Prepared by:Mrs Faraziehan Senusi

PA-A11-7C

Electrochemical Cells

Corrosion & Prevention

Chapter 3Oxidation and Reduction

Oxidation-Reduction Concepts

Voltaic Cell

Electrolytic Cell

Reference: Chemistry: the Molecular Nature of Matter and Change, 6th ed, 2011, Martin S. Silberberg, McGraw-Hill

– Purpose of a voltaic cell is to convert the free energy of a spontaneous reaction into the kinetic energy of electrons moving through an external circuit (electrical energy)

– Electrical energy is proportional to the difference in the electrical potential between the two cell electrodes

Cell Potential, Ecell

or voltage of the cell or electromotive force (emf)

Cell Potential

– Positive Cell Potential – Electrons flow spontaneously from the negative electrode (Anode) to the positive electrode (Cathode)

– Negative cell potential – is associated with a “nonspontaneous” cell reaction

– Cell potential for a cell reaction at equilibrium would be “0”

cellE < 0 for a nonspontaneous process

cellE = 0 for an equilibrium process

cellE > 0 for a spontaneous process

Cell Potential

Units of Cell Potential– The SI (metric) unit of electrical charge is the:

Coulomb (C)– The SI (metric) unit of current is the:

Ampere (A)

– The SI (metric) unit of electrical potential is the:“Volt (V)”

– By definition, the energy released by a potential difference of one volt moving between the anode and cathode of a voltaic cell releases 1 joule of work per coulomb of charge

1 coulomb1 ampere = 1A = 1 C / s

second

1 J 1 J1 volt = 1 C =

C V

Cell Potential

Standard Cell Potential

– Eocell – The potential measured at a specific

temperature (298 K) with no current flowing and all concentrations in their “Standard States”

• 1 atm for gases

• 1 M for solutions

• Pure solids for electrodes

2+ 2+ ocellZn(s) + Cu (aq; 1M) Zn (aq; 1M) + Cu(s) E = 1.10V

When zinc-copper cell diagrammed previously, begins operating under standard state conditions, when [Zn2+] = [Cu2+] = 1M, the cell produce 1.10 V at 298 K.

Standard Cell Potential

Standard Electrode Half-Cell Potentials– Eo

half-cell – Potential associated with a given half-cell reaction

(electrode compartment) when all components are in “Standard States”

– Standard Electrode Potential for a half-cell reaction, whether anode (oxidation) or cathode (reduction) is written as a “reduction”

– Ex.

would be written:

2+ -Cu (aq) + 2e Cu(s) [reduction - cathode] 2+ -Zn(s) Zn (aq) + 2e [oxidation - anode]

2+ - o oCopper cathodeCu (aq) + 2e Cu(s) E (E ) [reduction]

2+ - o oZinc anodeZn (aq) + 2e Zn(s) E (E ) [reduction]

Standard Electrode Half-Cell Potentials

– Electrons flow spontaneously from Anode (negative) to Cathode (positive)

– Cathode must have a more “Positive” Eohalf-cell than the

Anode– For a “positive” Eo

cell

– The standard cell potential is the difference between the standard electrode potential of the “Cathode” (reduction) half-cell and the standard electrode potential of the “Anode” (oxidation) half-cell

– Standard half-cell potentials are “intensive” properties, thus their values do NOT have to be adjusted for stoichiometry (# of moles)

o o ocell cathode (reduction) anode (oxidation)E = (E - E ) > 0

o o ocell copper ZincE = E - E

Standard Electrode Half-Cell Potentials

• What does the voltmeter reading signify??

It signifies the standard cell potential (Ecell) of the voltaic cell.

Eg: The Ecell of the zinc-copper voltaic cell is 1.1 V

The magnitude of a cell’s potential measures the spontaneity of its redox reaction.

E°cell = E°cathode − E°anode

where E°anode is the standard potential at the anode and

E°cathode is the standard potential at the cathode as given

in the table of standard electrode potential.

Standard Cell Potential

Practice Problem

Write out the overall equation for the cell reaction and determine the standard cell potential for the following galvanic cell

[Eo (Ag+/Ag) = 0.80 V; Eo (Ni2+/Ni) = - 0.26 V]

+ -2Ag (aq) + 2e 2Ag(s) [reduction (cathode)] 2+ -Ni(s) Ni (aq) + 2e [oxidation (anode)]

+ 2+Ni(s) + 2Ag (aq) Ni + 2Ag(s)

2+ +Ni(s) Ni (aq) Ag (aq) Ag(s)I I

+ -2Ag (aq) + 2e 2Ag(s) [reduction (cathode)] 2+ -Ni (aq) + 2e Ni(s) [reduction (anode)]

o o ocell Silver NickelE = E - E

ocellE = 0.80 - (-0.26) = 1.06 V

Write both reactions in ‘reduction’ form

= Eo cathode – Eo anode

• To know the standard potential of each electrode (in each half-cell), a reference electrode in which its standard potential (E) is assigned as 0.00V is created.

• This reference electrode uses hydrogen, and is known as the standard hydrogen electrode (SHE).

Standard Hydrogen Electrode

The Standard Hydrogen Electrode– Half-cell potentials are not absolute quantities– The values found in tables are determined relative to a

“Standard”– The Standard Electrode potential is defined as zero

(Eoreference) = 0.00

– The “standard reference half-cell” is a standard “Hydrogen” electrode

– Specially prepared Platinum electrode immersed in a1 M aqueous solution of a strong acid through which H2 gas at 1 atm is bubbled + - o

2 reference2H (aq; 1 M) + 2e H (g); 1 atm) E = 0.00V

Standard Hydrogen Electrode

Standard Hydrogen Electrode

Reference Half-Cell and Unknown Half-Cell– The “Standard” electrode can act as either the “Anode” or

the “Cathode”

– Oxidation of H2 (lose e-) at anode half-cell and reduction of unknown at cathode half-cell

– Reduction of H+ (gain e-) at cathode half-cell and oxidation of unknown at anode half-cell

o o ocell unknown unknownE = E - 0.00 V = E

o o ocell unknown unknownE = 0.00 V - E = - E

Standard Hydrogen Electrode

Eo cell = Eo cathode – Eo anode = Eo unknown – Eo reference

Eo cell = Eo cathode – Eo anode = Eo reference– Eo unknown

Voltmeter reading

Or,

To know the value of Eo zinc from the value of Eocell

Eo cell = Eo cathode – Eo anode = Eo reference– Eo zinc

Eozinc = Eo reference – Eo cell = 0.000 – 0.763 = – 0.763 V

Voltmeter reading

Or,

To know the value of Eo copper from the value of Eocell

Eocell = Eo

cathode – Eoanode = Eo

copper – 0.000

Eocopper = Eo

cell = 0.337 V

• Relative Strength of Oxidizing and Reducing Agents

• The more positive the Eo value, the more readily the reaction occurs– Strength of Oxidizing Agents – Cu2+ > H+ > Zn2+

– Strength of Reducing Agents – Zn > H2 > Cu

• Oxidizing agents decrease in strength as the value of Eo decreases, while the strength of the reducing agents increases as the value of Eo decreases

• Cu2+ is the stronger Oxidizing agent• Zn metal is the stronger Reducing agent

2+ - oCu (aq) + 2e = Cu(s) E = 0.34 V+ - o

22H (aq) + 2e = H (g) E = 0.00 V2+ - oZn (aq) + 2e = Zn(s) E = - 0.76 V

Strength of Oxidizing and Reducing Agents

(Better tendency to undergo reduction)

emf Series– All values are relative to the “standard hydrogen

(reference) electrode– All reactions are written as “reductions”– F2 is strongest oxidizing agent (high, positive Eo)

• Fluorine is very electronegative• It is easily reduced (gain e-) to form weak reducing

agent, F- (reluctant to lose electrons)– Li metal is strongest reducing agent (low, negative Eo)

• has low ionization potential• easily oxidized (loses e-) to form strong oxidizing agent,

Li+ (reluctant to gain electrons)

Table of Standard Electrode Potentials

Increasing strength as reducing agent (oxidation)Table of Standard Electrode Potentials

(The emf Series)

All Values are relative to the “standard hydrogen (reference) electrode

All reactions are written as “reductions”

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Practice Problem

• The Nernst Equation

What if the voltaic cell is NOT at standard states?

The reaction quotient, Q involves a ratio of concentrations or pressures of products to those of reactants, each raised to the power indicated by the coefficient in the balanced equation. The Q expression that is used in the Nernst equation is the thermodynamic reaction quotient; it can include both concentrations and pressures.

Practice Problem

Solution :The zinc–hydrogen cell operated with oxidation at the zinc electrode and reduction at the hydrogen electrode, with a standard cell potential of 0.763 V.

Practice Problem