Electrochemistry

  • The study of chemical reactions that produce electricity or use electricity to drive a desired product (the interchange of chemical and electrical energy).

  • Electrochemical reactions oftren result in electron transfer, often called redox reactions:

    • Oxidation: electrons are loss from reactant (reducing agent)

    • Reduction: electrons are gained from reactant (Oxidizing agent)

  • Redox reactions only occur in reactants and it always happens at the same time

Galvanic Cell

  • controlled environment for spontaneous redox reactions to occur

  • produces electricity spontaneously

  • basically it harnesses the electron transfer to create actual electricity

    • examples: Battery chargers

  • Salt bridge acts as a way to balance the build up of charge between cells, thus it also completes the circuit.

  • There are 2 ways for charges to be conducted

    • Metallic conduction, the charge travels through metal such as the wires connected in a galvanic cell

    • Electrolytic conduction, the charge travels through the ions in liquid ie the salt bridge

Half Cell (we focus on the metal)

  • Reduction (Cathode): Occurs when a dissolved metal ion collides with the electrode, gains electrons, and deposits onto the surface. [RED CAT]

  • Oxidation (Anode): Occurs when a neutral metal atom on the electrode surface loses electrons and enters the solution as an ion. [AN OX]

  • Equilibrium: Isolated half-cells quickly establish a dynamic equilibrium between the solid metal and its dissolved ions:

Standard Cell Notation

Anode Electrode | Anode Solution || Cathode Solution | Cathode Electrode

Cells are represented using a shorthand format following specific rules:

  • A single vertical line (|) or slash (/) represents a phase boundary (e.g., when the reactant to product partnership has different states solid-liquid etc)

  • A double vertical line (||) represents the salt bridge.

  • A semi-colon (;) represents a boundary when the reactant to product partnership are same states of matter

Cell Potential and half cell potential

  • is the energy or force needed for the electron to move from one place to to another ie from anode to cathode within a cell

  • measured in volts as 1v = 1Jule/coulomb

  • Voltage is the driving force of the reactions

  • Cell potential disregards coefficient of equation since its intensive

  • Higher Charge= reduction

  • Lower Charge= oxidation

  • Cell potential is the difference in the potential for cathode and anode, half-cell potential are the individual cell potentials of both

    • To determine which one is the anode and cathode we can look at the half-cell potential as the lesser one will always be the anode (oxidized)

    • If half-cell potential is negative then it is easily oxidized, if half-cell potential is positive then it is easily reduced

Special cases

  • If both sides of your half reaction are lone Hydrogen then its half-cell potential is 0V

Spontaneousness

  • Predicting Spontaneity using E cell: Given any redox reaction, we can calculate spontaneity based on signs

  • Spontaneous Reaction: A reaction is spontaneous if E cell (positive value).

  • Non-spontaneous Reaction: A reaction is non-spontaneous if E cell (negative value). This means the reaction is spontaneous in the reverse direction.

  • Galvanic Cell Spontaneity: A galvanic cell is always spontaneous because it is inherently designed and set up so that its E cell is positive.