Chemistry Module 1: Equilibria Study Notes

Equilibria is fundamental in chemistry, referring to states where reactants and products of a reaction are present in concentrations that do not change over time.

Definition: Energy cannot be created or destroyed, only converted from one form to another.
Implications in Chemistry:
- The total energy of the universe remains constant during a chemical reaction.
- Energy can convert forms, e.g., chemical potential energy in bonds can turn into heat, light, or work.
- Example: Bond dissociation in HCl is endothermic, requiring heat to break bonds.
- Example: Burning methane is exothermic, releasing heat as chemical bond energy is converted into thermal energy.
- Note: This law does not provide insight into the spontaneity of reactions.

Definition: The entropy of the universe tends to increase.
Key Concepts:
- In every real process, the total entropy of the universe must either increase or remain the same during reversible processes.
- This law helps determine the feasibility of reactions based on entropy changes in the system and surroundings.

The First Law indicates energy balance.
The Second Law indicates increasing disorder within that balance.
- Example: In gases, heat disperses, increasing randomness.

Josiah Willard Gibbs: Recognized the relationship between enthalpy, entropy, and free energy.
Formula: The Gibbs Free Energy (G) is defined as $G = H - TS$ where:
- $H$ = Enthalpy
- $T$ = Temperature
- $S$ = Entropy
Gibbs Free Energy Change: $riangle G = riangle H - T riangle S$
- Represents the maximum non-volume expansion work extracted from a closed system at constant temperature and pressure.

At equilibrium, Gibbs Energy decreases as entropy increases.
If riangle G < 0 , the reaction is spontaneous.
A reaction will proceed towards minimizing Gibbs energy under constant temperature and pressure.
Standard Enthalpy of Reaction: $riangle H_r^{ heta}$ is the difference in molar enthalpy between products and reactants.
Standard Entropy of Reaction: $riangle S_r^{ heta}$ also calculated similarly.

Definition: Concentration is crucial; it defines how much of a substance is present in a mixture or solution (e.g., sugar in tea).
Common units: mol/L (molarity), mole fraction (mole/mole).

To find partial pressure $pJ$ , identify the mole fraction $xJ$ of the gas.
Partial molar properties reflect individual contributions to mixture properties.

Importance: Indicates how each substance in a mixture contributes to overall thermodynamic properties.
Key Concepts:
- The partial molar volume $V_J$ varies based on the composition of the mixture.
- For example:
- 1 mole of pure substance has a certain volume, while in a mixture, interactions alter that volume.

The partial molar Gibbs energy, defined as the ability of a substance to influence a reaction's tendency.
Significance: Chemical potential ( $ u$ ) dictates reaction directions and equilibriums based on concentration fluctuations.
Temperature and Pressure: Changes in pressure and temperature lead to shifts in potential and, consequently, in equilibrium.

Application of Raoult's Law (for solvents) and Henry's Law (for solutes) relates concentration to chemical potential.
Solvents can maintain stability via activity coefficients in real solutions versus ideal conditions.

Criterion for Spontaneity: The reaction Gibbs energy ( $riangle G$ ).
Spontaneity increases when riangle G < 0 , and reactions adjust until Gibbs energy minimizes.
Equilibrium constant, $K$ , links the reaction rates with concentrations of reactants/products.

Principle: A system at equilibrium will counter disturbances to restore balance.
Effects of changing concentrations, temperature, or pressure on equilibrium can be predicted.

Brønsted-Lowry Theory:
- Acids are proton donors; Bas are proton acceptors.
- Water can self-ionize, leading to hydronium ions influencing pH and equilibria.
Equilibrium Constants:
- $Ka$ for acids and $Kb$ for bases help assess strength.

Polyprotic acids can release multiple protons, altering the equilibrium dynamics based on pH.
Speciation: How different ion forms and their concentrations relate to pH changes in solution.

Understanding the point of neutralization and buffer systems is vital in controlling pH in chemical reactions.
The Henderson-Hasselbalch equation links pH to concentrations of acid and base forms.

These principles pave the way to understanding chemical equilibria, the influence of changes in conditions, and how systems can shift towards equilibrium or spontaneity.