Condensed States of Matter:

Liquids and Solids

Chapter 14

Condensed States

• Liquids and solids = condensed states because they have significantly higher densities than gases

Gas Liquid Solid

highly compressible

only slightly compressible

only slightly compressible

low density high density high density

fills container completely

does not expand to fill container -

has definite volume

rigidly retains its volume

assumes shape of container

assumes shape of container retains own shape

rapid diffusion slow diffusion

extremely slow diffusion - only

occurs at surfaces

high expansion on heating

low expansion on heating

low expansion on heating

KMT of Matter

• According to the kinetic molecular theory, the state of a substance at room temperature depends on the strength of the attractions between its particles.

Forces

• Intramolecular forces are between the atoms within a molecule = bonds

• Intermolecular forces are between molecules

Intramolecular Forces

• Covalent Bonds – between nonmetals; sharing of electrons

• Ionic Bonds – between metals and nonmetals; transfer of electrons to form ions

• Metallic Bonds – between metals; sea of electrons

Intermolecular Forces

1. Dipole-dipole attractionsa. Hydrogen bonding

2. Ion-dipole attractions3. London dispersion forces

Dipole-Dipole Attractions• Attractions due to permanent

dipoles in polar molecules• Remember, a dipole is

created when positive and negative charges are separated by some distance.

• Only 1% as strong as covalent or ionic bonds and even weaker if distance between charges increases.

Hydrogen Bonds

• Unusually strong dipole-dipole attractions that occur among molecules in which hydrogen is bonded to a highly

electronegative atom, such as F, N, or O

Ion-Dipole Attractions

• Polar molecules surround ions based on their attraction for the charge on the ion

London Dispersion Forces• Explain the attraction that

exists between non-polar molecules

• Even in these molecules the electrons are not uniformly distributed at every second

• Temporary dipolar arrangement of charge creates an instantaneous dipole.

London Dispersion Forces• Instantaneous dipoles can

induce similar dipoles in neighboring atoms

Kinetic

Kinetic EnergyIntermolecular

AttractionState at Room

Temp.

low high solids

medium medium liquids

high low gases

Kinetic energy determines if particles will overcome the intermolecular forces keeping

them together. Thus, higher attractions mean higher boiling points, because a higher kinetic

energy will be needed to overcome the attraction.

Properties of Solids

• Intermolecular forces keep the particles of a solid packed together tightly- Particles are highly ordered

with fixed positions- Only particle movement =

vibrations• Liquids have similar distance

between particles

Bonding in Solids

2 basic types of solids based on the nature of the particles that make them up:

1. Crystallinea. Atomicb. Ionicc. Covalent-network

2. Amorphous

Crystalline Solids• Solid in which the

representative particles are in a highly ordered, repeating pattern called a crystal

• Can be studied as unit cells, small representative units that repeat throughout the structure

Bonding in Crystalline Solids

The physical properties of solids, such as hardness,

electrical conductibility, and melting point, depend on the kind of particles that make up the solid and on the strength

of the attractive forces between them. (Includes ionic,

covalent, and atomic substances)

Covalent-Network Solids

• A type of crystalline solid in which strong covalent bonds forma a network extending throughout the solid

• Have very high melting points due to strong covalent bonds

Amorphous Solids

• Solids in which the arrangement of the representative particles lacks a regular, repeating pattern

• Also known as “supercooled liquids”- Liquids cooled until

viscosity becomes so high that no flow can occur

Amorphous Solids

• Get softer when heated instead of reaching an abrupt melting point like crystalline solids

• Examples: glass, rubber, some plastics

LiquidsThe physical properties of liquids are determined by the nature and strength of the intermolecular forces

present between their molecules

Properties of Liquids

• Viscosity – resistance to motion that exists between molecules of a liquid when they move past each other– Increased intermolecular forces

yield increased viscosity

• Glycerine has lots of hydrogen bonds

• Decreased temperatures yield increased viscosity

Properties of Liquids

• Surface tension – imbalance of attractive forces at the surface of a liquid that cause the surface to behave as if it had a film across it

»Explains “beading” of liquids»As with viscosity, increased

intermolecular forces or decreased

temperatures yield increased surface tension

Phase Changes

• Changes of state are physical changes, not chemical.– Intermolecular forces break or

form, not intramolecular• Phase changes occur when

energy enters or leaves a compound.

• Energy in: solid liquid gas• Energy out: gas liquid solid

Heating/Cooling Curves• Show the changes in

temperature as energy as heat is added to (or removed from) a substance and it changes states.

Heating/Cooling Curve

Phase Changes

• During phase changes, heat is still transferred, but no temperature change takes place.

Phase Diagrams

• Shows the state of matter of a substance at increasing pressures and temperatures.

• Can be used to determine the state of matter of a substance at a particular pressure and temperature.

Phase Diagram of Water

Energy Requirements

• Molar heat of fusion: the energy required to melt one mole of a solid

• Molar heat of vaporization: the energy required to vaporize one mole of a liquid

• Molar heats of vaporization are higher because more intermolecular forces will have to be overcome to go from a liquid to a gas than a solid to a liquid.

Vaporization

• Not all of the particles in a liquid have the same kinetic energy.

• Temperature is a measurement of the average KE.

• The particles of a liquid must be moving at a sufficient speed to overcome the intermolecular forces of the liquid to escape as a gas.

• Thus, evaporation is endothermic: it requires energy to occur.

Vapor Pressure

• Vapor pressure is the pressure exerted by a vapor in equilibrium with its liquid phase at a certain temperature.

• For a liquid in a closed container, vaporization and condensation will occur at an equal rate once the equilibrium vapor pressure is reached.

Vapor Pressure and Boiling Point

• When water begins to boil, bubbles of H2O gas are created as some of the molecules escape the intermolecular bonds holding the liquid together.

• The bubbles will expand and rise to the surface only if the pressure exerted by the water vapor in the bubble is greater than the atmospheric pressure.

Vapor Pressure and Boiling Point

• Thus, the vapor pressure of the water must be equal to atmospheric pressure before boiling can occur.

• At temperatures below 100oC, the vapor pressure of water is below 1 atm, so the atmospheric pressure prevents boiling from occurring.

Elevation and Boiling Point• What happens to

atmospheric pressure as you increase your elevation above sea level?

• What would you expect to happen to the boiling point of water at high elevations?