CHEM1010 W12.2

Spontaneity of Reactions

Spontaneous Processes

  • Definition: A spontaneous reaction occurs on its own, without external assistance.
  • Non-spontaneous Process: The reverse of a spontaneous process.
  • Important Considerations:
    • Spontaneity does not equate to speed; it concerns direction and extent.
    • Non-spontaneous reactions are possible with external assistance (force, work, power).
    • Experimental conditions (temperature, pressure) can influence spontaneity.
  • Examples:
    • Water freezing in a freezer (spontaneous).
    • Ice melting on a countertop (spontaneous).
    • Electrolysis of water to produce hydrogen and oxygen (non-spontaneous, requires electricity).

Enthalpy, Exothermic, and Endothermic Reactions

  • Exothermic reactions release heat from the system to the surroundings.
  • Endothermic reactions absorb heat from the surroundings to the system.
  • Not all exothermic reactions are spontaneous, and not all endothermic reactions are non-spontaneous.
  • Example: Dissolution of ammonium nitrate in water is a spontaneous, endothermic reaction.

Chance, Probability, and Microstates

  • Scenario: Gas distribution between two flasks connected by a tap.
    • Gas spontaneously flows from a full flask to an empty one until pressure is equalized.
    • Reverse process (gas returning to one flask) is non-spontaneous.
  • Macroscopic vs. Microscopic View:
    • Macroscopic: Overall behavior of the gas.
    • Microscopic: Behavior of individual particles.
  • Microstates: Possible arrangements of particles in a given macroscopic state.
    • Example: Two particles in two flasks have four possible microstates.
  • Probability: Spontaneous processes increase the number of possible microstates, making them statistically more likely.

Entropy ((S))

  • Definition: A measure of the number of microstates associated with a macroscopic state.
  • Relationship to Microstates: Quantified using the Boltzmann constant ((k))
    • S=k×number of microstatesS = k \times \text{number of microstates}
  • Types of Microstates:
    • Position of particles.
    • Motion of particles.
    • Vibration of bonds.
    • Rotation of bonds.
  • Entropy and States of Matter:
    • Solids (e.g., ice): Fewer microstates (ordered, crystalline lattice).
    • Liquids: More microstates (molecules can slide and rotate).
    • Gases: Even more microstates (greater movement, more positions).
  • Interpretation: Entropy is related to the spreading of energy or matter, increasing randomness or disorder.

Laws of Thermodynamics

  • First Law: Energy cannot be created or destroyed; it converts from one form to another.
  • Second Law: For any spontaneous process, the entropy of the universe (system + surroundings) increases.
    • \Delta S{\text{universe}} = \Delta S{\text{system}} + \Delta S_{\text{surroundings}} > 0
  • Third Law: Entropy is an absolute quantity with a zero point.
    • Zero entropy is a perfectly pure, crystalline substance at 0 Kelvin.

Gibbs Free Energy ((G))

  • Definition: A quantity derived from enthalpy ((H)) and entropy ((S)) to determine spontaneity at constant pressure and temperature.
  • Formula:ΔG=ΔH−TΔS\Delta G = \Delta H - T\Delta S
  • Spontaneity Criterion:
    • \Delta G < 0: Spontaneous reaction.
    • \Delta G > 0: Non-spontaneous reaction.
    • ΔG=0\Delta G = 0: Reaction at equilibrium.
  • Endothermic reactions can be spontaneous if the entropy increase is large enough that (\Delta G) is negative.