Chemistry: Buffer Solutions and Entropy

Chemical Energetics

• Atomisation enthalpy (ΔH_at): The enthalpy change when one mole of gaseous atoms is formed from the element in its standard state. ΔH_at is always positive (endothermic).
• Lattice dissociation enthalpy: The enthalpy change needed to convert one mole of solid crystal into its scattered gaseous ions. Lattice dissociation enthalpies are always positive (endothermic).
• Lattice formation enthalpy: The enthalpy change when one mole of solid crystal is formed from its scattered gaseous ions. Lattice formation enthalpies are always negative (exothermic).
• Ionic bonds are strong: When ions combine to form an ionic solid lattice, a large amount of energy is released.
• Ionic compounds are stable. The strong electrostatic forces of attraction between oppositely charged ions make ionic compounds more stable than their gaseous ions.

Factors Affecting Lattice Enthalpy

• Ionic charge: Increasing the ionic charge increases the attraction between the ions, leading to a larger, more negative lattice enthalpy.
• Ionic radius: Decreasing the ionic radius brings the ions closer together, increasing the attraction and leading to a larger, more negative lattice enthalpy.
• Charge density: Both ionic charge and ionic radius affect charge density. A higher charge density results in a greater lattice enthalpy.

Electron Affinity

• First electron affinity (EA₁): The enthalpy change when one mole of electrons is added to one mole of gaseous atoms to form one mole of gaseous ions each with a single negative charge. The first electron affinity is typically exothermic (negative) for most elements.
• Factors affecting EA:
  • Nuclear charge: An increase in nuclear charge increases the attraction between the nucleus and the incoming electron, resulting in larger, more negative affinities.
  • Distance: A greater distance between the nucleus and the outermost electron shell/orbital weakens the attraction, resulting in lower, less negative affinities.
  • Shielding: More electron shells increase the shielding effect, reducing the attraction and subsequently decreasing the electron affinity.

Trends in Electron Affinity of Group 16 & 17 Elements

• General trend: Moving down a group, EA₁ becomes less exothermic due to increased atomic size, greater distance between the nucleus and the outermost electron, and greater shielding.
• Fluorine exception: Fluorine has a lower EA₁ than chlorine due to its very small atomic radius resulting in high electron density, and therefore greater repulsion between the incoming electron and existing electrons.

Enthalpy Changes (Table of Definitions)

• A table including various enthalpy changes, such as formation, neutralisation, combustion, enthalpy of atomisation, ionisation enthalpy, hydration enthalpy, etc. along with their respective thermodynamic signs (positive or negative).

Enthalpy of Solution and Hydration

• Enthalpy of solution: The enthalpy change when one mole of an ionic substance dissolves in water to give a solution of infinite dilution.
• Hydration enthalpy: The enthalpy change when one mole of gaseous ions dissolves in sufficient water to give an infinitely dilute solution. Hydration enthalpy is always negative.
• Attraction between ions and water: When an ionic solid dissolves, positive and negative ions are formed in water. Water molecules form ion-dipole attractions with these ions.

Entropy

• Entropy (S): A measure of disorder in a system. Increasing disorder results in increased entropy, making the system energetically more stable.
• Gas > Liquid > Solid: Gases have the highest entropy, followed by liquids and then solids.
• Temperature change: Entropy generally increases with increasing temperature as the particles gain more energy and disorder increases.
• Change in state: Phase transitions (melting, vaporization, etc.) generally increase entropy. Dissolving a solid into a liquid, increases the entropy.

Gibbs Free Energy

• Gibbs free energy (ΔG): A measure of spontaneity. ΔG = ΔH - TΔS.
  • If ΔG is negative, the reaction is spontaneous at that particular temperature.
  • If ΔG is positive, the reaction is nonspontaneous at that particular temperature.
• Temperature change and feasibility: A reaction that is not feasible at low temperatures may become feasible at higher temperatures, or vice versa, depending on the relative enthalpy and entropy changes.
• Units of ΔG, ΔH, and ΔS are crucial.

Equilibria

• Conjugate acid-base pair: Compounds that differ by one proton.
• pH: The negative logarithm of the hydrogen ion concentration.
• pOH: The negative logarithm of the hydroxide ion concentration.

Acid Dissociation Constant (K_a)

• Ka: measure of the extent of dissociation of a weak acid; the smaller the Ka the weaker the acid.

Acidic Buffers

• Buffer solutions: Resist changes in pH when small quantities of acid or alkali are added.
• How Ka relates to buffer action: A buffer with a higher K_a value has greater resistance to pH change than a buffer with a lower K_a value.

Basic Buffers

• Buffer solutions: Resist changes in pH when small quantities of acid or alkali are added.
• How Ka relates to buffer action: A buffer with a higher K_a value has greater resistance to pH change than a buffer with a lower K_a value.

Amphoteric Buffers

• Amphoteric substances: Substancdes that can act as both acids and bases.
• Amino acids in blood: The amino and carboxyl groups in amino acids often act as acids and bases to maintain the blood pH.
• Carbonic acid/ bicarbonate buffer in blood: These two components buffer changes in blood pH.

Electrochemistry

• Methods for predicting substances liberated during electrolysis from various factors such as electrolyte state (molten or aqueous), redox series (electrode potential), and concentration.
• Balancing Redox half equations Details of steps in balancing atoms, water molecules, protons and electrons when balancing redox reactions in different scenarios.

Solubility Product (K_sp)

• Solubility Product (Ksp): A measure of the solubility of a sparingly soluble salt in water. A large Ksp means that the salt is more soluble.
• Sparingly Soluble Salts: Salts that dissolve only slightly in water.
• Common Ion Effect: Addition of a common ion to a saturated solution decreases the solubility of the salt.

Partition Coefficients

• Partition coefficient (K_pc): The ratio of concentrations of a solute in two immiscible solvents at equilibrium. The higher the K_pc, the greater the solute's solubility in the organic solvent.
• Factors affecting solubility.

This quiz explores key concepts related to buffer solutions, dissociation constants, and entropy in chemical systems. Understand how added acids affect equilibrium in buffers and the significance of entropy in measuring stability. Dive into the interactions between weak acids and their conjugate bases.