Exam III Review: Thermodynamics, Electrochemistry, Nuclear Chemistry, and Organic Chemistry

Chapter 7: Chemical Thermodynamics

Thermodynamic Quantities and Spontaneity

Thermodynamics studies the energy changes accompanying chemical and physical processes. Key thermodynamic quantities include enthalpy (ΔH), entropy (ΔS), and Gibbs free energy (ΔG), which help predict whether a process is spontaneous.

• ΔG (Gibbs Free Energy): Determines spontaneity. If ΔG < 0, the process is spontaneous; if ΔG > 0, nonspontaneous; if ΔG = 0, the system is at equilibrium.

• ΔH (Enthalpy): Heat absorbed or released at constant pressure. Exothermic reactions have ΔH < 0; endothermic reactions have ΔH > 0.

• ΔS (Entropy): Measure of disorder. Processes that increase disorder (ΔS > 0) are generally favored.

• Spontaneity: A process is spontaneous if it can occur without outside intervention. Spontaneity depends on both ΔH and ΔS, as described by ΔG.

Key Equation:

$\Delta G = \Delta H - T\Delta S$

• Standard State: Standard enthalpy (ΔH°), entropy (ΔS°), and free energy (ΔG°) values are measured under standard conditions (1 atm, 1 M, 25°C).

• Phase Changes: Entropy increases when a substance changes from solid to liquid to gas.

• Boiling Point: At the normal boiling point, ΔG = 0 for the phase change.

Example: For the melting of ice at 0°C, ΔH > 0 (endothermic), ΔS > 0 (increased disorder), and ΔG = 0 (equilibrium).

Calculating Thermodynamic Quantities

• ΔG from ΔH and ΔS: Use the equation above, ensuring units are consistent (ΔH in kJ, ΔS in J/K).

• ΔG° from Standard Free Energies of Formation:

  $\Delta G^\circ = \sum n\Delta G^\circ_{f,\text{products}} - \sum n\Delta G^\circ_{f,\text{reactants}}$

• ΔG from Equilibrium Constant (K):

  $\Delta G^\circ = -RT \ln K$

  • R = 8.314 J/(mol·K)

  • T = temperature in Kelvin

Example: If K > 1, ΔG° < 0 (spontaneous under standard conditions).

Entropy and the Second Law of Thermodynamics

• Second Law: The total entropy of the universe increases in a spontaneous process.

• ΔS_univ = ΔS_sys + ΔS_surr

Example: Dissolving NaCl in water increases the entropy of the system.

Chapter 20: Electrochemistry

Redox Reactions and Electrochemical Cells

Electrochemistry involves the study of redox (reduction-oxidation) reactions and their relationship to electricity. Redox reactions involve the transfer of electrons between species.

• Oxidation: Loss of electrons (increase in oxidation number).

• Reduction: Gain of electrons (decrease in oxidation number).

• Oxidizing Agent: Causes oxidation; is reduced.

• Reducing Agent: Causes reduction; is oxidized.

Balancing Redox Reactions:

• Assign oxidation numbers to all atoms.

• Identify species oxidized and reduced.

• Balance atoms and charges, often using the half-reaction method.

Electrochemical Cells

• Galvanic (Voltaic) Cell: Converts chemical energy to electrical energy; spontaneous redox reaction.

• Electrolytic Cell: Uses electrical energy to drive a nonspontaneous reaction.

• Cell Notation: Anode | Anode solution || Cathode solution | Cathode

Standard Electrode Potentials (E°): Measured relative to the standard hydrogen electrode (SHE, E° = 0 V).

Cell Potential:

$E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$

• Relationship to ΔG:

  $\Delta G^\circ = -nFE^\circ_{\text{cell}}$

  • n = number of moles of electrons transferred

  • F = Faraday's constant (96485 C/mol e⁻)

• Nernst Equation: Calculates cell potential under nonstandard conditions.

  $E = E^\circ - \frac{0.0592}{n} \log Q$

  • Q = reaction quotient

Example: In a Zn/Cu cell, electrons flow from Zn (anode) to Cu (cathode).

Electrolysis and Faraday's Laws

• Electrolysis: Nonspontaneous redox reactions driven by an external voltage.

• Faraday's Laws: The amount of substance produced at an electrode is proportional to the quantity of electricity passed.

• Calculating Mass:

  $\text{mass} = \frac{ItM}{nF}$

  • I = current (A), t = time (s), M = molar mass, n = electrons per ion, F = Faraday's constant

Example: Electrolysis of molten NaCl produces Na at the cathode and Cl₂ at the anode.

Chapter 21: Nuclear Chemistry

Types of Radiation and Nuclear Reactions

Nuclear chemistry studies changes in atomic nuclei, including radioactive decay and nuclear reactions.

• Alpha (α) Particle: Helium nucleus (⁴₂He), charge +2.

• Beta (β) Particle: Electron (⁰₋₁e), charge -1.

• Positron (β⁺): Positive electron (⁰₁e), charge +1.

• Gamma (γ) Ray: High-energy photon, no charge.

Balancing Nuclear Equations:

• Sum of atomic numbers and mass numbers must be equal on both sides.

• Identify missing particles or products by balancing these numbers.

Example: $^{14}_6\text{C} \rightarrow ^{14}_7\text{N} + ^{0}_{-1}\text{e}$ (beta decay)

Chapter 24: Organic Chemistry (Introduction)

Functional Groups and Nomenclature

Organic chemistry focuses on compounds containing carbon. Functional groups determine the chemical properties of organic molecules.

• Functional Group: Specific group of atoms responsible for characteristic reactions (e.g., alcohol, carboxylic acid, amine).

• Naming Organic Compounds: Use IUPAC rules to assign names based on the longest carbon chain and functional groups.

• Structural Formulas: Show the arrangement of atoms in a molecule.

Example: Ethanol (CH₃CH₂OH) contains a hydroxyl (–OH) functional group.

Summary Table: Key Thermodynamic and Electrochemical Equations

Equation	Description
$\Delta G = \Delta H - T\Delta S$	Gibbs free energy change
$\Delta G^\circ = -RT \ln K$	Standard free energy and equilibrium constant
$E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$	Standard cell potential
$\Delta G^\circ = -nFE^\circ_{\text{cell}}$	Relationship between free energy and cell potential
$E = E^\circ - \frac{0.0592}{n} \log Q$	Nernst equation (at 25°C)
$\text{mass} = \frac{ItM}{nF}$	Mass of substance produced in electrolysis