Solids, Liquids and Gases Solid Fixed Volume Rigid, definite shape Liquid Fixed Volume Indefinite Shape Particles easily glide past one another Gas Particles farther apart than in liquids No definite volume or shape Flowing and Compressible
The Kinetic Theory of Matter 1827: Robert Brown Scottish Botanist Studied pollen grains on water Brownian Motion: Constant, random motion of tiny chunks of matter Kinetic Energy: The energy of moving objects
Kinetic Model of Gases Particles move in a straight line until they strike a container wall or another gas particle. Particles do NOT lose speed when collisions are made (ELASTIC COLLISIONS) Gases fill container because each particle moves until it hits a wall and then changes direction
Kinetic Model of Liquids Particles maintain volume, but not shape Particles slide past one another, but they do not move as straight or as quickly as particles of gases.
Properties of Liquids Surface tension: the energy required to increase the surface area of a liquid by a unit amount. Viscosity: a measure of a liquid’s resistance to flow. Surface tension: The net pull toward the interior of the liquid makes the surface tend to as small a surface area as possible and a substance does not penetrate it easily. Viscosity: Related to mobility of a molecule (proportional to the size and types of interactions in the liquid). –Viscosity decreases as the temperature increases since increased temperatures tend to cause increased mobility of the molecule.
Intermolecular Forces Intermolecular forces: attractions and repulsions between molecules that hold them together. Intermolecular forces (van der Waals forces) hold molecules together in liquid and solid phases. Ion-dipole force: interaction between an ion and partial charges in a polar molecule. Dipole-dipole force: attractive force between polar molecules with positive end of one molecule is aligned with negative side of other. London dispersion Forces: interactions between instantaneously formed electric dipoles on neighboring polar or nonpolar molecules.
Kinetic Model of Solids Strong force between particles = rigid structure Particles do not slide past each other, but instead vibrate, or bounce back and forth between each other. Form a CRYSTAL LATTICE, or repeating, fixed, 3-D arrangement
Structure of Solids Types of solids: Crystalline – a well defined arrangement of atoms; this arrangement is often seen on a macroscopic level. Ionic solids – ionic bonds hold the solids in a regular three dimensional arrangement. Molecular solid – solids like ice that are held together by intermolecular forces. Covalent network – a solid consists of atoms held together in large networks or chains by covalent networks. Metallic – similar to covalent network except with metals. Provides high conductivity. Amorphous – atoms are randomly arranged. No order exists in the solid.
Changing States Evaporation: particles of a liquid form a gas by escaping from the surface VOLATILE liquids evaporate quickly Vapor Pressure: Pressure at liquid/gas equilibrium Boiling: Vapor Pressure is the pressure at the surface
Changing States Condensation: Gaseous particles come closer together and form a liquid Sublimation: particles of a solid escape from the surface and form a gas Deposition: particles of a gas are forced together under pressure and form a solid
Changing States Melting: Particles of a solid begin to lose their crystal lattice and slide past one another. Freezing: particles form crystal lattice and form a rigid structure
Changing States Heat of Vaporization Water vapor and Boiling water are at the same temperature Gaseous particles form because kinetic energy increases Joule: (J) SI unit of energy; energy required to lift 1 kg mass by 1 meter against gravity Energy absorbed when 1 kg of a liquid vaporizes at normal boiling point.
Changing States Heat of Fusion Melting Point: Temperature at which crystal lattice begins to disintegrate. Freezing Point = Melting Point Energy released as 1kg of a substance solidifies at its freezing point.
Heat of fusion for H 2 O: 80.0 cal/g A calorie is the amount of heat needed to raise the temperature of one gram of water 1 0 C 1.Convert heat of fusion to Joules (J). 1.00 cal = 4.18J 4.18J/1 cal X 80.0 cal = 334J 2. What is the molar heat of fusion for water ( in cal & in J) a) 80.0 cal/g X 18.00g= 1440cal or 1.44 kcal b) 334J X 18.00g= 6010J or 6.01kJ
Sample Problems Heat of Vaporization of water: 540. cal/g 1.Convert heat of vaporization to J & kJ 540.cal/g X 4.18J/cal = 2260J/g or 2.26kJ/g 2. What is the molar heat of vaporization for water? (cal & J) a)540. cal/g x 18.00g/mol = 9720 cal or 9.72kcal b) 2260J/g x 18.00g/mol = 40.7 kJ
Change of State Problems If you wanted to melt 16.00g ice at – 25.0 0 C what would you have to do first? Change the temperature to the melting temp: zero 0 Q = m ∆t Cp Q= heat M= mass Cp= Specific Heat Capacity
Specific Heat Capacity = the amount of heat needed to raise the temperature of 1.00g of a substance 1.00 0 C. Cp of water= 1.00 cal/g.K Cp of ice= 0.493 cal/g.K Cp of steam= 0.447 cal/g.K 1.To melt 16.00g of ice at -25.0 its temperature must be raised to zero C: Q = m ∆t Cp Q = (16.00g)(25.0)(0.493) Q= 197cal
2. How much energy is needed to melt 68.0 g of ice at 0 0 C into water at 0 0 C ? 3. How much energy is needed to change 42.0g of water at 100.0 0 C into steam at 100.0 0 C? 4. How much energy is needed to raise the temperature of 57.0g of water at 25.0 0 C to its boiling point? 5. How much energy is needed to convert 15.0g of ice at -6.00 0 C to water at 25.0 0 C?
The triple point of a substance indicates the temperature and pressure at which a solid, liquid and vapor can coexist at equilibrium. The critical point of a substance indicates critical temperature and critical pressure. Critical temperature is the temperature above which only the vapor can exist. Critical pressure is the lowest pressure (at the critical temperature) at which the substance can exist as a liquid. At any lower pressure only the vapor exists.