Kinetic Molecular Theory

Kinetic theory Describes how particles move as an ideal gas Can be adjusted to describe the motion of particles in all states of matter Points of the Kinetic Theory Gases consist of a large number of tiny particles that are far apart relative to their size Collisions between particles and between particles and the container wall are elastic Gas particles are in continuous, rapid, random motion There are no forces of attraction between gas particles The temperature of a gas depends on the average kinetic energy of the particles in the gas KE = 1/2mv2

Behavior of Gases The behavior of a gas is directly related to the five points of the theory Gases Expand (expansion) This is a result of no attractive forces and a constant, random motion Gases are fluids (fluidity) Particles can freely move past one another No attractive forces, tiny particles that are far apart, and constant motion Gases have a low density A lot of empty space with a relatively low mass Particles are far apart relative to their size

Behavior Continued Gases can be compressed Gases are far apart with empty space therefore they can be forced closer Gases have the ability to diffuse and effuse Due to the constant, random motion of gases particles, they will fill the volume of their container and no attractive forces they will mix together This allows them to mix evenly (diffusion) and escape through tiny openings in their containers (effusion)

Real Gases Real gases tend to deviate from ideal gases in that we do see some slightly attractive forces between particles Gases such as the noble gases will exhibit a more ideal nature than other gases due to their unreactive nature The temperature and pressure can also affect the ideal nature High temperatures and low pressures will typically exhibit an ideal nature

Liquids As temperature and pressure begin to change, the principles of the kinetic theory begin to breakdown Principle #4: There are no attractive forces between particles As kinetic energy decreases, attraction between particles will increase Particles in the liquid state lose kinetic energy and become attracted to each other which forces the particles to stay close together (fixed volume)

Behavior of Liquids Liquids are fluids Liquids have a high density Even though they lose kinetic energy, particles still have enough kinetic energy to rotate freely around each other Liquids have a high density With stronger attractive forces, particles are more tightly packet occupying a smaller volume Less empty space between particles Liquids are relatively incompressible Eliminate empty space Liquids can diffuse Particles can still move freely around one another so they can freely mix with other liquids

Behaviors Continued Liquids have surface tension Due to the attractive forces, liquid particles want to stay together so there is a force that is required to break them apart There is also an attraction between the liquid particles and solid particles known as capillary action Liquids have the ability to evaporate and boil Both boiling and evaporation are part of a process known as vaporization (liquid becoming a gas) Evaporation occurs only at the surface of the liquid Boiling can occur throughout the liquid Liquids can become solids by freezing

Solids With a continued loss of kinetic energy, the attractive forces between particles will continue to grow stronger pulling the particles closer together Eventually, the particles will be pulled into a fixed position and exhibit a repeating pattern Particles will just be vibrating in place

Behavior of Solids Solids have two types: crystalline and amorphous Crystalline solids consist of crystals which have an orderly, geometric, repeating pattern Amorphous solids have a random arrangement of particles Solids have a definite shape and volume Solids have a definite melting point Defined temperature required to break apart the attractive forces and allow the particles to move freely Solids have a high density and are incompressible Solids have a very low rate of diffusion The rate is millions of times slower in solids than in liquids

Crystalline Solids Crystal structure Three dimensional arrangement of particles This arrangement can be represented by a coordinate system called a lattice A unit cell is the smallest representation of this repeating pattern

Forces of attraction Crystals can be classified into four categories Ionic Crystals: repeating pattern of oppositely charged ions Covalent Network Crystals: each atom is bonded to the nearest atom by sharing electrons Metallic Crystals: Repeating pattern of metal cations that are surrounded by a sea of electrons Covalent Molecular crystals: Molecules are formed by covalently bonded atoms and then the molecules are held together by intermolecular attraction

Amorphous Solids Comes from the Greek word meaning “without shape” Glasses and Plastics (polymers) are good examples These solids are formed in a manner that does not allow them to crystallize whether it be by a cooling process (glasses) or molded under high temperature and pressure (plastics)

Changes of State Possible changes in state Solid to liquid (melting) Solid to gas (sublimation) Liquid to solid (freezing) Liquid to gas (vaporization) Gas to liquid (condensation) Gas to solid (deposition)

Equilibrium This is a dynamic condition in which two opposing changes occur at equal rates in a closed system Equilibrium can be establish between any two states that are changing at the temperature where that change occurs For example, a sealed liquid will continually build up vapor pressure at the surface of the liquid as it evaporates This process will continue until equilibrium vapor pressure was achieved Keeping in mind that equilibrium is a dynamic condition Temperature and pressure can change the equilibrium points

Boiling Boiling occurs when the partial pressure of the liquid is equal to the atmospheric pressure Boiling point is the temperature where boiling will occur Normal boiling point is the temperature at 1 atm of pressure (or 760 torr or 101.3 kPa) Once a substance has started to boil, all the energy that is added goes directly to converting the state The amount of energy needed to vaporize one mole of liquid at the boiling point is known as the molar enthalpy (heat) of vaporization Expressed as kJ/mol (kilojoules per mole)

Freezing and Melting Freezing point is the temperature at which a solid and liquid are in equilibrium Normal freezing point is also at 1atm Melting point is the same temperature just depends on whether energy is added or removed from our system Just like with boiling, energy will be added (melting) or removed (freezing) until the substance is entirely converted but the temperature will not change Water at 0ºC will remain at 0ºC until all the ice is melted or the water is frozen The amount of energy required to melt one mole of a solid at the melting point is the molar enthalpy (heat) of fusion Expressed as kJ/mol

Phase Diagram

Phase Diagram Triple Point Critical Point Critical temperature Temperature and pressure where all three state exist in equilibrium Critical Point Indicates where the critical temperature and critical pressure meet Critical temperature Temperature in which you can no longer have a substance in the liquid state Critical pressure Lowest pressure where a substance can still be a liquid at the critical temperature