Chemical vs. Electrochemical Reactions

Chemical reactions are those in which elements are

added or removed from a chemical species.

Electrochemical reactions are chemical reactions in

which not only may elements may be added or removed

from a chemical species but at least one of the species

undergoes a change in the number of valance electronS.

Corrosion processes are electrochemical in nature.

Chemical corrosion: Removal of atoms from a material by virtue of the solubility or chemical reaction between the material and the surrounding liquid.

EXAMPLES: Dezincification: A special chemical corrosion process by which both

zinc and copper atoms are removed from brass, but the copper is replated back onto the metal.

Graphitic corrosion: A special chemical corrosion process by which iron is leached from cast iron, leaving behind a weak, spongy mass of graphite.

2Chemical Corrosion

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Photomicrograph of a copper deposit in brass, showing the effect of dezincification (x50).

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Electrochemical corrosion - Corrosion produced by the development of a current in an electrochemical cell that removes ions from the material.

Electrochemical cell - A cell in which electrons and ions can flow by separate paths between two materials, producing a current which, in turn, leads to corrosion or plating.

Oxidation reaction - The anode reaction by which electrons are given up to the electrochemical cell.

Reduction reaction - The cathode reaction by which electrons are accepted from the electrochemical cell.

Electrochemical Corrosion

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The components in an electrochemical cell: (a) a simple electrochemical cell and (b) a corrosion cell between a steel water pipe and a copper fitting.

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The anode and cathode reactions in typical electrolytic corrosion cells: (a) the hydrogen electrode, (b) the oxygen electrode, and (c) the water electrode.

Electrochemistry

Thermodynamics at the electrode

Redox (Review)

Oxidation is...Loss of electrons

Reduction is...Gain of electrons

Oxidizing agents oxidize and are reducedReducing agents reduce and are oxidized

Redox Review (Cu-zn)

Zn displaces Cu from CuSO4(aq)

In direct contact the enthalpy of reaction is dispersed as heat, and no useful work is done

Redox process: Zn is the reducing agent Cu2+ is the oxidizing agent

eaqZnsZn 2)()( 2

)(2)(2 sCueaqCu

Separating the combatants

Each metal in touch with a solution of its own ions External circuit carries electrons transferred during the redox process A “salt bridge” containing neutral ions completes the internal circuit. The energy released by the reaction in the cell can perform useful work –

like lighting a bulb

Labelling the parts

Cell notation

Anode on left, cathode on rightElectrons flow from left to rightOxidation on left, reduction on rightSingle vertical = electrode/electrolyte boundaryDouble vertical = salt bridge

Anode:Zn →Zn2+ + 2e

Cathode:Cu2+ + 2e →Cu

Odes to a galvanic cell

Cathode Where reduction occurs Where electrons are

consumed Where positive ions migrate

to Has positive sign

Anode Where oxidation occurs Where electrons are

generated Where negative ions migrate

to Has negative sign

The role of inert electrodes

)(3)(2)( 23 aqFeaqFesFe

Not all cells start with elements as the redox agents Consider the cell

Fe can be the anode but Fe3+ cannot be the cathode. Use the Fe3+ ions in solution as the “cathode” with an inert metal such as Pt

Anode Cathode

Oxidation

Reduction

Connections: cell potential and free energy

VCJ 111

The cell in open circuit generates an electromotive force (emf) or potential or voltage. This is the potential to perform work

Energy is charge moving under applied voltage

Relating free energy and cell potential

nFEG

The Faraday :

F = 96 485 C/mole

Standard conditions (1 M, 1 atm, 25°C)

nFEG

Standard Reduction Potentials

The total cell potential is the sum of the potentials for the two

half reactions at each electrode

Ecell = Ecath + Ean

From the cell voltage we cannot determine the values of either

– we must know one to get the other

Enter the standard hydrogen electrode (SHE)

All potentials are referenced to the SHE (EH=0 V)

Unpacking the SHE

The SHE consists of a Pt electrode in contact with H2(g) at 1 atm in a solution of 1 M H+(aq).

The voltage of this half-cell is defined to be 0 V.

An experimental cell containing the SHE half-cell with other half-cell gives voltages which are the standard potentials for those half-cells

Ecell = 0 + Ehalf-cell

Zinc half-cell with SHE

Cell measures 0.76 VStandard potential for

Zn(s) = Zn2+(aq) + 2e : 0.76 V

Where there is no SHE

CuaquCaqZnZn )()( 22

In this cell there is no SHE and the measured voltage is 1.10 V

VEeaqZnsZn o 76.0,2)()( 2

)()()()( 22 sCuaqZnaqCusZn

VEsCueaqCu o 34.0),(2)(2

Standard reduction potentials

Any half reaction can be written in two ways:Oxidation:

M = M+ + e (+V)Reduction:

M+ + e = M (-V)Listed potentials are standard reduction potentials

Applying standard reduction potentials

Consider the reaction

What is the cell potential? The half reactions are:

E° = 0.80 V – (-0.76 V) = 1.56 V

NOTE: Although there are 2 moles of Ag reduced for each mole of Zn oxidized, we do not multiply the potential by 2.

)(2)()(2)( 2 sAgaqZnaqAgsZn

eaqZnsZn 2)()( 2 )()( sAgeaqAg

Extensive VS intensive

nFGE

Free energy is extensive property so need to multiply by no of moles involved

But to convert to E we need to divide by no of electrons involved

E is an intensive property

nFEG

The Nernst equation

QRTGG ln

Working in nonstandard conditions

QnFRTEE ln

QRTnFEnFE ln

QnEE log0592.0

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Electrode potential - Related to the tendency of a material to corrode. The potential is the voltage produced between the material and a standard electrode.

emf series - The arrangement of elements according to their electrode potential, or their tendency to corrode.

Nernst equation - The relationship that describes the effect of electrolyte concentration on the electrode potential in an electrochemical cell.

Faraday’s equation - The relationship that describes the rate at which corrosion or plating occurs in an electrochemical cell.

Summary

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The half-cell used to measured the electrode potential of copper under standard conditions. The electrode potential of copper is the potential difference between it and the standard hydrogen electrode in an open circuit. Since E0 is great than zero, copper is cathodic compared with the hydrogen electrode.

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Suppose 1 g of copper as Cu2+ is dissolved in 1000 g of water to produce an electrolyte. Calculate the electrode potential of the copper half-cell in this electrolyte.

The atomic mass of copper is 63.54 g/mol.

Example 22.1 SOLUTION

From chemistry, we know that a standard 1-M solution of Cu2+ is obtained when we add 1 mol of Cu2+ (an amount equal to the atomic mass of copper) to 1000 g of water. The atomic mass of copper is 63.54 g/mol. The concentration of the solution when only 1 g of copper is added must be:

Example

HalfCell Potential for Copper

From the Nernst equation, with n = 2 and E0 = +0.34 V:

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Polarization - Changing the voltage between the anode and cathode to reduce the rate of corrosion.

– Activation polarization is related to the energy required to cause the anode or cathode reaction

– Concentration polarization is related to changes in the composition of the electrolyte

– Resistance polarization is related to the electrical resistivity of the electrolyte.

Polarization

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Photomicrograph of intergranular corrosion in a zinc die casting. Segregation of impurities to the grain boundaries produces microgalvanic corrosion cells (x50).

Corrosion Cells

Galvanic cell (Dissimilar electrode cell) – dissimilar metals

Salt concentration cell – difference in composition of aqueous

environment

Differential aeration cell – difference in oxygen concentration

Differential temperature cell – difference in temperature

distribution over the body of the metallic material

Dissimilar Electrode Cell

When a cell is produced due to two

dissimilar metals it is called

dissimilar electrode cell

Dry cell

Local action cell

A brass fitting connected to a steel

pipe

A bronze propeller in contact with the

steel hull of a ship

Zn anode

HCl Solution

Cu cathode

Differential Temperature CellThis is the type of cell when two identical electrodes are

immersed in same electrolyte, but the electrodes are immersed into solution of two different temperatures

This type of cell formation takes place in the heat exchanger equipment where temperature difference exists at the same metal component exposed to same environment

For example for CuSO4 electrolyte & Cu electrode the electrode in contact with hot solution acts as cathode.

Salt Concentration Cell

Differential Aeration Cell

Corrosion at the bottom of the electrical poles

Local Action Cell