1 / 67

Exploring Molecular Structures in States of Matter

Delve into the world of states of matter, from solids to liquids to gases, and understand the properties of crystals, vapor pressure, evaporation, and boiling points through the lens of kinetic theory. Learn about crystal structures, amorphous solids, evaporation processes, vapor pressure measurements, and the factors affecting boiling points.

jameslarsen
Download Presentation

Exploring Molecular Structures in States of Matter

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chapter 13 States of Matter

  2. Kinetic Theory • Atoms are always in motion and are colliding • Atoms move faster as temperature increases • Heavier particles move slower than smaller less massive particles

  3. States of Matter • Solid • Liquid • Gas

  4. States of Matter • http://web.jjay.cuny.edu/~acarpi/NSC/2-matter.htm

  5. Solids • Have definite shape and volume • Particles are packed close/tight together in a form held together by molecular forces. • Divided into two categories • Crystalline – orderly arrangement • Iron, ice, and diamond • Amorphous – no order • Rubber and wax

  6. 13.3 A Model for Solids • The melting point (mp) is the temperature at which a solid changes into a liquid.

  7. 13.3 Crystal Structure and Unit Cells • In a crystal, the particles are arranged in an orderly, repeating, three-dimensional pattern called a crystal lattice.

  8. 13.3 Crystal Structure and Unit Cells • The shape of a crystal reflects the arrangement of the particles within the solid.

  9. 13.3 Crystal Structure and Unit Cells • Crystal Systems • A crystal has sides, or faces. Crystals are classified into seven crystal systems.

  10. 13.3 Crystal Structure and Unit Cells • These minerals show four out of the seven crystal systems.

  11. 13.3 Crystal Structure and Unit Cells • The smallest group of particles within a crystal that retains the geometric shape of the crystal is known as a unit cell. • A crystal lattice is a repeating array of any one of fourteen kinds of unit cells. • There are from one to four types of unit cells that can be associated with each crystal system.

  12. 13.3 Crystal Structure and Unit Cells • Three kinds of unit cells can make up a cubic crystal system.

  13. 13.3 Crystal Structure and Unit Cells • Allotropes • Allotropes are two or more different molecular forms of the same element in the same physical state. • Allotropes have different properties because their structures are different. • Only a few elements have allotropes.

  14. 13.3 Crystal Structure and Unit Cells • Carbon Allotropes

  15. 13.3 Crystal Structure and Unit Cells • Non-Crystalline Solids • An amorphous solid lacks an ordered internal structure. • Rubber, plastic, asphalt, and glass are amorphous solids. • A glass is a transparent fusion product of inorganic substances that have cooled to a rigid state without crystallizing.

  16. Liquid 13.2 • Has a definite volume but no shape • Take shape of the container they occupy. • Particles are tightly packed together, but are able to slide pass each other • Particles move fast enough to overcome molecular forces

  17. 13.2 A Model for Liquids • Substances that can flow are referred to as fluids. Both liquids and gases are fluids.

  18. 13.2 Evaporation • Evaporation • What is the relationship between evaporation and kinetic energy?

  19. 13.2 Evaporation • The conversion of a liquid to a gas or vapor is called vaporization. • When such a conversion occurs at the surface of a liquid that is not boiling, the process is called evaporation.

  20. 13.2 Evaporation • In an open container, molecules that evaporate can escape from the container.

  21. 13.2 Evaporation • In a closed container, the molecules cannot escape. They collect as a vapor above the liquid. Some molecules condense back into a liquid.

  22. 13.2 Evaporation • During evaporation, only those molecules with a certain highest kinetic energy can escape from the surface of the liquid.

  23. 13.2 Vapor Pressure • Vapor Pressure • When can a dynamic equilibrium exist between a liquid and its vapor?

  24. 13.2 Vapor Pressure • Vapor pressure is a measure of the force exerted by a gas above a liquid.

  25. 13.2 Vapor Pressure • A system is in equilibrium when the rate of evaporation of liquid equals the rate of condensation of vapor.

  26. 13.2 Vapor Pressure • Vapor Pressure and Temperature Change • An increase in the temperature of a contained liquid increases the vapor pressure. • The particles in the warmed liquid have increased kinetic energy. As a result, more of the particles will have the minimum kinetic energy necessary to escape the surface of the liquid.

  27. 13.2 Vapor Pressure

  28. 13.2 Vapor Pressure • Vapor Pressure Measurements • The vapor pressure of a liquid can be determined with a device called a manometer.

  29. 13.2 Vapor Pressure • Manometer

  30. 13.2 Boiling Point • Boiling Point • Under what conditions does boiling occur?

  31. 13.2 Boiling Point • When a liquid is heated to a temperature at which particles throughout the liquid have enough kinetic energy to vaporize, the liquid begins to boil.

  32. 13.2 Boiling Point • The temperature at which the vapor pressure of the liquid is just equal to the external pressure on the liquid is the boiling point (bp).

  33. 13.2 Boiling Point • Boiling Point and Pressure Changes • Liquids don’t always boil at the same temperature. • At a lower external pressure, the boiling point decreases. • At a higher external pressure, the boiling point increases.

  34. 13.2 Boiling Point • Altitude and Boiling Point

  35. 13.2 Boiling Point

  36. 13.2 Boiling Point • Normal Boiling Point • The normal boiling point is defined as the boiling point of a liquid at a pressure of 101.3 kPa.

  37. 13.2 Boiling Point

  38. Section 1 • Gases

  39. Gas • Are free to move in all directions • Gases expand to fill a space. • Particles move fast enough to break away from each other • Particles move rapidly

  40. Properties of Gases • Gases have no definite shape or volume • Particles move rapidly in all directions • http://web.jjay.cuny.edu/~acarpi/NSC/2-matter.htm • Gases are fluids • Gases have low density, b/c particles are far apart • Gases are compressible • Gases spread out easily and mix w/one another.

  41. Gases Exert Pressure • Since particles are always moving, they continuously bump each other and the walls of the container. • Balloon • Propane container

  42. Gas Laws • Gas laws – describe how the behavior of gases is affected by pressure and temperature. • Boyle’s Law • Charles’s Law • Gay-Lussac’s Law

  43. 13.1 Kinetic Theory and a Model for Gases • Particles in a gas are in rapid, constant motion.

  44. 13.1 Kinetic Theory and a Model for Gases • Gas particles travel in straight-line paths.

  45. 13.1 Kinetic Theory and a Model for Gases • The gas fills the container.

  46. 13.1 Gas Pressure • Gas Pressure • How does kinetic theory explain gas pressure?

  47. 13.1 Gas Pressure • Gas pressure results from the force exerted by a gas per unit surface area of an object. • An empty space with no particles and no pressure is called a vacuum. • Atmospheric pressure results from the collisions of atoms and molecules in air with objects.

  48. 13.1 Gas Pressure • Gas pressure is the result of simultaneous collisions of billions of rapidly moving particles in a gas with an object.

  49. 13.1 Gas Pressure • A barometer is a device that is used to measure atmospheric pressure.

  50. Section 4

More Related