General Chemistry II - Concepts Review

Lecture 24: Spontaneity

• Spontaneous Processes:

  • Definition: A process that occurs without outside intervention.
  • Characteristics: Speed or rate does not define spontaneity; spontaneous processes may occur slowly.

• Factors Predicting Spontaneity:

  • Energy Change
  • Temperature (∆H)
  • Entropy Change (∆S)

• Importance of Gibbs Free Energy (∆G):

  • The relationship: \Delta G = \Delta H - T \Delta S
  • For spontaneous processes: \Delta G < 0
  • Entropy definition: Measure of disorder; higher disorder results in greater entropy.

• Predicting spontaneity based on ΔS and ΔH:

  ΔS	ΔH	Result
  +	-	Spontaneous at all T
  +	+	Spontaneous at high T
  -	-	Spontaneous at low T
  -	+	Non-spontaneous at any T

• Laws of Thermodynamics:

  1. First Law: Energy of the universe is constant.
  2. Second Law: Spontaneous processes increase entropy.
  3. Third Law: At absolute zero (0 K), a perfect crystal has an entropy of zero.

Lecture 25: Free Energy

• Gibbs Free Energy:

  • Definition: Maximum non-expansion work from a closed system.
  • Formula: \Delta G^{\circ}{rxn} = \Sigma n \Delta G^{\circ}f(products) - \Sigma m \Delta G^{\circ}_f(reactants)

• Significance of Gibbs Free Energy:

  • \Delta G < 0 indicates spontaneous forward reaction.
  • \Delta G > 0 indicates spontaneous backward reaction.
  • At equilibrium: \Delta G = 0

• Hess's Law for Gibbs Free Energy:

  • Reverse reactions change the sign of ∆G.
  • Multiplying reactants by a factor multiplies ∆G by the same factor.

• Equilibrium Relations:

  • K = e^{-\Delta G^{\circ}/RT}

Lecture 26: Electrochemistry Introduction

• Electrochemistry:

  • Study of chemistry changes and electrical work interchanges.

• Oxidation & Reduction:

  • Oxidation: Loss of electrons; species that loses electrons is the reducing agent.
  • Reduction: Gain of electrons; species that gains electrons is the oxidizing agent.

• Electrochemical Cells:

  • Cathode: Region of reduction (negative).
  • Anode: Region of oxidation (positive).

Lecture 28: Galvanic Cells

• Differences between Galvanic and Electrolytic Cells:

  • Galvanic: Spontaneous; generates current.
  • Electrolytic: Needs external current to drive a non-spontaneous reaction.

• Cell Notation:

  • Format: Anode || Cathode
  • Example: \text{M(s) | Mm^{+}(aq) || Nn^{+}(aq) | N(s)}

• Function of Salt Bridge: Allows unintended ion movement balancing charge without electron build-up.

Lecture 29: Standard Reduction Potentials

• Cell Potential: Difference in electric potential between electrodes;

  • Formula: E^{\circ}(cell) = E^{\circ}(cathode) - E^{\circ}(anode)

• Nernst Equation:

  • Equilibrium: E^{\circ}(cell) = \frac{2.303RT}{nF} \log(K)
  • Non-standard: E(cell) = E^{\circ}(cell) - \frac{RT}{nF} \ln(Q)

Lecture 31: Nuclear Chemistry

• Nuclear Chemistry: Study of nucleus structure changes and chemistry.
• Radioactive Nuclei: Nuclei that spontaneously change structure emitting radiation.
• Nuclear Reactions vs Chemical Reactions:
  • Nuclear reactions release more energy; not sensitive to chemical environment.

Lecture 33: Nuclear Decay

• Half-Life: t_{1/2} = \frac{0.693}{k}

• First Order Process: Linear decay law relation: ext{Ln}(N{t}/N{0}) = -kt

• Decay: Nuclides want to decay to become more stable, often due to an unstable ratio of neutrons/protons.

• Example Equation: (N{0})(\frac{1}{2})^{n} = N{t}
  where n = number of half-lives.