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Lecture 10

The Liquid and Solid States

Shaun Williams, PhD

Examples of Solids, Liquids, and Gases

Photographs of a solid, a liquid, and a gas.

Changes of State

Liquid-Gas Phase Changes

  • Evaporation (vaporization)
    • At the surface of a liquid, some molecules may have sufficient kinetic energy to escape into the gas state.
    • Heat is required to maintain the temperature needed for evaporation.
      • Evaporation is an endothermic process.
An imagine depicting water molecules transitioning from the condensed liquid phase to the gas phase.

Liquid-Gas Phase Changes (cont.)

Images showing that in a closed container holding liquid water, water vapor builds up over time.

Equilibrium

Boiling Point

  • The temperature at which boiling occurs
    • Boiling occurs when the vapor pressure equals the external pressure of the atmosphere
  • Normal boiling point
    • Occurs when the atmospheric pressure is exactly 1 atm
A plot showing three different compounds vapor pressures as a function of temperture.

Liquid-Solid Phase Changes

Liquid-Solid Phase Changes (cont.)

Freezing

A photograph of an ice cube melting and a image of molecular ice crystal and liquid water.

Melting

  • Melting
    • Phase change from solid to liquid
    • Reverse of freezing
    • Opposite process of freezing
    • Also called fusion
  • Melting point
    • Same temperature as freezing point
Images showing solid water transitioning to them more disordered liquid water and crystalling sodium chloride transitioning to the more dissordered liquid salt.

Solid-Gas Phae Changes

  • Sublimation
    • Evaporation of a solid
    • Occurs when a solid has a high vapor pressure
    • Solid can change directly from a solid to a gaseous state without going through the liquid state
  • Deposition
    • Can go directly from gas to solid without passing through the liquid state
    • Reverse of sublimation
A photograph of gaseous iodine depositing on the bottom of a cold evaporating dish.

Cooling Curve

A plot of the temperature on a sample over time as heat is removed at a constant temperature. The temperature drops except when the phase changes are occuring (temperature is constant during phase changes).

Heating Curve

A plot of the temperature on a sample over time as heat is added at a constant temperature. The temperature increases except when the phase changes are occuring (temperature is constant during phase changes).

Energy Changes

Intermolecular Forces

A photo showing the yellowing chlorine gas, reddish-brown bromine liquid, and purple iodine crystals.

The Forces

Type of Force Type of Interaction Occurrence
London dispersion force A temporary dipole in one molecule induces the formation of a temporary dipole in a nearby molecule and is attracted to it. All atoms and molecules
Dipole-Dipole Force Polar molecules (permanent dipoles) attract one another Polar molecules
Hydrogen-Bonding Force Two dipoles, one containing hydrogen to an electronegative element and the other containing an electronegative element, attract one another. Polar molecules containing unpaired molecules and a hydrogen bonded to nitrogen, oxygen, or fluorine

London Dispersion Forces

London Dispersion Forces (cont.)

A series of images showing electron clouds fluctuating in time and forming a short lived bulge and therefore a small dipole moment.

Dipole-Dipole Forces

  • Attraction between polar molecules
  • Occurs when the partially positive end of one molecule attracts the partially negative end of another molecule
  • Generally stronger than London dispersion forces
An image showing the dipole moment on one SO2 molecule attracting the dipole moment on a neighboring SO2 molecule.

Boiling Points

A plot showing the boiling points of groups 17 diatomics, group 14 with hydrogen, and group 18 atoms. In each case the trend has the lighter species having a lower boiling point and the general trend is linear.

Hydrogen Bonding

  • Special type of dipole-dipole force
  • Only occurs in molecules that contain hydrogen bonded to a small, highly electronegative element
  • Stronger than a regular dipole-dipole force
  • Important force in living systems by stabilizing molecular shapes
A plot showing the boiling points of molecules involving groups 14, 15, 16, and 17 with hydrogen atoms. The molecules H2O, HF, and NH3 boiling at considerably higher temperatures than their family's linear trend predicts.

Hydrogen Bonding Examples

Images of hydrogen bonding in water, ammonia in water, ammonia by itself, and hydrogen fluoride.

Trends in Intermolecular Forces

  • Remember:
    • London forces exist between all atoms and molecules.
    • Dipole-dipole moments exist in only polar compounds.
    • Hydrogen bonds only exist in polar compounds that contain hydrogen.
  • In terms of strength (magnitude), intermolecular forces compare as shown below:
All molecules have London Dispersion Forces, polar molecules have dipole-dipole forces, and polar molecules with HF, HO, or HN bonds have hydrogen bonding.

\[ \text{London Disp.} \lt \text{Dipole-Dipole} \lt \text{Hydrogen Bonds} \]

Properties of Liquids

  • Are related to the distance between particles and to intermolecular forces
  • Particles in a liquid are much closer together than the particles in a gas
  • Liquid particles are not fixed, as they are in a solid
  • Three common properties of liquids:
    • Density
    • Viscosity
    • Surface tension
A photograph of ice cubes melting.

Density

Viscosity

Surface Tension

  • The amount of work required to increase the surface area of a liquid by a unit amount
  • Causes a liquid surface to behave like a stretched membrane
  • The greater the intermolecular forces in a liquid, the greater the surface tension
  • Surface tension decreases as temperature increases
An image showing the hydrogen bonding of a water molecule in the center of a liquid versus a water molecule at the surface of the liquid (at the surface it has half as many hydrogen bonds.

Meniscus

A picture of the concave surface of water in a test verus the convex surface of mercury in a test tube.

Properties of Solids

Amorphous and Crystalline Solids

Amorphous and Crystalline Solids Examples

A picture of quartz crystal showing the appearance of a crystalling solid and another photograph of broken window glass showing the random nature of the fractures indicating the amorphous nature of window glass.

Crystals and Crystal Lattices

Closest-Packed Arrangements

Photographs of the surface of a pineapple, the cells in a bee hive, the celling in a leaf, and the dimples on a golf ball. Each is an example of packing things into a limited space.

Types of Crystalling Solids

Types of Crystals

Type of Solid Fundamental Particles Attractive Forces Properties
Metallic Atoms Attractions between nuclei and delocalized electrons Low melting point & soft; or high melting point & hard; good heat & electrical conductors; malleable & ductile
Ionic Cations and Anions Ionic bonds High melting point; hard, brittle; nonconductors when solid; electrical conductors when melted
Molecular Polar Molecules Dipole-dipole forces Low to moderate melting point; variable hardness; may be brittle; nonconductors
Nonpolar Molecules London dispersion forces Low melting point; soft; poor heat conductors; electrical insulators
Netword Atoms Covalent bonds Very high melting point; very hard; somewhat brittle; non- or semiconductors

Metallic Solids

  • Valence electrons move freely through all parts of a metal
  • Attractions between atoms of a metal are delocalized, and therefore it is easy to move atoms by applying force
  • Ductile and malleable
A photograph of copper, zinc, and brass metals.

Alloys

A image of the mixed atoms in a titanium alloyed with nickel.

Ionic Solids

  • Contain cations and anions arranged in crystalline solids
  • Electrostatic forces (ionic bonds) hold together ionic crystals
  • High melting points
  • Hard and brittle
  • Solids are not electrical conductors, but melted or dissolved, they become good conductors
A image of crystals of sodium chloride, zinc sulfide, calcium fluoride, and cesium chloride.

Superconductors

An image of a generic superconductor made up of yttrium, barium, copper, and oxygen.

Molecular Solids

Example of a Molecular Solid

An image of carbon dioxide molecules arranged in a crystal that is dry ice.

Network Solids

  • Consists of a giant molecule that forms the entire crystal
  • Formed by metalloids or carbon
  • Strong covalent bonds connect the atoms in a network solid
  • Poor electrical conductors
  • High melting points
  • Very hard
  • Have very stable, three–dimensional structures
  • Many have a diamond structure, or some derivative of it
A photograph of a diamond, a network solid where all the carbon atoms are covalently bound to four neighboring carbon atoms.

Example of a Network Solids

A photograph of sand paper, the sand on which is silicon dioxide (a network solid). A photograph of graphite in pencil lead, which is a network solid of carbon where each carbon atom is covalently bonded to three neighbors on a plane.

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