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Chapter 9- The States of Matter

Chapter 9- The States of Matter. Gases indefinite volume and shape, low density. Liquids definite volume, indefinite shape, and high density. Solids definite volume and shape, high density Solids and liquids have high densities because their molecules are close together. Kinetic Theory. l l

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Chapter 9- The States of Matter

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  1. Chapter 9- The States of Matter • Gases indefinite volume and shape, low density. • Liquids definite volume, indefinite shape, and high density. • Solids definite volume and shape, high density • Solids and liquids have high densities because their molecules are close together.

  2. Kinetic Theory l l are evidence of this. • Kinetic theory says that molecules are in constant motion. • Perfume molecules moving across the room

  3. The Kinetic Theory of GasesMakes three assumptions about gases • A Gas is composed of particles • usually molecules or atoms • Considered to be hard spheres far enough apart that we can ignore their volume. • Between the molecules is empty space.

  4. The particles are in constant random motion. • Move in straight lines until they bounce off each other or the walls. • All collisions are perfectly elastic

  5. Kinetic Energy and Temperature • Temperature is a measure of the Average kinetic energy of the molecules of a substance. • Higher temperature faster molecules. • At absolute zero (0 K) all molecular motion would stop.

  6. High temp. % of Molecules Low temp. • Kinetic Energy

  7. High temp. Low temp. % of Molecules Few molecules have very high kinetic energy • Kinetic Energy

  8. High temp. % of Molecules Low temp. Average kinetic energies are temperatures • Kinetic Energy

  9. Temperature • The average kinetic energy is directly proportional to the temperature in Kelvin • If you double the temperature (in Kelvin) you double the average kinetic energy. • If you change the temperature from 300 K to 600 K the kinetic energy doubles.

  10. Temperature • If you change the temperature from 300ºC to 600ºC the Kinetic energy doesn’t double. • 873 K is not twice 573 K

  11. Liquids • Particles are in motion. • Attractive forces between molecules keep them close together. • These are called intermolecular forces. • Inter = between • Molecular = molecules

  12. Breaking intermolecular forces. • Vaporization - the change from a liquid to a gas below its boiling point. • Evaporation - vaporization of an uncontained liquid ( no lid on the bottle ).

  13. Evaporation • Molecules at the surface break away and become gas. • Only those with enough KE escape • Evaporation is a cooling process. • It requires heat. • Endothermic.

  14. Condensation • Change from gas to liquid • Achieves a dynamic equilibrium with vaporization in a closed system. • What is a closed system? • A closed system means matter can’t go in or out. (put a cork in it) • What is • “dynamic equilibrium?”

  15. Vaporization • Vaporization is an endothermic process - it requires heat. • Energy is required to overcome intermolecular forces • Why we sweat.

  16. Energy needed to overcome intermolecular forces % of Molecules T1 Kinetic energy

  17. % of Molecules • At higher temperature more molecules have enough energy • Higher vapor pressure. T2 Kinetic energy

  18. 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.

  19. Solids • Intermolecular forces are strong • Can only vibrate and revolve in place. • Particles are locked in place - don’t flow. • Melting point is the temperature where a solid turns into a liquid.

  20. The melting point is the same as the freezing point. • When heated the particles vibrate more rapidly until they shake themselves free of each other. • Ionic solids have strong intermolecular forces so a high mp. • Molecular solids have weak intermolecular forces so a low mp.

  21. Energy and Phase Change • Heat of vaporization energy required to change one gram of a substance from liquid to gas. • Heat of condensation energy released when one gram of a substance changes from gas to liquid.

  22. Energy and Phase Change • Heat of fusion energy required to change one gram of a substance from solid to liquid. • Heat of solidification energy released when one gram of a substance changes from liquid to solid.

  23. Heating Curve for Water 120 Steam Water and Steam 100 80 60 Water 40 20 Water and Ice 0 Ice -20 40 120 0 220 760 800 Heat of Vaporization

  24. Heating Curve for Water 120 Steam Water and Steam 100 80 60 Water 40 20 Water and Ice 0 Ice -20 40 120 0 220 760 800 Heat of Fusion

  25. Heating Curve for Water 120 Steam Water and Steam 100 80 60 Water 40 20 Water and Ice 0 Ice -20 40 120 0 220 760 800 Steam Water Slope = Specific Heat Ice

  26. Heating Curve for Water 120 Steam Water and Steam 100 80 60 Water 40 20 Water and Ice 0 Ice -20 40 120 0 220 760 800 Both Water and Steam

  27. Heating Curve for Water 120 Steam Water and Steam 100 80 60 Water 40 20 Water and Ice 0 Ice -20 40 120 0 220 760 800 Ice and Water

  28. Calcualting Energy • Three equations • Heat = specific heat x mass x DT • Heat = heat of fusion x mass • Heat = heat of vaporization x mass

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