Kinetic Molecular Theory & Intermolecular Forces

• The Kinetic Molecular Theory (KMT) is founded on the idea that all states of matter are composed of molecules possessing kinetic energy.
• KMT suggests that the motion, arrangement, and amount of kinetic energy differ among molecules in gases, liquids, and solids.

KMT and States of Matter

• Gas: No definite shape or volume.
• Liquid: Definite volume, but no definite shape.
• Solid: Definite shape and volume.

Intermolecular Forces

• Intermolecular forces include London dispersion forces, dipole-dipole forces, ion-dipole forces, and hydrogen bonding.

London Dispersion Forces (LDF)

• LDF are weak attractions explaining interactions between nonpolar molecules.
• Dispersion forces increase with molecular mass but decrease as distance increases.

Dipole-Dipole Forces

• Dipole-dipole interaction is the attraction between bond dipoles in different molecules.
• Bond dipoles stem from unequal electron sharing in covalent bonds.

Ion-Dipole Forces

• These forces exist when polar molecules are attracted to ions.
• The positive pole is attracted to anions, while the negative pole is attracted to cations.
• An example is dissolving NaCl in H₂O.

Hydrogen Bonding

• Hydrogen bonds form a bridge between two highly electronegative atoms (F, O, or N) and involve covalent bonds with other hydrogens.
• Molecules with hydrogen bonded to F, O, or N exhibit significant hydrogen bonding ability.

Properties of Liquids

• Properties of liquids are affected by: Surface Tension, Viscosity, Capillary Action, Evaporation, Vapor pressure, and Boiling point

Surface Tension

• Surface tension is the force causing surface molecules of a liquid to tighten their hold, creating a thin membrane effect.
• High surface tension is found in substances with strong intermolecular attractions.
• Surface tension decreases as temperature increases.

Viscosity

• Viscosity is a liquid's resistance to flow.
• High viscosity liquids are "thick," while low viscosity liquids are "thin."
• Viscosity is affected by intermolecular forces, such as London dispersion forces.

Capillary Action

• Capillary action is the spontaneous rise of a liquid in a narrow tube.
• It results from cohesive forces within the liquid and adhesive forces between the liquid and container walls.

Evaporation

• Evaporation is the escape of molecules from the surface of a liquid.
• It is a direct result of molecular motion.

Vapor Pressure

• Vapor pressure results when a liquid vaporizes in a closed container, saturating the space above with vapor, leading to equilibrium between liquid and vapor.
• Rates of evaporation and condensation become equal, stabilizing liquid and vapor amounts.

Boiling Point

• Boiling point is the temperature at which a liquid's vapor pressure equals the external pressure.

Matter in the Solid Phase

• Solid particles are strongly attracted, keeping them in fixed positions, hence solids have definite shapes and volumes.
• Particles in solids possess limited kinetic energy, with movement restricted to vibrational motion.
• Increased temperature in solids increases particle kinetic energy, potentially causing movement but within attractive force limits.
• Increased temperature causes expansion, but the volume change is small.
• External pressure has little effect on a solid's volume; solids are incompressible.
• Sufficient heat breaks attractive forces, allowing particles to become mobile like liquid particles, and the solid melts.
• Melting point is the temperature a solid melts from solid to liquid.
• Stronger attractive forces correlate with higher melting points.
• Heat of Fusion is the heat required to completely melt a solid at its melting point.

Vapor Pressure of a Solid

• Solid particles may gain enough energy to escape the surface and become gas or vapor without turning into liquid.
• Sublimation is the change from solid to gas without becoming liquid.
• Equilibrium is established between a solid and its vapor when a solid that sublimes at an appreciable rate is placed in a closed container at a constant sublime temperature.

Classes of Solids Include:

• Amorphous solids
• Crystalline solids

Amorphous Solids

• Amorphous solids do not maintain a consistent form and have randomly arranged particles.
• Examples include asphalt, rubber, glass, and plastic.
• They are formed by rapid cooling of liquids, which freezes the particles in a disordered state.
• Show short order in arrangement of particles
• These are sometimes called pseudo solids or super cooled liquids.

Crystalline Solids

• Crystalline solids have orderly arrangements of atoms, ions, or molecules.
• Crystals exhibit regular shapes reflecting internal particle arrangement.
• Show long order in arrangement of particles
• Called true solids because they have true properties of solids.

Four types of crystalline solids:

• Metallic crystals
• Ionic crystals
• Molecular crystals
• Covalent crystals

Metallic Crystals/Solids

• Held together by metallic bonds.
• Examples include iron, gold, and silver.
• Made up of metal atoms.
• Have high melting points.
• Malleable and ductile.
• Good conductors of electricity.

Ionic Crystals/Solids

• Held together by ionic bonds.
• Examples include NaCl, CaCO₃, and MgO.
• Made up of metal and nonmetal atoms.
• High melting points.
• Brittle.
• Insulators in solid-state but conductors when molten.

Molecular Crystals/Solids

• Held together by covalent bonds.
• Examples include ice, sugar, and carbon dioxide.
• Made up of molecules.
• Low melting points.
• Soft in nature.
• Insulators.

Covalent Crystals/Solids

• Held together by covalent bonds.
• Examples include diamond, graphite, and silicon dioxide.
• Made up of nonmetal atoms.
• High melting points.
• Hard in nature.
• Poor conductors or semiconductors of electricity.