Overview of Electrochemistry

The branch of chemistry, which deals with the study of the conversion of chemical energy into electrical energy or viceversa and their mutual relationship, is called Electrochemistry.

Electrochemical cell:
An electrochemical cell is a device that converts chemical energy into electrical energy or vice versa through redox reactions.

Galvanic (Voltaic) cell:
It converts chemical energy into electrical energy. Electric current is produced by a spontaneous redox reaction (e.g., Daniell cell).

Working of a Daniell cell:

Zinc acts as the anode and undergoes oxidation:
Zn → Zn²⁺+ 2e⁻

Copper acts as the cathode and undergoes reduction:
Cu²⁺ + 2e⁻→ Cu

Salt bridge role:
The salt bridge allows the movement of ions to maintain electrical neutrality and completes the circuit without mixing the solutions.

Electrolytic cell:
It converts electrical energy into chemical energy by forcing a non-spontaneous redox reaction, commonly used in electrolysis and extraction of metals.

The tendency of an electrode to lose or gain electrons when it is in contact with its own ions in solution is called electrode potential.

1. Electrode potential develops due to the separation of charges at the interface between the metal and the electrolyte solution.
2. When neither the metal atoms nor the metal ions undergo any change, no electrode potential is developed.
3. If metal atoms lose electrons and pass into solution as positive ions, the process is called oxidation.
4. During oxidation, electrons accumulate on the metal surface, causing the electrode to acquire a negative potential.
5. If metal ions from the solution gain electrons and get deposited as metal atoms, the process is called reduction.
6. During reduction, the electrode develops a positive potential due to the loss of electrons.
7. The magnitude of electrode potential depends on the nature of the metal and the concentration of its ions in solution.
8. Standard electrode potential is the electrode potential measured under standard conditions of temperature, pressure, and concentration.
9. Standard conditions include a temperature of 298 K, ion concentration of 1 mol L⁻¹, and gas pressure of 1 atm.
10. The standard hydrogen electrode is used as the reference electrode, and its standard electrode potential is assigned a value of zero.

The electromotive force (EMF) of a galvanic cell is defined as the difference of electrical potential which causes the flow of current from one electrode to another when virtually no current is drawn from the cell.

• Electrical conduction occurs due to the movement of charge carriers, which may be electrons or ions.
• In metallic conduction, electric current flows due to the movement of free electrons present in metals and alloys.
• Metallic conduction does not involve any chemical change and no transfer of matter takes place.
• The electrical conductivity of metals decreases with increase in temperature due to increased resistance.
• Electrolytic conduction occurs due to the movement of ions in molten electrolytes or their aqueous solutions.
• Electrolytes do not conduct electricity in the solid state because ions are not free to move.
• During electrolytic conduction, chemical decomposition of the electrolyte occurs (electrolysis).
• The electrical conductivity of electrolytes increases with increase in temperature due to increased ionisation and ionic mobility.

Substances which allow the passage of electric current through them are called electrical conductors.

Substances which do not allow the passage of electric current through them are called insulators.

The flow of electric current through a substance due to the movement of charge carriers is called electrical conduction.

Substances in which electric current is conducted by the movement of free electrons are called metallic conductors.

The conduction of electric current in metals due to the flow of free electrons without any chemical change is called metallic conduction.

Substances which conduct electricity in molten state or aqueous solution due to the movement of ions and undergo chemical change are called electrolytic conductors.

The conduction of electric current through an electrolyte due to the movement of ions accompanied by chemical change is called electrolytic conduction.

Substances which dissociate into ions in molten state or aqueous solution and conduct electricity are called electrolytes.

Substances which do not dissociate into ions and do not conduct electricity in aqueous solution are called non-electrolytes.

The fraction of total molecules of an electrolyte that dissociate into ions in solution is called the degree of dissociation.

Electrolytes which dissociate almost completely into ions in aqueous solution are called strong electrolytes.

Electrolytes which dissociate only partially into ions in aqueous solution are called weak electrolytes.

Specific resistance or resistivity is the resistance of a conductor having unit length and unit cross-sectional area.

Conductance is the measure of the ease with which electric current flows through a conductor and is defined as the reciprocal of resistance.

Specific conductivity is the reciprocal of resistivity and represents the ability of a material to conduct electric current.

Equivalent conductivity is the conducting power of all the ions furnished by one gram equivalent of an electrolyte present in a definite volume of the solution.

Molar conductivity is the conducting power of all the ions furnished by one mole of an electrolyte present in a definite volume of the solution.

• Ohm’s law states that the potential difference across a conductor is directly proportional to the current flowing through it, provided temperature remains constant.
• Mathematically, Ohm’s law is expressed as V ∝ I or V = I R
• The constant of proportionality \[R=\frac{V}{I}\] is called the resistance of the conductor, and its SI unit is ohm (Ω).
• Ohm’s law can also be stated as current is directly proportional to potential difference and inversely proportional to resistance of the conductor.

\[R=\rho\frac{l}{A}\]

or

\[\rho=\frac{RA}{l}\]

Unit: ohm cm (Ω cm) or ohm m (Ω m)

\[C=\frac{1}{R}\]

Unit: ohm⁻¹ or siemens (S)

\[\kappa=\frac{1}{\rho}\]

Unit: S cm⁻¹ or S m⁻¹

\[\Lambda_{eq}=\frac{\kappa}{c_{eq}}\]

\[\Lambda_m=\frac{\kappa}{c_m}\]

• Statement of the law:
  Kohlrausch’s law states that the molar conductivity of an electrolyte at infinite dilution is equal to the sum of the molar conductivities of its constituent ions, each multiplied by the number of ions present.

• Independence of ionic contribution:
  At infinite dilution, each ion contributes independently to the molar conductivity, irrespective of the nature of the other ion present.

• Mathematical expression:

  \[\Lambda_m^\circ=v_+\lambda_+^\circ+v_-\lambda_-^\circ\]​

  where λ^∘₊ and λ^∘_- are the ionic conductivities of cations and anions, and ν₊,ν₋ are their numbers.

• Importance/Application:
  Kohlrausch’s law is used to calculate the molar conductivity at infinite dilution of weak electrolytes using data of strong electrolytes.

The cell which converts electrical energy into chemical energy is called an Electrolytic cell and the process which converts electrical energy into chemical energy is called Electrolysis.

• The amount of substance liberated or deposited at an electrode during electrolysis is directly proportional to the quantity of electricity (charge) passed through the electrolyte.
• Since charge $Q=I\times t$, the mass deposited WWW is proportional to current and time.
• Mathematically, the law is expressed as W ∝ Q or W = Z I t
• The constant Z is called the electrochemical equivalent, defined as the mass of substance liberated by a current of 1 ampere flowing for 1 second.

• When the same quantity of electricity is passed through different electrolytes connected in series, the masses of substances liberated are directly proportional to their equivalent masses.
• Mathematically, \[\frac{W_1}{W_2}=\frac{E_1}{E_2}\] where W is mass deposited and E is equivalent mass.
• Electrochemical equivalent Z is related to equivalent mass by \[Z=\frac{E}{F}\]
• Faraday’s constant (F) is the charge carried by one mole of electrons and is equal to 96500 coulombs.

• Primary cells are electrochemical cells in which the cell reaction occurs only once and cannot be reversed by external electrical energy; hence, they cannot be recharged.
• Dry cell (Leclanché cell) uses a zinc container as anode and a graphite rod surrounded by MnO₂ and carbon as cathode, with a paste of NH₄Cl and ZnCl₂ as electrolyte; its cell potential is about 1.25–1.5 V.
• In a dry cell, zinc is oxidised at the anode and MnO₂ is reduced at the cathode; the cell has a short life due to the acidic electrolyte corroding the zinc container.
• Mercury cell uses zinc–mercury amalgam as anode and HgO as cathode with KOH–ZnO paste as electrolyte; it provides a constant voltage (~1.35 V) and is widely used in small devices like watches and hearing aids.

• Secondary cells are electrochemical cells that can be recharged by passing direct current in the opposite direction and can be used repeatedly; they are also called storage cells or accumulators.
• A lead storage cell consists of lead (Pb) as anode, lead dioxide (PbO₂) as cathode, and 38% sulphuric acid (H₂SO₄) as the electrolyte, with a cell potential of about 2.0 V.
• During discharge of a lead storage cell, both electrodes are converted into lead sulphate (PbSO₄) and sulphuric acid gets diluted, indicating discharge of the battery.
• Recharging is done by passing direct current in the reverse direction, which converts PbSO₄ back to Pb at the anode and PbO₂ at the cathode, restoring the original concentration of sulphuric acid.
• A nickel–cadmium (Ni–Cd) cell uses cadmium as anode, nickel oxide as cathode, and KOH solution as electrolyte, provides an EMF of about 1.4 V, has a long life, and can be recharged many times.

The cells which convert the chemical energy of a fuel directly into electrical energy are called fuel cells.

When a metal is exposed to certain environment, it may come in combined state and get destructed due to its reaction with the environment and its surface may become rough. This phenomenon is called corrosion.

or

The slow but spontaneous destruction of metals due to their interaction with certain environment is called corrosion.

• Electrochemical nature:
  Rusting of iron occurs by an electrochemical process in the presence of moisture, oxygen, and dissolved carbon dioxide, where different regions of the iron surface act as miniature galvanic cells.

• Anodic reaction (oxidation):
  At the anodic regions, iron gets oxidised to ferrous ions:

  Fe(s) → Fe²⁺(aq) + 2e⁻

• Cathodic reaction (reduction):
  At the cathodic regions, dissolved oxygen is reduced in the presence of hydrogen ions to form water:

  O₂(g) + 4H⁺(aq) + 4e⁻ → 2H₂O(l)

• Formation of rust:
  The ferrous ions are further oxidised to ferric ions, which combine with oxygen and water to form hydrated ferric oxide (Fe₂O₃ ⋅ xH₂O), known as rust; the process is accelerated by the presence of electrolytes and impurities.

1. Barrier protection:
   Corrosion can be prevented by placing a protective barrier between the metal and the atmosphere using paint, oil, grease, or chemical coatings, which stop contact with air and moisture.
2. Coating with non-corroding metals:
   Rusting can be prevented by coating iron with non-corroding metals like nickel or chromium (electroplating), which protect the surface from oxidation.
3. Galvanisation (sacrificial protection):
   Iron is protected by coating it with zinc, which being more electropositive than iron, gets oxidised first and thus saves iron even if the coating is scratched.
4. Cathodic protection:
   In this method, iron is connected to a more reactive metal like zinc or magnesium, which acts as a sacrificial anode and corrodes instead of iron.
5. Anti-rust solutions:
   Corrosion can be prevented by using anti-rust solutions such as phosphate or chromate solutions, which form an insoluble protective film on the metal surface.