Intermolecular_Forces__Liquids_and_Solids

Unit 7: States of Matter and Intermolecular Forces

Intramolecular vs Intermolecular Forces

• Intramolecular Forces:

  • Attractive forces within molecules, referred to as "bonds".

  • Responsible for chemical properties.

• Intermolecular Forces:

  • Forces existing between molecules.

  • Account for physical properties.

Intermolecular Forces (IMFs)

• Higher IMFs lead to:

  • Elevated melting points of solids.

  • Higher boiling points of liquids.

• A gas condenses to a liquid at low temperature due to decreased kinetic energy, making IMFs significant.

Differences among States: Solid, Liquid, Gas

• State of a substance depends on:

  • Kinetic energy of particles.

  • Strength of attractions between particles.

Types of Intermolecular Forces (IMFs)

Ion-Dipole Forces

• Exist between an ion and a polar molecule.

• Example: NaCl dissolving in water (negative end of water attracted to positive sodium ion).

Dipole-Dipole Forces

• Found between neutral polar molecules in solid or liquid states.

• Develop when two polar molecules are in proximity.

London Dispersion Forces

• Result from attractions between nonpolar molecules forming temporary dipoles.

• Can occur when the positive nucleus of one atom attracts the electrons of another.

• Present for all atoms/molecules; weakest of IMFs.

Hydrogen Bonding

• A strong type of dipole-dipole force.

• Exists between H atoms bonded to F, O, or N.

• Small size of H allows close proximity to nonbonding electron pairs of electronegative atoms.

• Responsible for unique properties of water, such as high boiling point and low density in solid form.

Polarizability and Dispersion Forces

• Polarizability:

  • Degree to which a dipole can be induced in nonpolar species.

  • Increases with greater mass.

Classification of IMFs

• Hydrogen Bonding:

  • Example: Water (H₂O) - the hydrogen atoms form hydrogen bonds with oxygen atoms in neighboring water molecules.

• London Dispersion:

  • Example: Iodine (I₂) - temporary dipoles can occur between nonpolar iodine molecules, leading to weak intermolecular forces.

• Ion-Dipole:

  • Examples:

    • Ionic bonding (e.g., KBr in water)

• Dipole-Dipole:

  • Examples: H₂S, CH₃Cl

Evaporation and Temperature

• Observation exercise with liquids:

  • Record behaviors when drops are placed alone and together.

  • Develop an experiment to determine polarity of a third liquid (C).

IMF Question

• Order halogens by increasing boiling points, explaining reasoning. The halogens can be ordered by increasing boiling points as follows: Fluorine (F₂) < Chlorine (Cl₂) < Bromine (Br₂) < Iodine (I₂). This trend can be explained by the molecular weight and the strength of the London dispersion forces, which increase as the size and mass of the molecules increase.

• Fluorine has the lowest boiling point because it is the lightest and has the weakest dispersion forces due to low molar mass.

• Chlorine is heavier than fluorine, resulting in stronger dispersion forces and thus a higher boiling point.

• Bromine, being more massive, has even stronger London forces, leading to an increased boiling point.

• Iodine has the highest boiling point among them due to its large size and high molecular mass, which enhance the dispersion forces considerably.

Temperature Definitions

• Melting Point (MP): Temperature where solid converts to liquid.

• Freezing Point (FP): Temperature where liquid turns to solid.

• Boiling Point (BP): Temperature where liquid forms gas.

• Condensing Point (CP): Temperature where gas converts to liquid.

Vapor Pressure

• Equilibrium Vapor Pressure: Occurs when vapor produced equals the rate of condensation in a closed container.

• Low molar mass and weak IMFs lead to high vapor pressure; increasing temperature raises vapor pressure of liquids/solids.

Properties of Liquids

• Viscosity: Resistance of a liquid to flow; increases with molar mass.

• Surface Tension: Energy to increase surface area; arises from IMFs.

• Cohesion: IMFs binding similar molecules together.

• Adhesion: IMFs binding different substances together.

• Volatility: Tendency of a substance to vaporize.

Phase Changes

• Heating curves represent temperature vs. heat added.

• Lines BC and DE correspond to phases undergoing transitions.

Critical Temperature and Pressure

• Critical Temperature: Maximum temperature for liquid existence; above it, only gas phases exist.

• Critical Pressure: Needed for liquefaction at critical temperature.

• Nonpolar and low molar mass substances generally have low IMFs, critical temperatures, and pressures; polar substances and higher molar mass materials show opposite properties.

Phase Diagrams

• A graphical representation of solid, liquid, and gaseous phase equilibria of a substance.

• Evaluating system changes upon adjustments in pressure and temperature at given points.

Polar vs. Nonpolar Molecules

A molecule is considered polar when it has a net dipole moment due to the presence of polar bonds and an asymmetrical shape. This typically occurs when:

• There is a significant difference in electronegativity between the bonded atoms, leading to an unequal sharing of electrons.

• The molecular geometry does not allow the dipoles to cancel each other out (e.g., water (H₂O) is polar).

In contrast, a molecule is classified as nonpolar when:

• The electronegativity difference between the atoms is negligible, resulting in an even distribution of electron density.

• The molecule has a symmetrical shape that allows dipoles to cancel each other out (e.g., carbon dioxide (CO₂) is nonpolar, despite having polar bonds).

Understanding the polarity of a molecule is crucial as it influences its interactions, solubility, and overall behavior in different environments.