Electrochemistry

"Electrochemistry" 

Electrochemistry is branch of analytical chemistry which deals with the interconversion of electrical and chemical energy.

Types of Electrochemical Reaction:

Electrochemical reaction is also called redox reaction and is further divided into 2 types:

1. Spontaneous Redox Reaction:

Spontaneous Redox Reaction do not require external source of energy. It has sufficient chemical energy to operate the cell.

2. Non-spontaneuos Redox Reaction :

Non-spontaneuos redox reaction require external source of energy. It runs in the opposite direction of Spontaneous Redox Reaction. 

Redox Reaction:

Redox Reactions are pH dependent. 

There are 2 types of redox reactions, which are discussed below:

1. Reduction Reaction:

Gain of electron or decrease in oxidation number is called reduction reaction. 

2. Oxidation Reaction:

Loss of electron or increase in oxidation number is called oxidation reaction. 

Reducing Agents:

Agents which reduce others and oxidize itself are called reducing agents. 

Oxidizing Agents:

Agents which oxidize others and reduce itself are called oxidizing agents. 

Electrodes:

There are 2 types of electrodes compulsion used in electrochemistry. 

1. Inert Electrodes:

These electrodes don't take part in the reaction. They consume external source of energy. They are made of platinum(Pt), graphite, etc. 

2. Active Electrodes:

These electrodes take part in the reaction. They consume chemical energy of the Spontaneous Redox Reaction. They are made of copper, aluminum, zinc, etc. 

Electrolytic Cells:

Electrolytic cells used electrical energy to drive Non-spontaneuos reactions. 

For Non-spontaneuos redox reactions, we have to prepare:

• Suitable, insulating electrolytic tank made of Pyrex, wood, etc. 
• Electrode compulsion should be inert.
• Electrolyte in the form of molten, mixture of ionic compound and aqueous solution of ionic compound. 
• External source of energy.

Electrolysis of Molten NaCl:

For the electrolysis of molten NaCl took place in an insulated tank. First of we should keep in mind that the initiative reaction is always an oxidation reaction. 

When sodium ion collide with the negative electrode, the battery carries a large enough potential to force them to gain electrons to form sodium metal. The potential required to reduce sodium ion into metal is - 2.71V.

 

Now, the chloride ions move toward positive electrode where they oxidized to form chlorine gas. The petential required to oxidize chloride ion into gas is - 1.36V.

Now, the overall reaction is given by:

The potential required by battery to drive this reaction must be therefore equals to 4.07V.

Electrolysis of Mixture of Ionic compound:

For the electrolysis of mixture of ionic compounds, first we need to distinguish between the ions that which of them goes for oxidation and which of them goes for reduction. 

Now the ionic equations for the mixture of ionic compounds ( CaCl2 & KBr) are given below.

Ca+2 + 2e- → Ca

K+1 + e-  → K

2Cl- → Cl2 + 2e-

2Br- → Br2 + 2e-

For Oxidation:

We need to compare ion's electronegativities. Ion having less electronegativity goes for oxidation. Electronegativity of Cl>Br, therefore bromide ion should go for oxidation at anode.

2Br- → Br2 + 2e-

For Reduction:

We need to compare ion's ionization potential. Ion having greater I.P goes for reduction. Ionization potential of Ca>K, so calcium ion should go for reduction at cathode.

Ca+2 + 2e- → Ca

Overall Cell Reaction:

Ca+2 + 2Br- → Ca + Br2

Electrolysis of Aqueous Solution Of Ionic Compound:

For the electrolysis of aqueous solution of ionic compound, we need to compare their reduction potentials. Consider an electrolytic cell for the electrolysis of aqueous solution of Potassium Iodide.

Cathode Reactions:

 

The reduction potential of water is greater than potassium ion, so water is a strong oxidant. Therefore, water goes for reduction at cathode.

Anode Reactions:

The reduction potential of iodine ion is greater than water, so iodine is a strong reductant. Therefore, iodine goes for oxidation at anode.

Redox Reactions:

The potential required by the cell to operate the redox reaction is 1.37V.

Over-Voltage:

The increment of voltage for the production of gaseous product when aqueous solution is used, is called Over-voltage. This voltage is used for the dissociation of a molecule to its constituent (Water).

Faraday's Law of Electrolysis:

There are two laws given by Faraday on the Electrolysis.

Faraday's 1st Law of Electrolysis:

Faraday first law of electrolysis states that:

The mass of the substance liberated or deposited at the electrode is directly proportional to the quantity of charge that you are passing through the electrolytic solution.

m ∝ Q

m  ∝ It

m = Z'It

Where, Z' is the Electrochemical equivalent is that amount of substance liberated at the electrode when a current of 1 ampere is passed through the Electrolyte for 1 second.

Faraday's 2nd Law of Electrolysis:

Faraday 2nd law of electrolysis stated that:

When same quantity of electricity is passed through different electrolytes connected in series then the masses of the substances liberated at the electrode is equivalent to their electrochemical weight (equivalent).

 m = ZIt/F

Numerical of Faraday's Law of Electrolysis:

Electrochemical Cells:

And, the the purpose of salt bridge is explained in the video given below:

Electrochemical cells used chemical energy to operate spontaneous Redox reaction. In electrochemical cells, we use

• Two insulated tanks called half cells to separate oxidation and reduction. 
• We can use inert electrode but active electrode is preferable. 
• Maintain standard condition i.e 290K, 1atm, 1M.
• Reaction should be indirect redox i.e reduction and oxidation take place in different cells. 
• It should follow nomenclature LOAN ( Left Oxidation Anode Negative).
• If electrons flow from left to right then it is spontaneous direction and does not require external energy.
• Salt Bridge is used to maintain electrical neutrality and for the completion of circuit. 

Electrochemical Cell of Zinc and Copper:

For the electrochemical cell, we need to identify which metal goes for oxidation at anode and which metal goes for reduction at cathode. For this, we need to compare their potential.

The potential used by copper electrode is 0.34V and that of zinc electrode is - 0.76V. So, potential of Cu>Zn, therfore Zn metal goes for oxidation at anode and placed at left half cell. 

Left Cell:

In the left compartment, oxidation occurs by LOAN. So the ions present in the left compartment are:

Right Cell:

In the right compartment, it is obvious that reduction occurs. So the ions present in the right compartment are:

Overall Cell Reaction:

The potential produce in the oxidation of Zn electrode is - 0.76V. And the potential produce in the reduction of Cu electrode is 0.34V. So the overall reaction is:

Emf of the Cell:

Cell Notation:

Example:

 

Electrochemical Series:

Electrochemical series is a series in which reduction petential of electrodes are arranged in the increasing order (downward) or increasing order of oxidation potential (upward). 

Aplications of Electrochemical Series:

Some of the applications of electrochemical series are listed below:

• Determine the Emf of the cell. 
• Feasibility or spontaneity of the cell. 
• Reactivity of metals. 
• Identification of the metals having ability to produce hydrogen gas. 
• Relative strength of reducing and oxidizing agents. 

Corrosion:

The slow, harmful, undesirable and continuous reaction of a metal material in which it is eaten away by moisture or any chemical agents when exposed in open is called Corrosion. 

Theories Related to Corrosion:

1. Dry Theory of Corrosion:

• It does not require any conducting medium or occurs in the absence of moisture. 
• It involves direct attack on the metal surface and the process is slow. 
• The product is produced at the site of corrosion. 
• The process is uniform. 
• Due to the direct attack of air on the metal surface, it is also called oxide corrosion. 
• Types of metal oxide layers formed due to chemical corrosion are:

1. Stable oxide layer: protective layers on heavy metals like Pd, Hg, CR, Cd, etc. 
2. Unstable oxide layer: reaction is reversible and occurs in alkali metals, Au, etc. 
3. Volatile oxide layer: layer is volatile and evaporate and usually occurs in molybdenum. 
4. Porous oxide layer:layer formed has the ability to absorb moisture which form cracks and under metal layer is corroded and usually take place in Iron. 

Decarburization:

In steel industry, steel is made by combining iron and carbon in a specific ratio. When steel exposed to less acidic pH or HCl, H2S i.e compound containing hydrogen, then acid reacts with carbon to form methane gas which diffuse into metal matrix and exert pressure. Then remaining acid reacts with iron to form hydrogen gas. This hydrogen gas causes corrosion called Hydrogen embrittlment type corrosion. 

Piling Bedworth Rule:

Piling Bedworth rule says, we need to compare volume (thickness) of metal surface and metal oxide layer. 

• If MOL < ML, then it is unprotective or unstable layer of corrosion. 
• If MOL > ML, then it is protective or stable layer of corrosion. 

2. Wet (Electrochemical) Theory of Corrosion:

• It requires conducting medium i.e Electrolyte. 
• It is further divided into two types:

1. Hydrogen Evolution Type of Corrosion:

This type of corrosion causes "the displacement of hydrogen ions from the acidic solution by metal surface".

2. Oxygen Absorption Type of Corrosion:

The liberated electrons from anodic to cathodic area from metal surface gained by oxygen amd get reduced. 

The ferrous ion reacts with the resulting hydroxide ion. 

If enough oxygen is present then:

The product is called yellow rust. 

3. Acid Theory of Corrosion:

It says whenever there is an acidic medium present for corrosion, the rate become faster. 

Method to Save Metals from Corrosion:

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