Chapter 7
Electrochemistry and corrosion
A-Electrochemistry
Electric cells are composed of two electrodes–solid electrical
conductors and at least one electrolyte (aqueous electrical
conductor).
In current cells, the electrolyte is often a moist paste (just enough
water is added so that the ions can move). Sometimes one electrode
is the cell container.
The positive electrode is defined as the cathode and the negative
electrode is defined as the anode.
Electrochemical Reactions
In electrochemical reactions, electrons are transferred from one species to another.
Oxidation Numbers
In order to keep track of what loses electrons and what gains them, we assign
oxidation numbers.
1
Oxidation and Reduction
• A species is oxidized when it loses electrons.
Here, zinc loses two electrons to go from neutral zinc metal to the Zn2+ ion.
A species is reduced when it gains electrons.
Here, each of the H+ gains an electron and they combine to form H2.
H+ oxidizes Zn by taking electrons from it.
Zn reduces H+ by giving it electrons.
Assigning Oxidation Numbers
1. Elements in their elemental form have an oxidation number of 0.
2. The oxidation number of a monatomic ion is the same as its charge.
3. Nonmetals tend to have negative oxidation numbers, although some are
positive in certain compounds or ions.
2
Oxygen has an oxidation number of −2, except in the peroxide ion in which it
has an oxidation number of −1.
Hydrogen is −1 when bonded to a metal, +1 when bonded to a nonmetal.
Fluorine always has an oxidation number of −1.
The other halogens have an oxidation number of −1 when they are negative;
they can have positive oxidation numbers, however, most notably in
oxyanions.
4. The sum of the oxidation numbers in a neutral compound is 0.
5. The sum of the oxidation numbers in a polyatomic ion is the charge on the
ion.
Galvanic Cells:
Produces electrical current spontaneous by chemical reactions à Battery
Cu2+ + 2e- D Cu E0 = +0.34 V
Zn2+ + 2e- D Zn E0 = −0.76V
Cu2+ + Zn D Cu + Zn2+ E0 = 0.34 – (-0.76) = 1.10 V
Parts of the voltaic or galvanic cell…
Anode à the electrode where oxidation occurs
After a period of time, the anode may appear to become smaller as it falls into
solution.
Cathode à the electrode where reduction occurs
After a period of time it may appear larger, due to ions from solution plating onto
it.
Salt Bridge à a device used to maintain electrical neutrality in a galvanic cell.
This may be filled with neutral salt or it may be replaced with a porous barrier.
3
Electron Flow à always from anode to cathode (through the wire).
Anode Cathode
The diagram to the right illustrates what really happens when a Galvanic cell is
constructed from zinc sulfate and copper (II) sulfate using the respective metals as
electrodes.
Zn/Zn+2 is the anode
Zn Zn+2 + 2e- E° = +0.76 V
Cu/Cu+2 is the cathode
Cu+2 + 2e- Cu E° = 0.34 V
4
Zn Zn+2 + 2e- E° = +0.76 V
Cu+2 + 2e- Cu E° = 0.34 V
Cu+2 + Zn Cu + Zn+2 E°cell = 1.10 V
Shorthand Notation
Electromotive Force (emf)
• Water only spontaneously flows one way in a waterfall.
• Likewise, electrons only spontaneously flow one way in a redox reaction—
from higher to lower potential energy.
5
• The potential difference between the anode and cathode in a cell is called the
electromotive force (emf).
• It is also called the cell potential, and is designated Ecell.
Batteries
A battery is a group of galvanic cells connected in series
The potentials of the individual cells add to give the total battery potential
Secondary cells can be recharged by adding electricity
Figure 1: One of the Cells in a 12V Lead Storage Battery
Lead –acid battery:
If a number of cells are connected in series, the arrangement is called a battery.
The lead storage battery is one of the most common batteries that is used in the
automobiles. A 12V lead storage battery is generally used, which consists of six
cells each providing 2V. Each cell consists of a lead anode and a grid of lead
packed with lead oxide as the cathode. These electrodes are arranged alternately,
6
separated by a thin wooden piece and suspended in dil.H2SO4 (38%), which acts as
an electrolyte (Fig. 1).Hence it is called Lead-acid battery.
Anode: Pb
Cathode: PbO2
Electrolyte: H2SO4 (38%)
EMF=2V
To increase the current output of each cell, the cathode and the anode plates are
joined together, keeping them in alternate positions. The cells are connected
parallel to each other (anode to anode and cathode to cathode). The cell is
represented as
Pb | PbSO4 (s), H2SO4 (aq.) | PbSO4 (s), Pb
In the process of discharging, i.e. when battery produces current, the reactions at
the electrodes are as follows:
At anode:
Pb à Pb+2 + 2e-
Pb (s) + SO4 (aq.) à PbSO4 (s)
At cathode:
PbO2 (s) + SO4 (aq.) + 4H+ (aq.) + 2e– à PbSO4 (s) + 2H2O
Therefore, overall reaction is
Pb (s) + PbO2 (s) + 2H2SO4 (aq.) à 2PbSO4 (s) + 2H2O
During discharging the battery, H2SO4 is consumed, and as a result, the density of
H2SO4 falls; when it falls below 1.20 g/cm3, the battery needs recharging. In
Discharging, the cell acts as a voltoic cell where oxidation of lead occurs.
During recharging, the cell is operated like an electrolytic cell, i.e. electrical energy
is supplied to it from an external source. The electrode reactions are the reverse of
those that occur during discharge.
7
PbSO4 (s) + 2e– àPb (s) + SO4– – (aq.)
PbSO4 (s) + 2H2O à PbO2 (s) + 2H2SO4 + 2e–
2PbSO4 (s) + 2H2O à Pb (s) + PbO2 (s) + 2H2SO4 (aq.)
During this process, lead is deposited at the cathode, PbO2 is formed at the anode
and H2SO4 is regenerated in the cell.
Advantages: Lead acid batteries are used for supplying current to railways, mines,
laboratories, hospitals, automobiles, power stations, telephone exchange, gas
engine ignition, Ups (stand-by supplies). Other advantages are its rechargeability,
portability and Its relatively constant potential & low cost.
Disadvantages: Use of Conc.H2SO4 is dangerous; Use of lead battery is fragile.
B. CORROSION
What is corrosion?
Corrosion is a natural event
It represents a return of metals to their more natural state as minerals (oxides).
Metal Corrosion
The destruction of a material by chemical or electrochemical reaction to its
environment”
Typically a transfer of electrons from one metal to another through an
Oxidation-Reduction Reaction.
Oxidation - Reduction Reaction
Anodic metal gives up electrons (oxidation)
Cathodic metal accepts electrons (reduction)
Or gases accept electrons (reduction)
8
Al→ Al+ 3+3e−Fe→Fe+2+2e−
Cu2 ++2e−→Cu
2 H++2e−→H2( gas )
Acceleration of Corrosion
Physical Characteristics
time of exposure (more time, more corrosion)
exposed area (less, increases corrosion rate)
Environmental Characteristics
acidic environment
sulfur gas environment
temperature (high temps, more corrosion)
moisture (oxygenated moisture)
Passivation
Refers to a material becoming "passive" that is, being less affected by
environmental factors such as air and water. Passivation involves a shielding outer-
layer of base material,
Types of Passivation
A protective film in oxidizing atmospheres
Chromium, nickel, titanium, aluminum
Metal oxide layer adheres to parent metal
Barrier against further damage
Self-healing if scratched
Forms of Corrosion
Uniform corrosion of a single metal
Usually an electrochemical reaction.
Relatively slow and predictable.
Rusting of exposed steel, tarnished silver.
Easily corrected with coatings and regular maintenance.
9
Galvanic Corrosion
2 dissimilar metals, electrolyte, electrical connection and oxygen.
Pitting Corrosion
Localized corrosion forming holes.
Difficult to initially detect.
Reinforcement Corrosion
Corrosion Products
Fe + 2OH = Fe(OH)2
Oxidation of Fe(OH)2
Fe(OH)3 (rust)
Corrosion of Metals in Concrete
Concrete is Normally Highly Alkaline
Protects Steel from Rusting if Properly Embedded
10
If Corrosion Occurs, the Reaction Products are greater in volume than the
original steel.
Corrosion Initiation and Rate Depends On
Amount of Concrete Cover, Quality of Concrete
Details of Construction, & Exposure to Chlorides
Avoiding Corrosive Situations
Choose couple metals close on the galvanic series
Use large anode, and small cathode areas
Electrically insulate dissimilar metals
Connect a more anodic metal to the system
Corrosion Prevention
Coatings
Barrier films
Inhibitive Pigments
Sacrificial treatments
11
Paint
Active Cathodic Protection
Cathodic protection is an electrochemical means of corrosion control in which the
oxidation reaction in a galvanic cell is concentrated at the anode and suppresses
corrosion of the cathode in the same cell. This is achieved by placing a more easily
corroded metal to act as the anode of the electrochemical cell in contact with the
12