States of Matter and Intermolecular Forces

Kinetic Molecular Theory

  • Explains states of matter based on the idea that matter is composed of tiny particles in constant motion.
  • Helps explain observable properties and behaviors of solids, liquids, and gases.

Differences in States of Matter

  • Solid:
    • Definite volume and shape.
    • Particles packed in fixed positions.
    • Particles are not free to move nor flow
  • Liquid:
    • Definite volume but indefinite shape.
    • Particles close together but not in fixed positions.
    • Particles are free to move or flow.
  • Gas:
    • Neither definite volume nor definite shape.
    • Particles are at great distances from one another.
    • Particles are free to move or flow.

Phase Changes

  • Melting: Solid to liquid
  • Freezing: Liquid to solid
  • Vaporization: Liquid to gas
  • Condensation: Gas to liquid
  • Sublimation: Solid to gas
  • Deposition: Gas to solid

Intermolecular Forces

  • Hold molecules together; can be attractive or repulsive.
  • Accountable for the properties of substances.
  • Explain why substances exist as solids, liquids, or gases at room temperature.
  • Types:
    • London Dispersion Forces
    • Dipole-dipole
    • Hydrogen bonding

Types of Intermolecular Forces (In Increasing Strength)

  • London Dispersion Forces:
    • Interacting Particles: All molecules (especially nonpolar molecules).
  • Dipole-dipole:
    • Interacting Particles: Polar molecules.
  • Hydrogen Bonding:
    • Interacting Particles: N, O, or F (FON) bonded with H atom.

London Dispersion Forces

  • Weakest attractive force, formed due to temporary dipoles induced in non-polar molecules.
  • Temporary dipoles occur when electron clouds of nonpolar molecules are distorted by external electric fields (polarizability).
  • London dispersion forces become stronger as the atom or molecule gets larger (has more electrons).
    • Example:
    • Fluorine, chlorine are both gases at room temp.
    • Bromine is a liquid
    • Iodine is a solid

Dipole-Dipole Forces

  • Attractive forces existing between polar molecules (molecules with a dipole moment), such as HCl.
  • Unequal sharing of electrons between H and Cl atoms creates partial positive and partial negative poles (dipole).

Hydrogen Bonding

  • Special type of dipole-dipole interaction between the hydrogen atom (partial positive) in a polar molecule and highly electronegative atoms F, O, N (partial negative atom) in another molecule or the same molecule.
  • Common example: H2O (water).
  • Strongest intermolecular force.

Heating/Cooling Curves

  • Graphical representations of the correlation between heat input and the temperature of a substance.
  • Used to determine the melting point and boiling point of a substance at a single, constant pressure.

Structure of Heating Curve

  • Shows phase transitions (solid, liquid, gas) with plateaus at melting and boiling points.

Structure of Cooling Curve

  • Reverse of heating curve, showing phase transitions and condensation/freezing points.

Phase Diagrams

  • Depicts phase changes at all temperatures and pressures.

Structure of a Phase Diagram

  • Key points:
    • Triple point: all states exist in equilibrium
    • Critical point: supercritical fluid has properties of a gas and a liquid
    • Vaporization/condensation point
    • Melting/freezing point
    • Sublimation/deposition point

Endothermic vs. Exothermic

  • Endothermic (absorbs heat):
    • Melting, boiling, and sublimation
    • The surrounding area will feel cold
  • Exothermic (releases heat):
    • Freezing, condensation, and deposition
    • The surrounding area will feel warm
  • You need enough heat to break intermolecular forces between liquid water to get it into the gaseous state.

Heating/Cooling Curves and Endothermic/Exothermic Processes

  • Heating curves are endothermic, while cooling curves are exothermic.

Dynamic Equilibrium

  • Occurs in a closed system (sealed container) where the rate of vaporization of a substance equals the rate of condensation.
  • Molecules are constantly changing phase - "dynamic"
  • The total amount of liquid and vapor remains constant - "equilibrium"

Vapor Pressure

  • The pressure exerted by a substance when it is in dynamic equilibrium.
  • Vapor pressure must equal atmospheric pressure for boiling to occur.

Vapor Pressure and Intermolecular Forces

  • Different substances have different vapor pressures at a given temperature due to the strength of their intermolecular forces.
  • Molecules must have enough kinetic energy to overcome intermolecular forces to escape into the vapor phase.

Vapor Pressure and Volatility

  • Substances with strong intermolecular forces have lower vapor pressures and are less volatile.
  • Substances with weak intermolecular forces have higher vapor pressures and are more volatile.

Surface Tension

  • A force that tends to pull adjacent parts of a liquid’s surface together, thereby decreasing surface area to the smallest possible size.
  • Results from stronger attractive forces between particles of a liquid (polar molecules, liquid metals).
  • Water is a common example.

Viscosity

  • The resistance to flow.
  • Liquids which flow very slowly (glycerin, honey) have high viscosities.
  • Liquids which flow very readily (ether, gasoline) have low viscosities.