Condensed States of Matter: Liquids and Solids Chapter 14

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

only slightly compressible Gas Liquid Solid highly compressible only slightly compressible low density high density fills container completely does not expand to fill container - has definite volume rigidly retains its volume 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

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 Dipole-dipole attractions a. Hydrogen bonding Ion-dipole attractions 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

Intermolecular Attraction Kinetic Energy Intermolecular Attraction State at Room Temp. low high solids medium liquids gases Kinetic 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: Crystalline Atomic Ionic Covalent-network 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

Liquids The 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 g liquid g gas Energy out: gas g liquid g 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?