Phases of Matter and Phase Changes

• This lesson explores phases of matter, transitions between them, and equilibrium in phase changes, covering vapor pressure, boiling/freezing points, and energy in phase transitions.

Phases and Phase Changes

• A phase is a system part with uniform composition and properties.
• The 3 common phases are solid, liquid and gas.
• A system can have multiple phases, like ice and water.
• Phase changes transition a substance from one phase to another.
• Sublimation is the transition from solid to gas.
• Vaporization is the transition from liquid to gas.
• Melting is the transition from solid to liquid.
• Freezing is the transition from liquid to solid.
• Deposition is the transition from gas to solid.
• Condensation is the transition from gas to liquid.
• Condensation is the change of a gas to a liquid, like water vapor turning into liquid water on a cold surface.
• Vapor is a gas in contact with its liquid or solid phase, like water vapor above liquid water.

Equilibrium

• Equilibrium occurs when opposing changes happen at equal rates in a closed system.
• In a closed container with liquid water, evaporation and condensation occur simultaneously.
• Equilibrium is dynamic; processes still occur, but there's no net change in each phase's amount.
• Vapor pressure equilibriums are affected by concentration, temperature, and pressure.

Vapor Pressure

• Vapor pressure is measured using a mercury manometer.
• Equilibrium vapor pressure is pressure exerted by a vapor in equilibrium with its liquid at a given temperature, and is specific for each liquid.
• Volatile liquids evaporate easily due to weak attractions.
• Nonvolatile liquids don't evaporate easily due to strong attractions.
• Acetone is volatile because it evaporates quickly; motor oil is nonvolatile.

Energy and Phase Changes

• Energy can be absorbed or released in a change, but is never destroyed or created based on the Law of Conservation of Energy.
• Vaporization changes a liquid to a gas and includes boiling (within and at the surface) and evaporation (only at the surface).
• Boiling is the change of a liquid to vapor within the liquid and at its surface.
• Lower atmospheric pressure results in a lower boiling point.
• Boiling point is when a liquid's equilibrium vapor pressure equals atmospheric pressure.
• Normal boiling point is the boiling point at standard atmospheric pressure (1 atm, 760 torr, or 101.3 kPa).
• Water's normal boiling point is exactly 100°C.
• Boiling and evaporation are endothermic, requiring added energy.
• Condensation is exothermic, requiring energy removal.
• Heat of Vaporization is the amount of heat energy needed to boil 1 gram of liquid at its boiling point.
• Water's heat of vaporization is 2260 J/g.
• The equation to calculate heat energy needed or removed during vaporization is q = mHvap, where q = energy (J), m = mass (g), and Hvap = heat of vaporization (J/g).
• Freezing is the change from liquid to solid, involving energy loss.
• Normal freezing point is when solid and liquid are in equilibrium at standard atmospheric pressure (1 atm, 760 torr, or 101.3 kPa).
• At equilibrium, melting and freezing occur at equal rates.
• Melting is endothermic, requiring energy addition.
• Freezing is exothermic, requiring energy removal.
• Heat of Fusion is the heat energy to melt 1 gram of solid at its melting point.
• Water's heat of fusion is 334 J/g.
• The equation to calculate heat energy needed or removed during fusion is q = mHfus, where q = energy (J), m = mass (g), and Hfus = heat of fusion (J/g).
• Under low temperature and pressure, a solid exists in equilibrium with its vapor.
• Sublimation is the change from solid directly to gas.
• Deposition is the reverse process, gas directly to solid.

Heating/Cooling Curves and Phase Diagrams

• Heating/cooling curves graphically represent phase transitions as heat is added or removed.
• Phase diagrams graph temperature versus pressure showing conditions for phases of a substance.
• Water's phase diagram shows at 1.00 atm, the n.b.p. is 100.00°C, and the n.f.p. is 0.00°C.
• The triple point is when all three phases are in equilibrium.
• The critical point is the critical pressure and temperature for the substance.
• Critical Temperature: The temperature above which the substance cannot exist as a liquid.
• Critical Pressure: The lowest pressure at which the substance, at critical temperature, can exist as a liquid.

Comprehensive Review

• A phase is a uniform part of a system (solid, liquid, gas).
• Equilibrium occurs when opposing changes happen at the same rate.
• Vapor pressure is influenced by temperature, concentration, and pressure.
• Volatile liquids evaporate easily; nonvolatile liquids do not.
• Energy changes accompany phase transitions:
• Endothermic: melting, vaporization, sublimation
• Exothermic: freezing, condensation, deposition
• Calculations use heat of fusion and heat of vaporization (q = mH).
• Phase diagrams show phase conditions with key points like the triple point and critical point.

Understanding the Kinetic Molecular Theory of Gases

• Explores the Kinetic Molecular Theory (KMT) of gases.
• KMT helps understand gas behavior based on the motion and properties of particles.

Temperature, Pressure, and KMT

• Gas containers often have a warning, "Do not store at a temperature above 120°F (50°C)."
• Temperature is directly proportional to the average kinetic energy of particles based on the KMT.
• Gas pressure is caused by collisions of gas particles with each other and the container walls.
• Increased temperature leads to particles moving faster, resulting in more frequent and forceful collisions, increasing pressure.
• Exceeding pressure can cause the container to rupture.

Five Postulates of the Kinetic Molecular Theory

• The volume of individual gas particles is nearly zero, making gases mostly empty space.
• There are no forces of attraction between gas particles due to great distance.
• Temperature of a gas in Kelvin is directly proportional to the average kinetic energy of particles.
• Lighter gas particles have greater speed than heavier ones at the same temperature.
• If temperature doubles, kinetic energy doubles.
• Gas pressure is caused by collisions of gas particles, with increased collisions leading to increased pressure.
• Collisions between gas particles, other particles, and the container wall are elastic collisions.
• Particles bounce off with the same speed that they hit with resulting in pressure when hitting the container.
• Particles are in constant, rapid, straight-line motion.
• Gases expand to fill their containers and are fluids that flow.

Ideal vs. Real Gases

• The Kinetic-Molecular Theory of Gases applies to ideal gases.
• Some gases nearly behave as an ideal gas.
• In reality, there is no ideal gas!
• Gases deviate the most from ideal gas when pressure is high, or when temperature is low.
• At high pressure, gas particles are closer, and intermolecular forces become more significant.
• At low temperatures, gas particles move slower, allowing intermolecular forces to have a greater effect.

Expansion and Fluidity

• Gases do not have a definite shape or a definite volume.
• Gases completely fill the container in which they are enclosed.
• Gases move rapidly in all directions without any significant attraction between particles.
• Gravity keeps all of the gases in our atmosphere from taking off into space.
• Forces between gas particles are so insignificant that they move about freely.
• Gases, as well as liquids, flow, and both are fluids.

Diffusion and Effusion

• Diffusion is the spontaneous mixture of gas particles of two or more substances caused by random motion.
• Effusion is the process of gas particles passing through a tiny opening.
• Gases spread out and mix with other gases spontaneously, even without stirring.
• Effusion of gases is directly related to the velocities of the particles.
• Lighter particles effuse faster than heavier particles.

Density and Compressibility

• Gases have low densities due to large spaces between particles.
• Gases are easily compressed because of these large spaces.
• Compressing a gas increases its pressure because it reduces the volume.

Kinetic-Molecular Theory of Matter

• The kinetic-molecular theory of matter is based on the idea that particles of matter are always in motion.
• GAS: High freedom of movement
• LIQUID: Less freedom than gases
• SOLID: Limited freedom of movement; particles are fixed in place

Comprehensive Review

• Temperature and Kinetic Energy: Temperature is directly proportional to the average kinetic energy of gas particles.
• Pressure and Collisions: Gas pressure results from collisions of gas particles with each other and the container walls.
• Five Postulates of KMT: Understanding the assumptions of the KMT is crucial for explaining gas behavior.
• Ideal vs. Real Gases: Real gases deviate from ideal behavior under high pressure and low temperature.
• Expansion and Fluidity: Gases expand to fill their containers and flow freely.
• Diffusion and Effusion: Gases mix spontaneously (diffusion) and pass through small openings (effusion).
• Density and Compressibility: Gases have low densities and are easily compressible.
• States of Matter: The kinetic-molecular theory explains the properties of solids, liquids, and gases.