redox reaction to occur by

transferring electrons through an external connector.

There are two types Galvanic (Voltaic) cell Electrolytic cell

Galvanic Cells Electrochemical cells produce a current (flow of electrons) as a

result of a redox reaction

Product favored reaction.

Example: Batteries

These cells separate the oxidation & reduction half-reactions, by connecting them to a wire, so the electrons must travel through the wire.

That electron flow, is an electrical current, that can be sent through a circuit which could be part of an electrical device (radio, t.v.)

Basic Terms:

Half Cell -A half reaction & its electrode. -One half of a galvanic cell.

Anode -Electrode that is the source of the negative charge (designated by a

minus (-) sign

-Site of oxidation

Cathode -Electrode that is the source of the positive charge (designated by a positive (+) sign -Site of reduction

Salt Bridge -Connects the two half cells -Maintains the charge neutrality of each half-cell -Allows the flow of the ions with minimal mixing of the half-cell

solutions Gal vanic Cel l sGal vanic Cel l s

Zn

Zn2+ ions

Cu

Cu2+ ions

wire

saltbridge

electrons

AnodeAnodeOxidationOxidation

CathodeCathodeReductionReduction

ZnZn2++2e- Cu2++2e-Cu

AnionsCations

What Is Occurring?

-Zn is oxidized & is the reducing agent

Zn (s) Zn2+(aq) + 2e-

-The Zn electrode is losing mass as the Zn metal is oxidized to Zn2+ ions which go into solution.

-The [Zn2+] solution increases.

-Anions are flowing from the salt bridge toward the anode to balance the positive charge of the Zn2+ ions produced.

-With time Cu plates out onto the Zn metal strip & the Zn metal “disappears”

-Cu is reduced & is the oxidizing agent

Cu2+ (aq) + 2e- Cu (s)

-The Cu electrode is gaining mass as the Cu metal is reduced to Cu2+ ions which go into solution.

-The [Cu2+] solution decreases.

-Cations are flowing from the salt bridge toward the cathode to replace the positive charge of the Cu2+ ions consumed

-The reaction occurs spontaneously

- Electrons always flow anode to cathode.

Way to remember:

Gal vanic Cel l sGal vanic Cel l sA A RRed ed CCatat ate ate AAn n OOx.x.

EEDDUUCCTTIIOONN

AATTHHOODDEE

NNOODDEE

XXIIDDAATTIIOONN

-Line notation -A shorthand way to represent an electrochemical cell without

drawing a picture.

Gal vanic Cel l sGal vanic Cel l s

SolidSolid AqueousAqueous AqueousAqueous SolidSolid

AnodeAnode SaltSaltBridgeBridge

CathodeCathode

-A single line represents a phase change

-A double line represents the salt bridge

-If no solid metal is present, then you need an inert conductor (Pt).

Cu(s) Cu2+(aq) Fe2+(aq),Fe3+(aq) Pt(s)

-Pt is also used when gases are involved.

Gal vanic Cel l sGal vanic Cel l sExamplesExamples

HH++(aq(aq) + 2e) + 2e-- HH22(g)(g)

Zn(sZn(s) ) ZnZn2+2+ + 2e+ 2e--

Zn(sZn(s) Zn) Zn2+2+(aq) (aq) HH++(aq(aq) ) HH22 (g) (g) Pt(sPt(s))

RedRed

OxOx

CatCat

AnAn

Gal vanic Cel l sGal vanic Cel l s2Ti(s)+Sn2Ti(s)+Sn2+2+(aq)(aq) 2 Ti2 Ti++ ((aq)+Snaq)+Sn (s)(s)

SnSn2+2+(aq) + 2e(aq) + 2e-- SnSn (s)(s)2 2 Ti(sTi(s) ) 2 Ti2 Ti++ + 2e+ 2e--

Ti(sTi(s) ) TiTi++(aq(aq) Sn) Sn2+2+(aq) (aq) SnSn (s)(s)

RedRedOxOx

CatCatAnAn

Gal vanic Cel l sGal vanic Cel l sZn(sZn(s) Zn) Zn2+2+(aq) Fe(aq) Fe3+3+(aq),Fe(aq),Fe2+2+(aq) (aq) Pt(sPt(s))

FeFe3+3+(aq) + e(aq) + e-- FeFe2+2+(aq) (s)(aq) (s)

Zn (s) Zn (s) ZnZn2+2+(aq) + 2 e(aq) + 2 e--

Zn(sZn(s) +2 Fe) +2 Fe3+3+(aq)(aq) ZnZn2+2+(aq) +2 Fe(aq) +2 Fe2+2+(aq)(aq)

RedRed

OxOx

CatCat

AnAn

xamples

Cell Potential

Cell Potential, Cell Voltage or Electromotive Force (emf): The driving force pushing the electrons from the anode to the cathode.

Unit is the volt

1 Volt = 1 Joule/Coulomb

Voltmeter: Measures cell potential

Standard Reduction Potential

Cell Potential (E°) can be determined from the standard reduction potential (E°red) for the half reaction.

Reduction potential is the tendency for reduction to occur.

Standard reduction potential (E°) is the voltage associated with a reduction reaction at an electrode when all solutes are 1 M & all gases are at 1 atm.

Standard Hydrogen Electrode (E°=0): This is the reference electrode. Potentials of any other electrode are compared with the SHE.

2e- + 2H+(1 M) 2H2 (1 atm) E°= 0

A positive E° is a spontaneous reduction reaction.

A negative E° is a non-spontaneous reduction reaction or spontaneous oxidation (reverse the reaction.

St andar d Reduct ion St andar d Reduct ion Pot ent ialPot ent ial

HalfHalf-- ReactionReaction EE°° (V)(V)FF22 + 2e+ 2e-- 2F2F-- 2.872.87AuAu3+ 3+ + 3 e+ 3 e-- Au 1.50Au 1.50AgAg++ + e+ e-- AgAg 0.800.80CuCu2+2+ + 2e+ 2e-- Cu 0.34Cu 0.342H2H++ + 2e+ 2e-- HH22 0.000.00NiNi2+2+ + 2e+ 2e-- Ni Ni -- 0.230.23ZnZn2+2+ + 2e+ 2e-- Zn Zn -- 0.760.76AlAl3+ 3+ + 3e+ 3e-- AlAl -- 1.661.66LiLi++ + e+ e-- Li Li -- 3.053.05

Spontaneous Spontaneous ReductionReduction

SHESHENonNon--Spontaneous Spontaneous ReductionReductionSt andar d Reduct ion St andar d Reduct ion

Pot ent ialPot ent ialHalfHalf-- ReactionReaction ℰℰ°° (V)(V)FF22 + 2e+ 2e-- 2F2F-- 2.872.87AuAu3+3+ + 3 e+ 3 e-- Au 1.50Au 1.50AgAg++ + e+ e-- AgAg 0.800.80CuCu2+2+ + 2e+ 2e-- CuCu 0.340.342H2H++ + 2e+ 2e-- HH22 0.000.00Ni Ni NiNi2+2+ + 2e+ 2e-- +0.23+0.23Zn Zn ZnZn2+2+ + 2e+ 2e-- +0.76+0.76Al Al AlAl3+ 3+ + 3e+ 3e-- +1.66+1.66Li Li LiLi++ + e+ e-- +3.05+3.05

Spontaneous Spontaneous ReductionReduction

SHESHE

Spontaneous Spontaneous OxidationOxidation

St andar d Reduct ion St andar d Reduct ion Pot ent ialPot ent ial

HalfHalf-- ReactionReaction EE°° (V)(V)FF22 + 2e+ 2e-- 2F2F-- 2.872.87AuAu3+ 3+ + 3 e+ 3 e-- Au 1.50Au 1.50AgAg++ + e+ e-- AgAg 0.800.80CuCu2+2+ + 2e+ 2e-- Cu 0.34Cu 0.342H2H++ + 2e+ 2e-- HH22 0.000.00NiNi2+2+ + 2e+ 2e-- Ni Ni -- 0.230.23ZnZn2+2+ + 2e+ 2e-- Zn Zn -- 0.760.76AlAl3+ 3+ + 3e+ 3e-- AlAl -- 1.661.66LiLi++ + e+ e-- Li Li -- 3.053.05

Strongest Strongest Oxidizing Agent Oxidizing Agent

(Most Easily (Most Easily Reduced)Reduced)

Strongest Strongest Reducing Agent Reducing Agent (Most Easily (Most Easily Oxidized)Oxidized)

The half reaction with the greater reduction potential (further down on the table) will undergo reduction.

The half reaction with the smaller reduction potential (higher up on the table) will undergo oxidation.

Calculating Cell Potential

E°overall = E°oxidatio + E°reduction

Use the standard reduction potential table to find values.

For the oxidation reaction you flip the reaction and change the sign of E°.

Example: Calculate the cell potential for the following reaction: Zn + Cu2+ Zn2+ + Cu

Split into half reactions & look up their value on the standard reduction potential table.

Cu2+ + 2e- Cu E° = 0.34 V

Zn2+ +2e- Zn E°= - 0.76 V

-Decide which is the oxidation reaction. Flip the reaction & change the sign of E°

Cu2++ 2e- Cu E°red= 0.34 V

Zn Zn2++2e- E°ox =+ 0.76V

E°overall = E°oxidation + E°reduction = 0.76 V + 0.34 V = 1.10 V

Example 2: Calculate the cell potential for the following reaction: Cu + Ag+ Cu2+ + Ag Cu2+ + 2e- Cu E° = 0.34 V

Ag+ + e- Ag E° = 0.80 V

-Decide which is the oxidation reaction. Flip the reaction & change the sign of E°

CuCu2+ + 2e- E°ox= -0.34 V

Ag++ e- Ag E°red = 0.80 V

-Before you calculate E°overall the number of electrons must be equal.

-Multiply Ag’s reaction by 2.

-E° value doe not change.Cu Cu2++ 2e- E°ox= -0.34 V

2Ag++2e- 2Ag E°red= 0.80 V

Cu Cu2++ 2e- E°ox =- 0.34 V

2Ag++2e- 2Ag E°red = 0.80V

E°overall = E°oxidation + E°reduction = -0.34 V + 0.80 V = 0.46 V

Cell Potential & Equilibrium

As a the reaction proceeds, the concentrations of the solutions change, the driving force behind the reaction becomes weaker, and the cell potential eventually reaches zero.

When the cell potential equals zero, the reaction is at equilibrium.

Use the following equation to calculate equilibrium.

Cel l Pot ent ial & Cel l Pot ent ial & Equil ibr iumEquil ibr ium

Use the following equationUse the following equationto calculate equilibrium.to calculate equilibrium.

E°E° cellcell ==0 0.05920.0592

nnlog log KKeqeq

Number of ElectronsNumber of Electrons

Example: The E°cell for the reaction between Zn & Cu is 1.10 V. Calculate Keq.

Zn + Cu2+ Zn2+ + CuCel l Pot ent ial & Cel l Pot ent ial & Equil ibr iumEquil ibr ium

E°E° cellcell ==

0

0.05920.0592nn log log KKeqeq

1.10 V =1.10 V = 0.05920.059222

log log KKeqeq

=== 1.45 x 101.45 x 103737

Nernst Equation

Used to find the cell potential at any moment during a reaction or at conditions other than standard-state.

Ner nst Equat ionNer nst Equat ionEquation:Equation:

E°E° cellcell

0

-- RTRTnFnF

lnln QQEEcellcell ==

R = 8.31 J /mol KR = 8.31 J /mol KT = temperature in KelvinT = temperature in Kelvinn = number of electronsn = number of electronsF = 96,485 C/mol eF = 96,485 C/mol e--

Ner nst Equat ionNer nst Equat ion

0Q=Q=

Q is Reaction QuotientQ is Reaction Quotient

ProductsProductsReactantsReactants

Ner nst Equat ionNer nst Equat ion

0

What is the cell potential What is the cell potential of the following cell at 30°C?of the following cell at 30°C?

E° = 1.10 VE° = 1.10 V

Zn(sZn(s) Zn) Zn2+2+(1x10(1x10-- 55M) CuM) Cu2+2+(0.10M) Cu (s)(0.10M) Cu (s)

-First write the equation:

Zn(s) + Cu2+(aq) Cu(s) + Zn2+(aq)

-Then calculate Q

Ner nst Equat ionNer nst Equat ion

0

Q=Q=

Q is Reaction QuotientQ is Reaction Quotient[Zn[Zn2+2+]][Cu[Cu2+2+]]

==1 x 101 x 10-- 5 5 MM0.10 M0.10 M

== 0.0001 M0.0001 MNer nst Equat ionNer nst Equat ion

Equation:Equation:E°E° cellcell

0

-- RTRTnFnF

lnln QQEEcellcell ==

EEcellcell ==1.10 V1.10 V 8.31×3038.31×3032×96,4852×96,485

lnln 0.00010.0001--

== 1.381.38 VV

Batteries

Bat t er iesBat t er ies

0

Car Battery (Lead storage Car Battery (Lead storage battery)battery)

-- Anode: Anode: PbPb + HSO+ HSO44

-- PbSOPbSO44 + H+ H++ + 2e+ 2e--

-- Cathode: Cathode: PbOPbO2 2 +HSO+HSO44

-- + 3H+ 3H++ + 2e+ 2e-- PbSOPbSO44 + 2H+ 2H22OOBat t er iesBat t er ies

0

Cell: Cell: PbPb+ PbO+ PbO22+2H+2H++ +2 HSO+2 HSO44

-- 2 PbSO2 PbSO44 + 2H+ 2H22OO

2 volts per cell, 6 cells to a2 volts per cell, 6 cells to abattery battery 12 volt battery12 volt battery

Bat t er iesBat t er ies

0

Alkaline BatteryAlkaline Battery

-- Anode: Anode: Zn Zn ZnZn2+2+ + 2e+ 2e--

-- Cathode: Cathode: 2 MnO2 MnO22 + H+ H22O + 2eO + 2e--

MnMn22OO33 + 2OH+ 2OH--

-- 1.5 volts1.5 volts

Bat t er iesBat t er ies

0

Lemon BatteryLemon Battery

-- Anode: Anode: Zn Zn ZnZn2+2+ + 2e+ 2e--

-- Cathode: Cathode: CuCu2+2+ + 2e+ 2e-- CuCu

-- 1.1 volts1.1 volts

Electrolytic cell

Electrical energy is used to produce chemical change.

Non-spontaneous reaction -An external power source is used to force the reaction to occur.

Uses:-Charging (rechargeable) batteries-Electroplating-Producing or purifying metals