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Chapter 12-13

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  1. Chapter 12-13 States of Matter

  2. States of Matter • Types of atoms (composition) and arrangement (structure) determine chemical properties of matter

  3. States of Matter • Composition and structure also affect the physical properties of matter

  4. Solids, Liquids and Gases • It is easy to see the difference between solids and liquids but gases are different

  5. Gases • Why are physical properties of gases different from solids and liquids?

  6. Kinetic Molecular Theory • A model that describes the behavior of gases in terms of particles in motion • The model makes several assumptions about the size motion and energy of gas particles

  7. Assumption 1: Size • Gas particles are small and separated from one another • The volume of the particle is small compared to the space in between • There are no significant attractions or repulsions among gas particles

  8. Assumption 2: Motion • Gas particles are in constant random motion • They move in a straight line until they run into something • Collisions are elastic=no kinetic energy is lost

  9. Assumption 3: Energy • Mass and velocity of particle determine kinetic energy • The bigger the mass the bigger the KE. KE=1/2 mv2 • Temperature is a measure of KE

  10. Explanation of Gas Behavior • Diffusion is the movement of one material through another. • Because of the space between particles gases diffuse easily

  11. Effusion: Graham’s Law • Similar to diffusion, effusionis the rate that a gas moves through a tiny opening. • Graham determined that the rate that the gas effused was related to its molar mass Rate of effusion is proportional to 1 molar mass • To compare the diffusion of two gases • Ratea = molar massb • Rateb molar massa

  12. Explanation of Gas BehaviorKMT explains • Density is mass per volume • Gases occupy large volumes with very little mass

  13. Explanation of Gas Behavior • Compression and Expansion the empty space found between gas particles allows for compression and expansion

  14. Gas Pressure • Pressure is defined as the force per unit area. The earth is surrounded by the atmosphere which exerts pressure on the surface of the earth. • Units of atmospheric pressure are 1 atm = 101.3kPa = 760 mm Hg = 760 torr

  15. Measuring Pressure • Barometer Measures the force exerted on the surface of a pool of mercury

  16. Measuring Pressure Manometers measure the pressure in a vessel relative to the atmosphere

  17. Converting Pressure Units 1 atm = 101.3kPa = 101325 Pa = 760 mm Hg = 760 torr The pressure of N2 gas was measured to be 750 mm Hg. What is the pressure of the gas in kPa? 750 mm Hg101.3kPa = 760 mm Hg

  18. Daltons Law of Particle Pressure • The pressure total is equal to the sum of the individual pressure that make up the gas mixture • Ptotal = P1 + P2 +P3 + … Pn

  19. Daltons Law of Particle Pressure • A mixture of O2, CO2, and N2 has a total pressure of 0.97 atm. What is the partial pressure of O2 if CO2 has a pressure of 0.70 atm and N2 has a pressure of 0.12 atm? 0.97 atm = 0.70atm + 0.12atm + xatm x= .15atm

  20. Forces of Attraction • Intramolecular forces: Occur between atoms to form molecules. Examples include Ionic, Covalent and metallic bonds • Intermolecular forces: Occur between molecules. Examples include dispersion, dipole-dipole and hydrogen bonding.

  21. Forming a dipole • A dipole is a charged region of a molecule. (some regions of the molecule have a partially positive or negative charge.) This occurs when there is a shift in the electron density d- O H H d+ d+ Dipole Neutral

  22. Intermolecular Forces Intermolecular Forces • Dispersion Forces Involve temporary dipole-dipole interactions d- d+ d- d+

  23. Intermolecular Forces • Dipole- Dipole forces involve molecules with permanent dipoles. An example would be HCl:

  24. Intermolecular Forces • Hydrogen Bonding is a dipole-dipole force that involves H bonded to nitrogen, oxygen or a halogen. • Hydrogen bonding is very strong. An example of this is water

  25. Effects of Intermolecular Forces • Solids and Liquid tend to have strong molecular forces. • Water is composed of 2 gases. When combined to make a molecule with strong molecular forces water exists at room temperature as a liquid. • Phase changes involve forming or breaking intermolecular forces.

  26. The Liquid State • Intermolecular forces play a part in: • Surface Tension • Capillary Action • Viscosity

  27. Surface Tension • All particles are not attracted equally in a solution. • Particles in the middle are attracted by those above and below while those on the surface are attracted to those below.

  28. Surface Tension • So the surface stretches over the top • ST is a measure of the inward pull by particles in the interior

  29. Capillary Action • Capillary action is a result of the different degrees of attraction between the container and the liquid.

  30. Capillary Action • If the attraction to the container is greater than the attraction to other water particle the water will travel up the capillary

  31. Viscosity • The resistance of a liquid to flow. • Particles in a liquid are close enough to each other to have intermolecular forces involved. • The stronger the forces the higher the viscosity. • The size and shape of the molecules and the temperature effect viscosity.

  32. Viscosity • Viscosity of motor oil increases in the summer.

  33. Surface Tension Activity • Obtain a 250 or 400mL beaker with water in it • Float the paper clip on the surface of the water • Use a dropper to add one drop of water containing detergent to the beaker. Observe what happens

  34. Answer these questions for Homework • Is a paper clip likely to be more or less dense than water? • How does the shape of the pin help it float • Hypothesize about the reason for the pin’s behavior before and after you added the detergent.

  35. Change of State/Phase Changes • NO bonds are broken when a compound changes state.

  36. Phase Changes That REQUIRE energy Melting, Evaporation, Sublimation and Boiling

  37. Vaporization • Vaporization is the process by which gas particles escape the surface of a solid or a liquid. This process can occur through: • Evaporation • Sublimation • Boiling

  38. Melting • The energy absorbed by the ice is used to break hydrogen bonds that held the ice together. The temp at which this happens is the melting point

  39. Evaporation • Requires energy to change from a liquid to a gas • This is how we cool ourselves

  40. Sublimation

  41. Boiling

  42. Vapor Pressure • Vapor pressure is the pressure of the gas over a liquid in equilibrium. • The rate of evaporation = rate of condensation

  43. Vapor Pressure • Vapor pressure is determined principally by the size of the intermolecular forces in the liquid. • Vapor pressure increases significantly with temperature.

  44. Boiling • If the temperature of a liquid increases, the molecules of water gain kinetic energy and the vapor pressure increases.

  45. Boiling • When the vapor pressure of the liquid = the pressure of the surrounding atmosphere then the boiling will occur.

  46. Boiling At high altitudes, the boiling point of liquids is lower than at sea level. In Denver, Colorado, water will boil at about 94°C. Do not confuse boiling with cooking. Cooking pasta in Denver is a slower process because the water is at a lower temperature. Also, realize that water boiling rapidly is no hotter than water boiling slowly.

  47. Boiling • The temperature of a boiling liquid never rises above its boiling point. No matter how much heat is applied, the liquid only boils faster, not hotter.

  48. Boiling • Compounds with a high degree of intermolecular forces will have high boiling points.

  49. Phase Changes that RELEASE energy Condensation, Depositon, Freezing