Chapter 11 - Liquids and Intermolecular 

• Intermolecular - Forces between molecules.
• ==Boiling== breaks intermolecular forces.

Properties that reflect intermolecular forces

• Boiling point
• Melting point
• Viscosity
• Surface tension
• Capillary action
• Water has a ==high surface tension==.

States of matter

• Gas
• Liquid
• Solid
  • ==Gases and liquids== are called fluids.
  • ==Liquids and solids== are called condensed states, they have strong forces.

Intermolecular forces (Weakest to strongest)

• Dispersion forces or London forces - Only occurs in non-polar molecules. It is constantly shifting to a different set of temporary forces.
  • The tighter the molecules, the lower the surface area they have, and the lower their boiling point is.
  • Ex. Neopentane has a ==lower boiling point== than pentane because neopentane has a lower surface area.
• Dipole-dipole interactions - They form permanent dipoles. Occurs in polar molecules.
  • Bad interactions - 2 positive or 2 negative banging into each other.
• Hydrogen bonding - Strongest force. It can only be formed with a hydrogen atom bonding with a nitrogen, oxygen, or fluorine atom.
• Crystalatus - Smallest defined unit that repeats inside a molecule.
• Water is the only liquid that freezes from the top.

Ion-dipole interactions

• Found in solutions of ions.
• Can only occur with polar compounds.

Relative strengths of intermolecular forces

• When 2 molecules have comparable moral masses and shapes, dispersion forces are equal.
• When 2 molecules have very different molar masses and there's no H-bonding, dispersion force determines the substance with stronger attractions.

Properties affected by intermolecular forces

• Viscosity - Resistance of a liquid to flow. Increases with strong forces, decreases with higher temperature.
• Surface tension - Water acts as if it has a skin be of the extra forces on the surface allowing water to bead up when in contact with nonpolar surfaces.
• Capillary action - The rise of liquid up narrow tubes.
  • Cohesive forces - Intermolecular forces that bind similar molecules to one another.
  • Adhesive forces - Intermolecular forces that bind a substance to a surface.

Phase changes

• Phase change - Conversion from one state to matter to another.
  • Melting / Fusion - Solid to liquid, endothermic.
  • Freezing - Liquid to solid, exothermic.
  • Vaporization - Liquid to gas, endothermic.
  • Condensation - Gas to liquid, exothermic.
  • Sublimation - Solid to gas, endothermic.
  • Deposition - Gas to solid, exothermic.
• 

Heating Curves

• Heating curve - Graph of temperature (y) and the heat added (x).
• 

Vapor pressure

1. As temperature increases, more molecules are able to have enough energy to become a gas.

• P = nRT/V
• P = MRT
  • M - molarity
  • R - gas constant
• Vapor pressure - How much of a liquid evaporates at a certain pressure.
• 
• At any temperature, some liquid molecules have enough energy to escape the surface and become a gas.

Vapor pressure curves

• 

• Natural log of the vapor pressure - It’s inversely proportional to its temperature.

• Clausius-Clayperon equation

  • We can find ΔH of vaporization if we know the vapor pressure and the temperature at one point.
  • We can find the pressure at point 1 when we know the ΔH of vaporization and the temperature at point 2.

• Formula simplification

• 

Phase diagrams

• Phase diagram - A graph that shows the states of matter under conditions of temperature and pressure.
• 
• Triple point - The point where all three states of matter coexist.
• Critical point - The point at which no amount of pressure alone can liquify the gas.
  • Here you can’t tell the difference between a gas and a liquid.
• PHASE DIAGRAM OF WATER
• 
• Water can melt only by pressure.
• The slope of the melting curve is negative, meaning that as the pressure goes up, the melting point goes down.
• CARBON PHASE DIAGRAM
• 
• Carbon has two triple points.
  • Triple points are always between the liquid and gas or between two solids.