Atkins P., Jones L., Laverman L. - Chemical Principles_ The Quest for Insight (Part 1) - General Chemistry (6th edition).pdf 2

6.2 Ion–Dipole Forces

• Intermolecular Interactions: Different types of interactions summarized; ionic interactions > dipole interactions > covalent bonds.

• Formation of Condensed Phases: Attractive forces pull molecules together, while repulsion dominates at short separation.

6.2.1 Ionic Interactions in Water

• Ionic Solids: Held together by strong electrostatic attractions.

• Water and Ionic Solids: When an ionic solid is added to water, water molecules surround each ion, leading to dissolution.

• Hydration: The process where water molecules cluster around solute ions is a result of the polar character of H2O.

  • Cation: Water molecules orient with O atoms (partially negative) facing the cation.

  • Anion: H atoms (partially positive) of water molecules are attracted to the anion.

• Ion-Dipole Interaction: Hydration is an example of this type of interaction due to attraction between ions and polar water molecules.

6.2.2 Potential Energy of Interaction

• Equation: The potential energy decreases as EP = - (z! * !)/(r^2)

  • Where z is the charge number and ! is the dipole moment of water molecules.

• Dependence on Distance: Ion-dipole interactions act over a shorter range compared to ion-ion interactions.

6.2.3 Hydration Specifics

• Hydration Strength: Depends on ion size; smaller cations attract water molecules more effectively.

  • Examples: Li+ and Na+ are hydrated, while larger cations such as K+ are not as much.

• Charge Impact: Higher charge leads to stronger hydration, observed when comparing Ba2+ and K+.

6.3 Dipole–Dipole Forces

• Polar Molecules: Exhibit dipole-dipole interactions, where opposite charges attract each other.

• Energy Dependence: The interaction strength is proportional to the distance r. Doubling distance increases separation impact on strength by 23 (8 times weaker).

• Gas Phase Behavior: Rapid rotation of polar molecules decreases strong dipole-dipole interactions due to cancellations of partial charges.

• Potential Energy Formula:Ep = ((!1^2 * !2^2)/r^6).

  • Doubling the distance reduces interaction strength by a factor of 64.

6.3.1 Boiling Point Predictions

• Stronger Interactions and Boiling Points: Higher boiling points correlate with stronger intermolecular forces (dipole-dipole).

6.4 London Forces

• Attraction Between Nonpolar Molecules: London forces arise due to instantaneous dipole moments.

• Electron Distribution: Electrons in nonpolar molecules are symmetrically distributed, leading to fleeting dipoles.

• Instantaneous Dipole Moments: Net attraction occurs as fluctuating dipoles induce further dipoles in neighboring molecules.

6.4.1 London Interaction Strength

• Potential Energy Formula: EP ∝ 1/r^6, meaning strength decreases rapidly with distance.

6.5 Hydrogen Bonding

• Polar Molecules: Certain hydrogen compounds (H2O, NH3, HF) exhibit strong hydrogen bonds due to electronegative atoms.

• Exceptionally High Boiling Points: Hydrogen bonds account for anomalies in boiling points among specific binary hydrogen compounds.

6.5.1 Biological Importance

• Hydrogen Bond Implications: Critical to the structure and function of biological molecules (proteins, DNA).

6.6 Repulsions

• Molecular Size Influence: Close proximity leads to repulsive forces based on orbital overlap due to Pauli exclusion principle.

6.7 Liquid Structure

• Molecular Movement in Liquids: Molecules are close yet mobile, leading to short-range order.

6.8 Viscosity and Surface Tension

• Viscosity: Resistance of a liquid to flow; higher viscosity indicates stronger intermolecular forces.

  • Water vs. hydrocarbons: Hydrogen bonds increase water's viscosity.

• Surface Tension: Caused by cohesive forces in a liquid, significantly higher in water due to hydrogen bonding.

6.9 Classification of Solids

• Types of Solids: Crystalline (orderly) vs. amorphous (random); metals vs. ionic, network, variety in properties.

6.10 Molecular Solids

• Molecular Solid Properties: Generally softer and have lower melting points compared to ionic and network solids.

6.11 Network Solids

• Structural Composition: Atoms covalently bonded through the solid, usually very hard with high melting points.

6.12 Metallic Solids

• Metallic Bonding: Cations in a sea of electrons give metals distinct properties like conductivity, malleability.

6.13 Unit Cells

• Understanding Crystal Structure: Unit cells define the arrangement of particles in a crystal lattice allowing for systematic representations.

6.14 Ionic Structures

• Section Overview: Ionic structures depend on packing of varying radii ions ensuring overall neutrality.

  • Examples: Rock-salt and cesium chloride structures demonstrate 3D arrangements effectively.

6.15 Liquid Crystals

• Properties and Applications: Flow like liquids but organized like solids; used in electronic displays due to their responsiveness.

6.16 Ionic Liquids

• Characteristics: Liquid salts that remain liquid at room temperature, reducing carcinogenic vapors while dissolving organic materials.