Transcript Electro Chemistry

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
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 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.
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Photomicrograph of a copper deposit in brass, showing the effect of
dezincification (x50).
Electrochemical Corrosion
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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.
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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 reduced
Reducing 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
2
Zn( s)  Zn (aq)  2e
Cu 2 (aq)  2e  Cu(s)
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 right
 Electrons flow from left to right
 Oxidation on left, reduction on right
 Single vertical = electrode/electrolyte boundary
 Double 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
Fe(s)  2Fe3 (aq)  3Fe2 (aq)
 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
Catho
de
Oxidati
on
Reduct
ion
Connections: cell potential and free
energy
 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
1J  1C 1V
Relating free energy and cell
potential
The Faraday :
F = 96 485 C/mole
Standard conditions (1 M, 1 atm, 25°C)
G  nFE
G  nFE
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 V
Standard potential for
Zn(s) = Zn2+(aq) + 2e : 0.76 V
Where there is no SHE
 In this cell there is no SHE and the measured voltage is 1.10 V
2
2
Zn Zn (aq) C u (aq) Cu
2
2
Zn( s )  Cu (aq)  Zn (aq)  Cu ( s )
Zn(s)  Zn 2 (aq)  2e, E o  0.76V
2
Cu (aq)  2e  Cu(s), E  0.34V
o
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
Zn(s)  2 Ag  (aq)  Zn 2 (aq)  2 Ag (s)
 What is the cell potential?
 The half reactions are:
Ag  (aq)  e  Ag ( s)
Zn( s )  Zn 2 (aq)  2e
 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.
Extensive VS intensive
 Free energy is extensive property so need to multiply by no of moles
involved
G  nFE
 But to convert to E we need to divide by no of electrons involved


G
E 

 E is an intensive property
nF
The Nernst equation
 Working in nonstandard conditions

G  G  RT ln Q

 nFE  nFE  RT ln Q

E  E  RT

nF
ln Q
E  E  0.0592 log Q
n
Summary
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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.
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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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Example
HalfCell Potential for Copper
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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:
From the Nernst equation, with n = 2 and E0 = +0.34 V:
Polarization
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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.
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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
Zn anode
 Local action cell
 A brass fitting connected to a steel pipe
 A bronze propeller in contact with the
steel hull of a ship
Cu cathode
HCl Solution
Differential Temperature Cell
 This 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