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States of Matter

States of Matter

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States of Matter

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  1. States of Matter • There are three main states of matter Kinetic molecular Theory (KMT) helps us explain the states of matter.

  2. Kinetic Molecular Theory (KMT) KMT states that particles of matter are in constant motion. Molecular motion is defined by: Kinetic Energy = ½ mv2 Kinetic energy is a reflection of the substances temperature. (increase temp = increase kinetic energy) Some assumptions about KMT as they apply to ideal gases: Gas particles are spaced far apart Gas particles collide with “elastic” collisions Particles are in constant, rapid and random motion Gases are not attracted to each other. The temperature of a gas reflects the kinetic energy of the gas particles.

  3. States of matter • Properties of Gases: • Expansion- gases have no definite shape or volume. • Fluidity- gas particles will easily glide past one another • Low density- gases have 1/1000 the density of other substances • Compressible- gas particles can be pushed closer together • Diffusion- gases can easily disperse and mix in space. • Effusion- gases can pass through a tiny opening.

  4. Deviation from ideal • Comparison of gases:

  5. Effusion and Diffusion: • Effusion- gas particles escape through a small opening. example of effusion • Diffusion-ability of molecule to mix.

  6. States of Matter: • Properties of Liquids: • Liquids are more dense than gases due to attraction of particles. • Liquids are much less compressible than gases, less empty space. • Liquids have fluidity- able to glide past each other and diffuse. • Liquids can boil and evaporate. • Liquids can be formed into solids.

  7. States of Matter • Properties of Solids: • Definite shape and volume-organized arrangement of particles • Definite melting point- energy requirement to overcome forces holding the particles together • Greatest density/ incompressible- little/no empty space • Low rate of diffusion- fixed position of particles will not allow them to glide and mix. • Organized/crystalline arrangement Simulation and states of matter.

  8. Intermolecular Forces (IMF) • Intermolecular forces- describe the interaction between molecules and influence many properties including: • Melting point (solid ↔ liquid) • Boiling point (liquid ↔ gas) • Vapor pressure (gas phase equilibrium of a liquid)

  9. IMF & Properties of Water • Water molecules are polar and have strong intermolecular attractions. • Unique properties of water due to IMF: • High surface tension • Capillary action • High boiling point (considering MWt.) • Density of solid < Liquid state

  10. Surface Tension  tendency of a liquid to contract in area  result of molecular attractions  a sphere has smallest surface area for a given volume  water has very high surface tension  explains why needles can float on water

  11. Capillary Action  rising of a liquid in a narrow tube  caused by cohesion and adhesion  adhesion - attractive force between two unlike substances  cohesion - attractive force between like substances

  12. States of Matter and IMF • IMF defines the energy (temp) requirements for a material to change phase phases. • Some of these relationships can be interpreted by understanding the following graphs: • Vapor Pressure Curves • Phase Diagrams

  13. Vapor Pressure: • Equilibrium vapor pressure is the pressure exerted by a vapor at equilibrium with its corresponding liquid at a given temperature (as gas molecules enter the headspace, the Vapor pressure increased) • General trend : IMF =  Vapor pressure

  14. What is Pressure: • Is defined as force/area. • SI unit = Pascal's • Atmospheric Pressure- collision of air molecules with earth surfaces.  attitude =  atmospheric pressure • Barometer- used to measure atmospheric pressure in mm Hg. Useful relationships: 1 atm = 760 mmHg = 760 torr= 101.3 kPa

  15. Vapor Pressure Curve: • Describes the boiling point (temp) of a liquid at a given atmospheric pressure. • Vapor pressure increases with increasing temperature. • IMF = Vapor pressure and IMF = Vapor pressure

  16. Boiling • A liquid boils when the vapor pressure = the external pressure • Normal Boiling point is the temperature a substance boils at 1 atm pressure. • The temperature of a liquid can never rise above it’s boiling point.

  17. Changing the Boiling Point – lowering it • Lower the pressure (going up into the mountains). • Lower external pressure requires lower vapor pressure. • Lower vapor pressure means lower boiling point. • Food cooks slower.

  18. Changing the Boiling Point – raising it • Raise the external pressure (Use a pressure cooker). • Raises the vapor pressure needed. • Raises the boiling point. • Food cooks faster.

  19. Effect of Pressure on Boiling Point

  20. Melting Vaporization Freezing Condensation Phase Changes Solid Gas Liquid

  21. Sublimation Vaporization deposition Melting Solid Gas Liquid Freezing Condensation

  22. Phase Diagram • Represents phases as a function of temperature and pressure. • Critical temperature: temperature above which the vapor can not be liquefied. • Critical pressure: pressure required to liquefy AT the critical temperature. • Critical point: critical temperature and pressure (for water, Tc = 374°C and 218 atm).

  23. Phase changes by Name

  24. Water