Introduction to Electrochemistry
Electrochemistry is the branch of physical chemistry which deals with the relationship between electrical energy and chemical changes taking place in redox reactions.
Electrolytes are substances whose aqueous solution undergo decomposition into ions when electric current is passed through them. The whole process is known as electrolysis or electrolytic decomposition.
Electrolytic Cell (Voltameter)
The device in which the process of electrolysis or electrolytic decomposition is carried out is known as an electrolytic cell or voltameter.
Electrode Reactions
Primary products are formed directly from electrode reactions. These may undergo further changes to form secondary products.
Preferential Discharge Theory
If more than one type of ion is attracted towards a particular electrode, then the ion is discharged which requires least energy or ions with lower discharge potential or which occur low in the electrochemical series.
The potential at which the ion is discharged or deposited on the appropriate electrode is termed the discharge or deposition potential (D.P.).
Discharge Potential Order
Products of Electrolysis
| Electrolyte | Electrode | Product at Cathode | Product at Anode |
|---|---|---|---|
| Aqueous NaOH | Pt or Graphite | H2 | O2 |
| Fused NaOH | Pt or Graphite | Na | O2 |
| Aqueous NaCl | Pt or Graphite | H2 | Cl2 |
| Fused NaCl | Pt or Graphite | Na | Cl2 |
| Aqueous CuSO4 | Pt or Graphite | Cu | O2 |
| Aqueous CuSO4 | Cu electrode | Cu | Cu oxidized to Cu2+ |
| Dilute H2SO4 | Pt electrode | H2 | O2 |
| Aqueous AgNO3 | Pt electrode | Ag | O2 |
| Aqueous AgNO3 | Ag electrode | Ag | Ag oxidized to Ag+ |
Applications of Electrolysis
Thickness of Coated Layer
Mass of deposited substance = density × volume = d × a × b × c
Using Faraday's law: d × a × b × c = (I × t × E) / 96500
Faraday's Laws of Electrolysis
Given by Michael Faraday in 1833, these laws govern the deposition of substances on electrodes during electrolysis.
Electrochemical Equivalent (ECE): The mass of the ion deposited by passing a current of one Ampere for one second (by passing 1 Coulomb of electricity). Unit: gram per coulomb.
1 Coulomb = 6.28 × 1018 electrons
1 electronic charge = 1.6 × 10-19 Coulomb
For Gaseous Electrolyte Product
Where V = Volume of gas evolved at STP
Ve = Equivalent volume
Quantitative Aspects
1 Faraday will reduce:
• 1 mole of monovalent cation
• 1/2 mole of divalent cation
• 1/3 mole of trivalent cation
• 1/n mole of n-valent cations
Conductors and Conductivity
Conductivity Terms
Molar Conductivity (Λ)
Λ = (κ × 1000) / M (if M in mol/L)
Unit: S cm2 mol-1
Equivalent Conductivity (Λe)
Where C is concentration in gram equivalent per litre
Unit: Ohm-1 cm2 (gm equiv-1)
Experimental Measurement
Cell constant = l/a (in cm-1)
Factors Affecting Electrolytic Conductance
Degree of Dissociation (α)
Where Λc = molar conductance at concentration C
Λ0 = molar conductance at infinite dilution
Migration of Ions and Transport Number
Electricity is carried through electrolyte solution by migration of ions.
Ionic Mobility
Unit: Ohm-1 cm2 or V-1 S-1 cm2
λa or λc = ta or tc × λ0
Absolute ionic mobility = Ionic mobility / 96,500 (Unit: cm sec-1)
Kohlrausch's Law
At infinite dilution, the molar conductivity of an electrolyte can be expressed as the sum of the contributions from its individual ions.
Where ν+, ν- = number of cations and anions
λ+∞, λ-∞ = molar conductivities at infinite dilution
Applications
Galvanic Cell
A galvanic cell or voltaic cell is a device in which chemical energy is converted into electrical energy by a spontaneous redox reaction.
Electrode Reactions
Daniell Cell
Anode: Zn → Zn2+ + 2e-
Cathode: Cu2+ + 2e- → Cu
Overall: Zn + Cu2+ → Zn2+ + Cu
Standard Electrode Potential
The electrode potential is the potential difference set up between the electrode and its electrolyte. The standard electrode potential (E°) is measured at 25°C with 1M solution and 1 atm pressure.
Standard Hydrogen Electrode (SHE)
Electrochemical Series
| Electrode Reaction | E° (V) |
|---|---|
| Li+ + e- ⇌ Li | -3.05 |
| K+ + e- ⇌ K | -2.93 |
| Ca2+ + 2e- ⇌ Ca | -2.87 |
| Na+ + e- ⇌ Na | -2.71 |
| Mg2+ + 2e- ⇌ Mg | -2.37 |
| Al3+ + 3e- ⇌ Al | -1.66 |
| Zn2+ + 2e- ⇌ Zn | -0.76 |
| Fe2+ + 2e- ⇌ Fe | -0.44 |
| Ni2+ + 2e- ⇌ Ni | -0.25 |
| Sn2+ + 2e- ⇌ Sn | -0.14 |
| Pb2+ + 2e- ⇌ Pb | -0.13 |
| 2H+ + 2e- ⇌ H2 | 0.00 |
| Cu2+ + 2e- ⇌ Cu | +0.34 |
| Ag+ + e- ⇌ Ag | +0.80 |
| Hg2+ + 2e- ⇌ Hg | +0.85 |
| Br2 + 2e- ⇌ 2Br- | +1.08 |
| Cl2 + 2e- ⇌ 2Cl- | +1.36 |
| Au3+ + 3e- ⇌ Au | +1.50 |
Nernst Equation
The Nernst equation gives the relationship between electrode potential and the concentration of ions in solution.
E = E° - (0.0591/n) log Q (at 25°C)
Where Q = reaction quotient
For Electrode Reaction
E = E° - (0.0591/n) log (1/[Mn+])
E = E° + (0.0591/n) log [Mn+]
For Cell Reaction
Ecell = E°cell - (0.0591/n) log ([C]c[D]d / [A]a[B]b)
Applications
Important Points to Remember
Key Points for JEE Main
- Electrolytic cell: Electrical → Chemical energy, Ecell = -ve, ΔG = +ve
- Galvanic cell: Chemical → Electrical energy, Ecell = +ve, ΔG = -ve
- Faraday's first law: W = Z × I × t
- Faraday's second law: W1/W2 = E1/E2
- 1 Faraday = 96500 C = Charge on 1 mole of electrons
- Kohlrausch's law: Λm∞ = ν+λ+∞ + ν-λ-∞
- Nernst equation: E = E° - (0.0591/n) log Q
- Standard hydrogen electrode potential = 0.00 V
- Electrochemical series: Lower E° → stronger reducing agent
- Higher E° → stronger oxidizing agent
Do's
Don'ts
JEE Main Weightage
This chapter typically carries 2-3 questions in JEE Main, covering electrolysis, conductance, electrochemical cells, and Nernst equation applications.