CHM 120

CHAPTER 21Electrochemistry:

Chemical Change and Electrical Work

Dr. Floyd BeckfordLyon College

REVIEWREVIEW• Oxidation: the loss of electrons by a species,

leading to an increase in oxidation number of one

or more atoms

• Reduction: the gain of electrons by a species,

leading to an decrease in oxidation number of one

or more atoms

• Oxidizing agents: the species that is reduced in a

redox reaction

• Reducing agents: the species that

is oxidized in a redox reaction

In acidic solution: add H+ or H2O only

In basic solution: add OH- or H2O only

  To balance O   To balance H

In acidic solution:

Add H2O and then

Add H+

  For each O needed

  For each H needed

In basic solution

1. add 2 OH- to the side needing O

and

and then

1. add 1 H2O to the

side needing H

and  2. add 1 H2O to the other

side

  2. add 1 OH- to the other

side 

THE HALF-REACTION METHODTHE HALF-REACTION METHOD• This method breaks the overall reaction into its

two components – half-reactions

• Each half-reaction is balanced separately and

then added

• Use the following guidelines to help

1. Write as much of the unbalanced net ionic

equation as possible

2. Decide which atoms are oxidized and which are

reduce – write the two unbalanced half-reactions

3. Balance by inspection all atoms in each half-

reaction except H and O

4. Use the rules mentioned previously to balance

H and O in each half-reaction

5. Make equal the number of electrons involved in

both half-reactions

• Take a look at the breathalyzer reaction

H+(aq) + Cr2O72-(aq) + C2H5OH(l)

Cr3+(aq) + C2H4O(l) + H2O(l)

Balance the following net ionic equation in basicsolution.

MnO4-(aq) + SO3

2-(aq) MnO42-(aq) + SO4

2-(aq)

ELECTROCHEMISTRYELECTROCHEMISTRY

• Deals with chemical changes produced by an

electric current and with the production of

electricity by chemical reactions

• All electrochemical reactions involve transfer

of electrons and are redox reactions

• EChem reactions take place in electrochemical

cell (an apparatus that allows a reaction to

occur through an external conductor)

ELECTROCHEMICAL CELLSELECTROCHEMICAL CELLS

Two types:

1. Electrolytic cells: - these are cells in which an

external electrical source forces a

nonspontaneous reaction to occur

2. Voltaic cells: - also called galvanic cells. In

these cells spontaneous chemical reactions

generate electrical energy and supply it to an

external circuit

• Electric current enters and exits the cell by

electrodes - electrodes are surfaces upon which

oxidation or reduction half-reactions occur

• Inert electrodes: - electrodes that don’t react

• Two kinds of electrodes:

1. Cathode: - electrode at which reduction

occurs (electrons are gained by a species)

2. Anode: - electrode at which oxidation occurs

(as electrons are lost by some species)

VOLTAIC CELLSVOLTAIC CELLS

• Cells in which spontaneous reactions produces

electrical energy

• The two half-cells are separated so that electron

transfer occurs through an external circuit

• Each half-cell contains the oxidized and reduced

forms of a species in contact with each other

• Half-cells linked by a piece of wire and a salt

bridge

• A salt bridge has three functions:

1. It allows electrical contact between the two

half-cells

2. It prevents mixing of the electrode solutions

3. It maintains electrical neutrality in each

half-cell as ions flow into and out of the salt

bridge

• Point 2 is important – no current would flow if

if both solutions were in the same cell

• Point 3 is also important – anions flow into the

oxidation half-cell to counter the build-up

of positive charge

• Current flow spontaneously from negative to

the positive electrode

• In all voltaic cells the anode is negative and the

cathode is positive

• In voltaic cells, voltage drops as the reaction

proceeds. When voltage = 0, the reaction is at

equilibrium

Zn Zn2+(1.0 M) Cu2+(1.0 M) Cu

Electrode

Salt bridge

Species (withconcentrations) in contact with electrodes

The Silver-Copper cell

• Composed of two half-cells:

1. A strip of copper immersed in 1 M CuSO4

2. A strip of silver immersed in 1 M AgNO3

• Experimentally we see:

: - Initial voltage is 0.46 volts

: - The mass of the copper electrode decreases

: - The mass of the silver electrode increases

: - [Cu2+] increases and [Ag+] decreases

Cu Cu2+ + 2e- (oxidation, anode)

2(Ag+ + e- Ag) (reduction, cathode)

2Ag+ + Cu Cu2+ + Ag (Overall cell reaction)

Cu |Cu2+(1.0 M) ||Ag+(1.0 M) | Ag

• Notice that in this case the copper electrode is

the anode

STANDARD ELECTRODE POTENTIALSSTANDARD ELECTRODE POTENTIALS

• Associated with each voltaic cell is a potential

difference called the cell potential, Ecell

• E measures the spontaneity of the cell’s redox

reaction

• Higher (more positive) cell potentials indicate a

greater driving force for the reaction as written

• All electrode potentials are measured versus the

Standard Hydrogen Electrode (SHE): E° = 0.00 V

• The E°cell calculated is for the cell operating

under standard state conditions

• For electrochemical cell standard conditions

are:

-solutes at 1 M concentrations

- gases at 1 atm partial pressure

- solids and liquids in pure form

• All at some specified temperature, usually 298 K

• The electrode potential for each half-reaction is

written as a reduction process

• The more positive the E° value for a half-

reaction the greater the tendency for the reaction

to proceed as written

• The more negative the E° value, the more likely

is the reverse of the reaction as written

Prediction of Spontaneity

1. First write the HR equation with the more

positive (less negative) E° for the reduction along

with its potential

2. Write the other HR as an oxidation and include

its oxidation potential

3. Balance the electron transfer

4. Add the reduction and oxidation HR and add

the corresponding electrode potentials to get the

overall cell potential, E°cell

• Important points to note:

1. E° for oxidation half-reactions are equal to

but opposite in sign to reduction half-reactions

2. Half-reaction potentials are the same

regardless of the species’ stoichiometric

coefficient in the balanced equation

E°cell > 0 Forward reaction is spontaneous

E°cell < 0 Backward reaction is spontaneous

E°E°cellcell, , G° and KG° and K

• From thermodynamics, we know that,

G° = -RT lnK

• We can relate E°cell to free energy for that cell

G° = -nFE°cell

n = number of moles of e-

So -nFE°cell = -RT lnK and

E°cell = (RT/nF) lnK

Forward reaction

G K Ecell

Spontaneous - > 1 +

Equilibrium 0 1 0

Non- spontaneous

+ < 1 -

(Standard state

conditions)

• Under nonstandard

conditions

G = -nFEcell

THE NERNST EQUATIONTHE NERNST EQUATION

• Usually concentrations of reactants differ from

one another and also change during the course

of a reaction

• As cell reaction proceeds, cell voltage drops so

that E°cell is different from Ecell

• E°cell and Ecell are related by the Nernst

Equation

Ecell = E°cell - (RT/nF) lnQ

Ecell = E°cell - (RT/nF) lnQ

E = potential under the nonstandard conditions

E° = standard potential

R = gas constant, 8.314 J/mol.K

T = absolute temperature

n = number of moles of electrons transferred

F = faraday, 96,485 J/V.mol e-

Q = reaction quotient

BATTERIESBATTERIES

• Two type of batteries:

: - Primary batteries cannot be “recharged”

• Once all the chemicals are consumed there is

no more chemical reaction

: - Secondary batteries can be regenerated

• Most common example is the lead storage

battery used to power automobiles

The Lead Storage BatteryThe Lead Storage Battery

• Composed of two alternating groups of Pb

plates; one group contains pure lead (anode) and

the other group contains PbO2 (cathode)

• The plates are immersed in 40 % sulfuric acid

• During discharge

Pb Pb2+ + 2e- (oxidation)

Pb2+ + SO42- PbSO4 (precipitation)

Net: Pb + SO42- PbSO4 + 2e- (anode)

• At the cathode

PbO2 + 4H+ + 2e- Pb2+ + 2H2O (reduction)

Pb2+ + SO42- PbSO4 (precipitation)

• Net reaction:

PbO2 + 4H+ + SO42- + 2e- PbSO4 + 2H2O

• Adding the HR for the two half-cells, gives

Pb + PbO2 + 4H+ + 2SO42- 2PbSO4 + 2H2O

E°cell = 2.041 V

• The battery can be recharged

Fuel Cells

• These are galvanic cells in which the reactants

are continuously supplied to the cell and the

products are continuously removed

• Best known example is the hydrogen-oxygen

fuel cell

• Hydrogen is fed into the anode compartment

and oxygen into the cathode compartment

• Oxygen is reduced at the cathode – porous

carbon doped with metallic catalysts

• At the anode hydrogen is oxidized to water

Anode: 2H2(g) + 4OH-(aq) 4H2O(l) + 4e-

Cathode: O2(g) + 2H2O(l) + 4e- 4OH-(aq)

Overall: 2H2(g) + O2(g) 2H2O(g)

CORROSIONCORROSION

• Ordinary corrosion is a redox process in

which metals are oxidized by oxygen in the

presence of moisture

• A point of strain on the surface of the metal

acts as an anode

• Areas on the metal surface exposed to air

serves as cathodes

Anode: Fe(s) Fe2+(aq) + 2e-

Cathode: O2(g) + 4H+(aq) + 4e- 2H2O(l)

4Fe(s) + O2(g) + 4H+(aq)

4Fe2+(aq) + 2H2O(l)

2Fe2+(aq) + 4H2O(l) Fe2O3•H2O(s) + 6H+

Rust

• Al also undergo corrosion – initial oxidation

is stopped by a layer of Al2O3

Corrosion prevention

1. Plating a metal with a thin layer of a less easily

oxidized metal

2. Allow a protective film to form naturally on the

surface of the metal

3. Galvanizing – coating the metal with zinc

4. Cathodic protection – connecting the metal to

a “sacrificial anode”

ELECTROLYTIC CELLSELECTROLYTIC CELLS

• Cells in which an electric current causes a

nonspontaneous reaction to occur – one common

process is called electrolysis

• In electrolytic cells the anode is the positive

electrode and the cathode is the negative

electrode

• Still : Anode = oxidation; cathode = reduction

The Down’ Cell: Electrolysis of molten NaCl

• Using graphite inert electrodes the following

observations are made

1. Chlorine, Cl2, is liberated at one electrode

2. Sodium metal forms at the other electrode

• Explanation

1. Chlorine is produced at the anode by the

oxidation of Cl- ions

2. Metallic sodium is formed by reducing Na+

ions at the cathode

• Electrons used at the cathode are

reproduced at

the anode

• The reaction is nonspontaneous and

electricity

is used to force the reaction to occur

2Cl- Cl2(g) + 2e-(oxidation, anode HR)

2(Na+ + e- Na(l) (reduction, cathode HR)

2Na+ + 2Cl- 2Na(l) + Cl2(g) Overall cell rxn.

________________________________________

Electrolysis of aqueous sodium chloride

• In an EChem cell containing aqueous NaCl

: - H2 gas is liberated at one electrode

: - Cl2 gas is liberated at the other electrode

: - Solution at the cathode is basic

• Rationalization

: - Chloride ions are oxidized at the anode and

H2O is reduced at the cathode

2Cl- Cl2 + 2e- (oxidation, anode)

2H2O + 2e- 2OH- + H2 (reduction, cathode)

2H2O + 2Cl- 2OH- + H2 + Cl2 Overall

• Sodium metal is more active than hydrogen

metal and liberates H2 from solution

• The hydroxide ions are responsible for the

basicity around the cathode

FARADAY’S LAWFARADAY’S LAW

• States that the amount of substance that

undergoes oxidation or reduction at each

electrode during electrolysis is directly

proportional to the amount of electricity that

passes through the cell

• One faraday = the amount of electricity that

reduces or oxidizes 1 equivalent of a substance

• One equivalent of any substance is

the amount

of that substance that supplies or

consumes one

mole of electrons

1F = 1mole of electrons

= 6.022 x 1023 e-

= 96,485 C

C = It

C= charge passed; I = current;

t = time (in seconds)