The document discusses corrosion, including its types, why we prevent it, and several related concepts from electrochemistry. It defines corrosion as the destruction of metals by chemical and electrochemical attack from the environment. There are two main types: chemical corrosion from oxygen and electrochemical corrosion, which occurs in the presence of moisture and produces an electrochemical cell. The document outlines several concepts to understand and control corrosion, including the Pilling-Bedworth rule, electrochemical series, Nernst equation, standard electrodes, and Pourbaix and Ellingham diagrams.

2.  What is corrosion ,its types & Why we prevent corrosion?
 Pilling bedworth rule.
 Electrochemical series & thermodynamic principles.
 Nernst equation & electrode potential of metals.
 Standard electrodes & reference electrode.
 E.M.F. & Galvanic series.
 Pourbaix diagram. & its importance for iron,aluminium,magnesium.
 Ellingham diagram.

3.  The process of destruction of metals & its alloys by chemical &
electrochemical attack through its environment starting from its surface is
called as *CORROSION*.
 E.g. Rusting of iron :Fe2O3*2H2O, Reddish brown color rust:
Fe2O3*3H2O,Green film of basic carbonate: CuCO3*Cu(OH)2.
 Generally metal undergo corrosion & convert into
oxides,hydroxide,carbonates,Sulphides etc.
 Another definition: It is the process by which the metal have tendency to go
back their combined state.
 N.B.: Metals are stable in their combined state(less energy , stable) ; not in
their elemental state(more energy , less stable ). So metal have tendency to
get back its original combined state.

4. 1. Chemical corrosion or Dry corrosion : It is the action of oxygen on metal
surface at high or low tempt. in the absence of moisture .
e.g. 2M 2M(n+) + 2ne- ( loss of e-) …..(1)
n/2 O2 + 2ne- n O(2-) ( gain of e-) …..(2)
--------------------------------------------
2M + n/2 O2  2M(n+) + n O(2-)
 2. Electro chemical corrosion or Wet corrosion : Wet Corrosion is the most
common form of corrosion. It occurs with the presence of moisture.
 It will occur if an “electrochemical cell” is produced.
 An electrochemical cell consists of an Anode, a Cathode, a Connection, and
an Electrolyte

6.  It may be stable : (1) non porous : It can't allow for the further corrosion.
E.g.: Al, Sn, Pb, Cu, Pt. (2) porous : It allow for the further corrosion. E.g. :
Na2O, K2O, Fe2O3. & all alkali earth metals.
 It may be unstable : (Au , Ag , Pt ) : The unstable metallic oxide layer
decomposes itself soon & it form unstable metallic oxide.
 It may be volatile : a volatile metallic oxide layer is formed . It undergo
further corrosion .

7.  Due to corrosion , thickness of metal decreases & hence loss of mechanical
strength takes place. There is a structural failure ( cracks ,voids ) also takes
place. So collapse of material takes place .
 So there is a reduction in the value of the material .
 It is estimated that 40% of failure in electronic equipments is due to metallic
corrosion.
 India has a tropical climates . So corrosion prob. in India is more than the cold
country.
 The loss of estimated is 250 cr per yr. in INDIA. The money spends for its
prevention is 50-70 cr .per yr. .
 So it is essential to know about the cause of corrosion & do its prevention.

8.  In case of stable oxide , to decide whether the oxides is porous or non porous
; the following pilling bedworth rule decides:
 According to this rule , * The smaller the specific volume ratio ( volume of
metal oxide / volume of metal ) , greater is the oxidation corrosion.
 Oxidation corrosion is inversely proportional to specific volume .
 It says that : (1) if the volume of metal oxide is less than its metal , then
oxide is porous . Hence further corrosion takes place .(2) if the volume of
metal oxide is at least slightly greater than its metal , then it becomes non
porous .So it act as a protective layer.

9.  I f we arrange the electrodes according to ascending series of their electrode
potential , we get a series of electrodes known as electrochemical series.

11.  It helps to understand the electrochemical behavior of corrosion reaction .
 It also gives how to control corrosion.
 Let us contract a reversible voltage cell or galvanic cell.

12.  In electrochemistry, the Nernst equation is an equation that relates the
reduction potential of an electrochemical reaction (half-cell or full cell reaction)
to the standard electrode potential, temperature, and activities (often
approximated by concentrations) of the chemical species undergoing reduction
and oxidation. It is the most important equation in the field of electrochemistry.
It is named after the German physical chemist who first formulated it, Walther
Nernst.[ It is clear that equation of reactant increases with increasing in
electrode potential & vice versa.

13.  In an electrochemical cell, an electric potential is created between two
dissimilar metals. This potential is a measure of the energy per unit charge
which is available from the oxidation/reduction reactions to drive the
reaction. It is customary to visualize the cell reaction in terms of two half-
reactions, an oxidation half-reaction and a reduction half-reaction.
 Reduced species -> oxidized species + ne-Oxidation at anode
 Oxidized species + ne- -> reduced species Reduction at cathode

14.  The cell potential (often called the electromotive force or emf) has a contribution
from the anode which is a measure of its ability to lose electrons - it will be called
its "oxidation potential". The cathode has a contribution based on its ability to
gain electrons s, its "reduction potential". The cell potential can then be written
 Ecell = oxidation potential + reduction potential
 If we could tabulate the oxidation and reduction potentials of all available
electrodes, then we could predict the cell potentials of voltaic cells created from
any pair of electrodes. Actually, tabulating one or the other is sufficient, since the
oxidation potential of a half-reaction is the negative of the reduction potential for
the reverse of that reaction. Two main hurdles must be overcome to establish such
a tabulation
 The electrode potential cannot be determined in isolation, but in a reaction with
some other electrode.
 The electrode potential depends upon the concentrations of the substances, the
temperature, and the pressure in the case of a gas electrode.

15.  A Standard Hydrogen Electrode (SHE) is an electrode that scientists use for
reference on all half-cell potential reactions. The value of the standard
electrode potential is zero, which forms the basis one needs to calculate cell
potentials using different electrodes or different concentrations. It is
important to have this common reference electrode just as it is important for
the International Bureau of Weights and Measures to keep a sealed piece of
metal that is used to reference the S.I. Kilogram
 What is a SHE made of?
 SHE is composed of a 1.0 M H+(aq) solution containing a square piece
of platinized platinum (connected to a platinum wire where electrons can be
exchanged) inside a tube. During the reaction, hydrogen gas is then passed
through the tube and into the solution causing the reaction:
 2H+(aq) + 2e- <==> H2(g).
 Platinum is used because it is inert and does not react much with hydrogen.

17.  Calomel electrode: This reference electrode consists of a mercury and
mercury-chloride molecules. ( used as second electrode)
 This electrode can be relatively easier to make and maintain compared to the
SHE. It is composed of a solid paste of Hg2Cl2 and liquid elemental mercury
attached to a rod that is immersed in a saturated KCl solution.
 It is necessary to have the solution saturated because this allows for the
activity to be fixed by the potassium chloride and the voltage to be lower and
closer to the SHE. This saturated solution allows for the exchange of chlorine
ions to take place. All this is usually placed inside a tube that has a porous
salt bridge to allow the electrons to flow back through and complete the
circuit.
 It contains mercury which is much more health hazards .

18.  Electromotive force, abbreviated emf (denoted E {displaystyle {mathcal
{E}}} and measured in volts),[1] is the electrical intensity or "pressure"
developed by a source of electrical energy such as a battery or generator.[2] A
device that converts other forms of energy into electrical energy (a
"transducer") provides an emf at its output
 In electromagnetic induction, emf can be defined around a closed loop of
conductor as the electromagnetic work that would be done on an electric
charge (an electron in this instance) if it travels once around the loop.

19.  Metals arranged according to standard
 potential values.
 More positive → noble metals
 More negative → active metals
 Only useful to predict which metal is anodic
 to other.
 Valid when activity of metal ions in
 equilibrium are unity i.e. 1.
 Alloys are not included (Only pure metals
 are considered).

21.  Arrangements of both metals and alloys.
 Well representative of particular
 environment.
 More appropriate for practical situation.
 It explain oxidation & reduction of alloys.

22.  In electrochemistry, a Pourbaix diagram, also known as a potential Vs pH
diagram, EH-pH diagram or a pE/pH diagram, maps out possible stable
(equilibrium) phases of an aqueous electrochemical system. Predominant ion
boundaries are represented by lines. As such a Pourbaix diagram can be read
much like a standard phase diagram with a different set of axes. Similarly to
phase diagrams, they do not allow for reaction rate or kinetic effects.
 The diagrams are named after Marcel Pourbaix (1904–1998), the Russian-born,
Belgian chemist who invented them

23.  Predict the spontaneous direction of the reaction .
 Estimate the composition of corrosion product .
 Predict the environmental condition that will prevent the corrosion.