Electrochemistry Lecture Notes

These vocabulary flashcards cover the fundamental principles of electrochemistry, including redox reactions, cell components, standard potentials, and electrochemical applications like batteries and electrolysis.

Electrochemistry

The study of oxidation-reduction reactions where the energy released by a spontaneous reaction is converted to electricity or electrical energy is used to cause a nonspontaneous reaction.

Oxidation half-reaction

The component of a redox reaction that involves the loss of electrons.

Reduction half-reaction

The component of a redox reaction that involves the gain of electrons.

Oxidation number

The charge an atom would have in a molecule (or an ionic compound) if electrons were completely transferred.

Free elements

Elements in an uncombined state, such as $\text{Na}$, $\text{Be}$, $\text{K}$, $\text{Pb}$, $\text{H}_2$, $\text{O}_2$, and $\text{P}_4$, which always have an oxidation number of zero.

Galvanic Cells

Electrochemical cells that use a spontaneous redox reaction to generate electricity.

Anode

The electrode in an electrochemical cell where oxidation occurs.

Cathode

The electrode in an electrochemical cell where reduction occurs.

Salt bridge

A device used in a galvanic cell to maintain electrical neutrality by allowing the flow of ions between the oxidation and reduction half-cells.

Electromotive force (emf)

The difference in electrical potential between the anode and cathode, also referred to as cell voltage or cell potential ($E^0_{\text{cell}}$).

Standard reduction potential ($E^0$)

The voltage associated with a reduction reaction at an electrode when all solutes are $\text{1 M}$ and all gases are at $\text{1 atm}$.

Standard hydrogen electrode (SHE)

A reference electrode with a reduction potential ($E^0$) defined as exactly $0\,V$ for the reaction $2e^- + 2H^+ (1\,M) \rightarrow H_2 (1\,atm)$.

Diagonal rule

A method used to compare standard reduction potentials to predict which species will oxidize others under standard-state conditions.

Faraday constant ($F$)

The charge of one mole of electrons, valued at approximately $96,500\,J\,V^{-1} \cdot mol^{-1}$ or $96,500\,C/mol$.

Nernst equation

An equation relating the cell potential ($E$) to the standard cell potential ($E^0$) and the reaction quotient ($Q$): $E = E^0 - \frac{RT}{nF} \ln(Q)$.

Mercury Battery

A type of battery using a zinc-mercury amalgam anode and a mercury(II) oxide cathode in a basic environment.

Lead storage battery

A battery commonly used in vehicles where the anode is lead ($\text{Pb}$) and the cathode is lead dioxide ($\text{PbO}_2$) in a sulfuric acid solution.

Fuel cell

An electrochemical cell that requires a continuous supply of reactants, such as $\text{H}_2$ and $\text{O}_2$, to maintain its function.

Corrosion

An electrochemical process that results in the deterioration of metals, such as the formation of rust on iron.

Cathodic Protection

A method of protecting a metal structure, such as an iron storage tank, from corrosion by making it the cathode of an electrochemical cell through the use of a sacrificial anode like magnesium ($\text{Mg}$).

Electrolysis

The process in which electrical energy is used to drive a nonspontaneous chemical reaction.