Electrochemistry Lecture Notes

Vocabulary and key concepts from Chapter 3: Electrochemistry, covering electrochemical cells, conductivity, electrolysis laws, and batteries.

Electrochemistry

The branch of chemistry which deals with the study of production of electricity from energy released during spontaneous chemical reactions and the use of electrical energy to bring about non-spontaneous chemical transformations.

Metallic Conductors (Electronic Conductors)

Substances which allow electric current to pass through them by the movement of electrons, such as metals.

Electrolytic Conductors (Electrolytes)

Substances which allow the passage of electricity through their fused state or aqueous solution and undergo chemical decomposition, such as aqueous solutions of acids, bases, and salts.

Strong electrolytes

Electrolytes that completely dissociate or ionise into ions, such as $HCl$, $NaOH$, and $K_2SO_4$.

Weak electrolytes

Electrolytes that dissociate partially (degree of dissociation $\alpha < 1$), such as $CH_3COOH$, $H_2CO_3$, and $NH_4OH$.

Standard cell

A cell of almost constant electromotive force (emf), with the Weston standard cell being the most common example.

Daniell Cell

An electrochemical cell of zinc and copper metals where the cathode is represented on the RHS and the anode on the LHS.

Salt-bridge

A device containing a solution of strong electrolyte (e.g., $KNO_3$, $KCl$) that completes the circuit and maintains electrical neutrality on both sides of an electrochemical cell.

Transport number (Transference number)

The fraction of the total current carried by an ion, where the sum of the cation transport number ($n_c$) and anion transport number ($n_a$) equals 1.

Electrode Potential

The tendency of an electrode to lose or gain electrons when in contact with the solution of its ions; it is an intensive property expressed in volts.

Oxidation potential

The tendency of an electrode to lose electrons, which is inversely proportional to the concentration of ions in the solution.

Reduction potential

The tendency of an electrode to gain electrons, which by IUPAC convention is alone called electrode potential unless otherwise mentioned, represented as $E^\text{o}_{red} = -E^\text{o}_{oxidation}$.

Standard electrode potential ($E^\text{o}$)

The potential difference developed between a metal electrode and a solution of ions of unit molarity ($1\text{ M}$) at $1\text{ atm}$ pressure and $25^{\text{o}}\text{C}$ ($298\text{ K}$).

Reference Electrode

An electrode of known potential used to determine the absolute value of electrode potential, categorized as primary (Hydrogen electrode) or secondary (Calomel electrode).

Standard hydrogen electrode (SHE)

A reference electrode also known as NHE consisting of a platinum wire and foil in $1\text{ M HCl}$ with hydrogen gas at $1\text{ atm}$ and $298\text{ K}$, with a fixed potential of zero at all temperatures.

Electromotive Force (emf)

The difference between the electrode potentials of two half-cells that causes the flow of current from an electrode at higher potential to one at lower potential.

Electrochemical Series

The arrangement of electrodes in increasing order of their standard reduction potentials.

Nernst Equation

The equation that provides the relationship between the concentration of ions and the electrode potential.

Standard Gibbs free energy change ($\Delta G^\text{o}$)

Related to the equilibrium constant ($K_c$) by the equation $\Delta G^\text{o} = -2.303RT \text{ log}(K_c)$.

Conductance ($G$)

The ease of flow of electric current through a conductor, defined as the reciprocal of resistance ($G = \frac{1}{R}$) with units in $\text{ohm}^{-1}$ or $\Omega^{-1}$.

Specific Conductivity ($K$)

The reciprocal of specific resistance; it decreases on dilution because the concentration of ions per cc decreases.

Molar Conductivity ($\Lambda_m$)

The conductivity of all ions produced when $1\text{ mole}$ of an electrolyte is dissolved in $V\text{ mL}$ of solution, calculated as $\Lambda_m = \frac{k \times 1000}{M}$.

Equivalent conductivity ($\Lambda_{eq}$)

The conducting power of all ions produced when $1\text{ g-equivalent}$ of an electrolyte is dissolved in $V\text{ mL}$ of solution, calculated as $\Lambda_{eq} = \frac{k \times 1000}{N}$.

Debye-Huckel Onsagar equation

An equation relating molar conductivity ($\Lambda_m$) at a particular concentration to molar conductivity at infinite dilution ($\Lambda^0_m$): $\Lambda_m = \frac{\Lambda^0_m - b\text{ }\times\text{ }\text{ }\text{ }\text{ }\text{ }\text{ }\text{ }\text{ }\text{ }\text{ }\text{ }\text{ }\text{ }\text{ }\text{ }\text{ }\text{ }\text{ }\text{ }\text{ } \sqrt{C}}{\text{ }}$.

Kohlrausch’s Law

States that at infinite dilution, the molar conductivity of an electrolyte is the sum of the ionic conductivities of its cations and anions.

Electrolysis

The process of decomposition of an electrolyte when electric current is passed through either its aqueous solution or molten state.

Over voltage (Bubble voltage)

The extra voltage required for the liberation of gases during electrolysis compared to the theoretical value of the standard electrode potential.

Discharge potential

The minimum potential that must be applied across the electrodes to bring about electrolysis and the subsequent discharge of the ion on the electrode.

Faraday’s First Law of Electrolysis

States that the amount of substance ($W$) deposited or liberated at the cathode is directly proportional to the quantity of electricity ($Q$) passed through the electrolyte ($W = Z \times I \times t$).

Electrochemical equivalent ($Z$)

The constant representing the amount of substance deposited or liberated by passing $1\text{ A}$ current for $1\text{ sec}$ ($1\text{ coulomb}$).

Faraday’s Second Law of Electrolysis

States that when the same quantity of electricity is passed through different electrolytes, the amounts of substances deposited are directly proportional to their equivalent weights.

Primary Batteries

Batteries in which the chemical reaction occurs only once and cannot be reused after becoming dead, such as the Leclanehe (dry) cell or Mercury cell.

Secondary Batteries

Cells that can be recharged and reused repeatedly by reversing the cell reactions, such as Lead storage batteries or Nickel-cadmium storage cells.

Fuel Cells

Galvanic cells that produce electrical energy directly from the combustion of fuels like $H_2$, $CH_4$, or $CH_3OH$; they are pollution-free and highly efficient.

Corrosion

The slow formation of undesirable compounds like oxides, sulphides, or carbonates on a metal surface through reactions with moisture and atmospheric gases.

Cathodic protection

A method of corrosion prevention where a metal is protected by connecting it to another metal that is more easily oxidised.