1) Explain the Galvanic cell with suitable diagram, reaction and procedure.
Ans: A Galvanic cell, also known as a voltaic cell, is an electrochemical cell that produces electrical energy by means of a spontaneous redox reaction. The cell consists of two half-cells, each with an electrode and an electrolyte solution. When the two half-cells are connected by a salt bridge or porous barrier, electrons flow from one half-cell to the other, producing a potential difference and generating an electrical current.
Here is a diagram of a simple Galvanic cell:
The half-reactions that take place in the cell are:
At the cathode: Cu2+ + 2e- -> Cu
At the anode: Zn -> Zn2+ + 2e-
Overall reaction: Zn + Cu2+ -> Zn2+ + Cu
Here, the anode is made of zinc (Zn) metal, which oxidizes and loses electrons, while the cathode is made of copper (Cu) metal, which gains electrons and reduces. The electrolyte solutions contain the corresponding metal ions, Zn2+ and Cu2+.
The procedure for setting up a Galvanic cell is as follows:
Prepare the two half-cells by placing a strip of the anode metal (in this case, Zn) in a solution containing its corresponding metal ions (Zn2+), and a strip of the cathode metal (Cu) in a solution containing its corresponding metal ions (Cu2+).
Connect the two half-cells with a salt bridge, which can be made of a piece of filter paper soaked in a solution of a salt, such as KCl.
Connect the two electrodes with a wire, creating a complete circuit.
Measure the voltage difference between the two electrodes with a voltmeter, which gives an indication of the cell’s potential difference or electromotive force (EMF).
2) Derive the Nernst’s equation for reduction reaction.
Ans: The Nernst equation relates the electrode potential of an electrochemical cell to the concentrations of the species involved in the cell reaction. The Nernst equation for a reduction reaction can be written as:
E = E° – (RT/nF) ln([Ox]/[Red])
where:
E is the electrode potential of the reduction reaction
E° is the standard electrode potential of the reduction reaction
R is the gas constant (8.314 J/mol-K)
T is the temperature in Kelvin
n is the number of electrons involved in the reduction reaction
F is the Faraday constant (96,485 C/mol)
[Ox] is the concentration of the oxidized species (reactant)
[Red] is the concentration of the reduced species (product)
To derive the Nernst equation for a reduction reaction, we start with the general expression for the Gibbs free energy change of the reaction:
ΔG = -nFE
where:
ΔG is the Gibbs free energy change of the reaction
n is the number of electrons involved in the reduction reaction
F is the Faraday constant
E is the electrode potential of the reduction reaction
ΔG° = -nFE°
Rearranging this equation, we get:
E° = ΔG°/(-nF)
At non-standard conditions, we can use the Gibbs-Helmholtz equation to relate the change in Gibbs free energy to the change in enthalpy and entropy of the system:
ΔG = ΔH – TΔS
At constant temperature, we can write:
ΔG = ΔH – TS
Dividing both sides by -nF, we get:
ΔG/(-nF) = -ΔH/(-nF) + (T/nF)ΔS
The left-hand side of this equation is equal to the electrode potential (E), and the right-hand side can be written in terms of the standard enthalpy change (ΔH°) and entropy change (ΔS°) of the reaction:
E = E° – (RT/nF) ln([Ox]/[Red])
where:
E° is the standard electrode potential
R is the gas constant
T is the temperature in Kelvin
n is the number of electrons involved in the reduction reaction
F is the Faraday constant
[Ox] is the concentration of the oxidized species (reactant)
[Red] is the concentration of the reduced species (product)
3) Discuss the various types electrodes used in electrochemistry.
Ans:In an electrochemical cell, there are two electrodes, positive and negative.
• Each electrode constitutes a half cell or a single electrode.
• Although a number of electrodes are possible but the more important of these electrodes are grouped into the following types:
1. Metal-metal ion electrodes
2. Non-metal ion electrode
Ni(s)/Ni2+ (aq)
Oxidation reaction: Zn(s) – 2e- → Zn2+(aq)
Reduction reaction: Zn2+(aq)+2e-→Zn(s)
Pt/H2(g)/H+(aq)
These electrodes are reversible with respect to anions.
Oxidation reaction: 2Br- (aq)-2e-→Br2(g)
Reduction reaction: Br2(g)+2e-→2Br- (aq)
3. Metal -metal insoluble salt anion electrode:
Eg: When a metal electrode is coated with a thin layer of its sparingly soluble salt and keeping this electrode in the solution containing common anion this electrode is formed.
Ag/AgCl/Cl- (aq)
Pb/PbSO4/SO4 2-aq)
4. Metal – amalgum electrode:
• When a metal is dissolved in mercury, the solution is known as amalgum.
• Keeping a metal rod in amalgum, metal amalgam electrode is formed.
Eg:Cd(s)/(CdinHg) Pb(s)/(PbinHg)
• The advantage of this type of electrode is highly reactive metal are not used directly as an electrode but they are used in a form of amalgum.
5. Oxidation-reduction(redox)electrodes:
• These are electrodes in which the emf arises from the presence of ions of a substance in two different oxidation states.
• These electrodes are setup by dipping an inert metal like gold or platinum in to a solution containing ions in two different oxidation states of the substance.
• Eg. A platinum wire immersed in a solution of ferrous and ferric ions or stannous and stannic ions constitutes a redox electrode.
• These electrodes are represented as
Pt/(Fe2+(aq),Fe3+(aq))
Oxidation reaction: Fe2+-e→Fe3(aq)
Reduction reaction: Fe3++e-→Fe2+(aq)
4) Discuss the various types of Dry corrosion in detail.
Ans: The corrosion which takes place in the absence of moisture or water is called Dry corrosion. It is uniform corrosion. This happens due to the direct interaction of atmospheric components such as O2, X2, SO2, NO2, H2S, etc with the metallic materials in the absence of moisture. This can be discussed under basic three categories.
1) Corrosion due to oxygen
Alkali and alkaline earth metals react with oxygen at room temperature. Other metals except Ag, Au, Pt react with oxygen at elevated temperature.
Further interaction depends on the nature of the oxide film formed.
Stable film: If the oxide film is stable , it remains adhered to the surface strongly. This prevents further penetration of oxygen to the base metal. Egg. In case of oxides of Al, Pb, Cu etc.
Unstable film: If the oxide formed is unstable, it dissociates back to metal and oxygen. Therefore corrosion does not take place. E.g., Oxides of Ag, Au, Pt.
Volatile: If the oxide film is volatile, it gets volatized to expose new surface to interact with oxygen. Here corrosion continues to eat the whole metallic structure. Therefore corrosion is rapid and continuous, e.g., oxide of molybdenum.
Porous: Here the oxide film has pores or channels through which oxygen slowly diffuses to interact with the base metal. Therefore in this case corrosion is slow but continuous. E.g., Oxide of iron
2) Corrosion due to other corrosive gases:
The extent of corrosion depends on the chemical affinity between the base metal and the atmospheric components.
3) Liquid metal corrosion
This type of corrosion occurs when a molten liquid is passed continuously on a solid Metal/alloy surface. The corrosion is attributed to either dissolution of the molten liquid Into the solid metal phase.
5) Explain the Waterline Corrosion and Pitting Corrosion with suitable Diagram and reactions.
Ans: Waterline Corrosion : When water is stored in a steel tank, corrosion take place along a line just below the level of water meniscus. In this case, the area above the water line having more access to Oxygen acts as cathode, whereas the area below the water line having relatively less Access to oxygen acts as anode.
Pitting Corrosion: he metal surface having pits or cavities undergoes corrosion due to development of separate anodic and cathodic areas. Here pits act as anode with respect to the normal portion. A typical pitting corrosion cell in iron surface is given below. Here corrosion is due to oxygen concentration cell.
