Chapter 10: Liquids, Solids, and Phase Changes

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Chapter 10 Liquids, Solids, and Phase Changes
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Introduction
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Introduction

• Gases
• Gas particles act independent of one another
• Attractive forces are very weak
• Particles are free to move randomly
• Occupy whatever space available
• Liquids and solids are different from gases in
  that they have strong Attractive Forces between
  molecules.

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Bonding in Molecules

• Ionic bond Cation anion.
• Covalent bond Sharing the electrons between
  atoms.

Polar covalent bond Nonpolar covalent bond
Difference in Electronegativity
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Electronegativity Scale
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Polar Covalent Bonds
Bond dipole A polar bond has two polar ends a
positive end () and a negative end (-). It is
often represented by an arrow with a cross at one
end (positive end) to indicate the direction of
electron displacement.
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Polar Molecules

• Polar Molecules
• Just as bonds can be polar, molecules as a whole
  can be polar
• Net sum of individual bond polarities and
  lone-pair contributions

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• Example 1Tell which of the following
  compounds are likely to have a dipole moment and
  show the direction of each.
• a) SF6 b) CHCl3
• c) CH2Cl2 d) CH2CH2

Molecules having a symmetrical plane are
non-polar molecules !!(net sum of bond dipole is
zero)
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Intermolecular Forces

• intermolecular forces as a whole are usually
  called
• Van der Waals forces

all are electrical in origin and result from the
mutual attraction of unlike charge or mutual
repulsion of like charges
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Van der Waals forces

• Four main types
• Dipole-dipole
• Ion-dipole
• Dispersion forces
• Hydrogen bonding
• Not every molecule has all of four types of
  forces.
• Then, for a specific molecule, which types of
  forces does it exhibit?

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Dipole-dipole

• Neutral but polar molecules experience
  dipole-dipole forces as a result of electrical
  interactions among dipoles on neighboring
  molecules.

repulsive
attractive
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Dipole-dipole

• Neutral but polar molecules experience
  dipole-dipole forces

These forces are weak 3-4 kJ/mol and only
significant if molecules are close
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Ion-dipole

• Ion-dipole force is the result of electrical
  interactions between an ion and the partial
  charges on a polar molecule

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Ion-dipole

• Ion-dipole force occurs between an ion and a
  polar molecule

Particularly important in aqueous solutions of
ionic substances such as NaCl, in which polar
water molecules surround the ions
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London Dispersion Forces

• Result from the motion of electrons

instantaneous dipole can induce temporary dipoles
in neighboring molecules
Every molecule has London dispersion forces.
More electrons a molecule has the stronger the
dispersion forces
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Hydrogen Bonding

• Attractive interaction between a hydrogen atom
  bonded to
• a very electronegative atom (O, N, F) and an
  unshared
• electron pair on another electronegative atom (O,
  N, F)

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Summary
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• Example 5 Identify the likely kinds of
  intermolecular forces in the following
• A)   HCl
• B) CH3CH3
• C) CH3NH2
• D) Kr

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• Example 6. Of the substances Ar, Cl2, CCl4 and
  HNO3 which has
• a)  The largest dipole-dipole forces?
• b)   The largest hydrogen-bond forces?
• c) The smallest dispersion forces?
  

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Properties of Liquids

• Viscosity
• The measure of a liquids resistance to flow

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Viscosity

• Substances with small non-polar molecules
• have weak intermolecular forces and low
• viscosities (free flowing)
• More polar substances have stronger
• intermolecular forces and have higher
• viscosities

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Properties of Liquids

• Surface Tension
• The resistance of a liquid to spread out and
  increase its surface area

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Surface Tention
Surface molecules feel attractive forces on only
one side and are drawn in toward the
liquid Interior molecules are drawn equally in
all directions
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Capillary Action

• capillary action is the ability of a liquid to
  flow up a thin tube against the influence of
  gravity
• the narrower the tube, the higher the liquid
  rises
• capillary action is the result of the two forces
  working in conjunction, the cohesive and adhesive
  forces
• cohesive forces attract the molecules together
• adhesive forces attract the molecules on the edge
  to the tubes surface

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Capillary Action

• the adhesive forces pull the surface liquid up
  the side of the tube, while the cohesive forces
  pull the interior liquid with it
• the liquid rises up the tube until the force of
  gravity counteracts the capillary action forces

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Meniscus

• the curving of the liquid surface in a thin tube
  is due to the competition between adhesive and
  cohesive forces
• the meniscus of water is concave in a glass tube
  because its adhesion to the glass is stronger
  than its cohesion for itself
• the meniscus of mercury is convex in a glass tube
  because its cohesion for itself is stronger than
  its adhesion for the glass
• metallic bonds stronger than intermolecular
  attractions

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Phase Changes

• Physical form changes but chemical identity does
  not change
• Fusion (melting) solid ? liquid
• Freezing liquid ? solid
• Vaporization liquid ? gas
• Condensation gas ? liquid
• Sublimation solid ? gas
• Deposition gas ? solid

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Thermochemistry in Phase Changes
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Evaporation of liquids
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Vapor Pressure
Vapor Pressure The pressure exerted by gas
molecules in equilibrium with liquid at
constant temperature.
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Vapor Pressure

• Numerical value of Vapor Pressure depends on
• a)      Magnitude of intermolecular forces
• The smaller the forces the higher the vapor
  pressure, loosely held molecules escape
  easily
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• b)      Temperature
• The higher the temperature, the higher the
  vapor pressure, larger fraction of molecules
  have sufficient kinetic energy to escape

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Heating Curve for H2O
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Boiling Point

• 1. Boiling point when the vapor pressure of a
  liquid rises to the point where it becomes equal
  to the external pressure
• Normal boiling point
• 2. The temperature at which boiling occurs when
  the external pressure is exactly 1 atm.

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Phase Diagram
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Phase Diagrams

• Triple Point The only condition under which all
  three phases can be in equilibrium with one
  another.
• Critical Temperature (Tc) The temperature above
  which the gas phase cannot be made to liquefy at
  any pressure.
• Critical Pressure (Pc) The minimum pressure
  required to liquefy a gas at its critical temp.
• Supercritical Fluid Neither true liquid nor
  true gas
• Normal boiling and melting point always at 1 atm

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• Example 7
• Can you label the following?
• a) solid region
• b) Liquid region
• c)  Gas region
• d)  Normal boiling point
• e)  Normal melting point
• f)   Triple point
• g)  Supercritical fluid region
• h) Critical point, what is the critical pressure
  and temperature

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Summary

• Properties of liquids Viscosity surface
  tension.
• Phase changes and Phase diagram.
• Vapor pressure and normal boiling/melting point.