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Chapter 11. States of Matter

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  1. Chapter 11.States of Matter

  2. States of Matter State is Determined by: Chemical Identity Temperature Pressure

  3. States of Elements

  4. Kinetic Molecular Theory of Matter The Kinetic Molecular Theory of Matter is an explanation of the behavior of matter, based on the idea that the particles that make up the matter are always in motion. The particles can be atoms, molecules, or ions.

  5. Kinetic Molecular Theory of Matter Matter is composed of tiny particles; the size of the particles is fixed for each sub-stance. 2. The particles are in constant random mo-tion and therefore possess kinetic energy (energy of motion, which can be trans-ferred by collisions).

  6. Kinetic Molecular Theory of Matter The particles interact with each other by means of electrostatic attractions and repulsions, and therefore possess potential energy (stored energy, pos-sessed by matter as a result of its po-sition, condition, and or composition).

  7. Kinetic Molecular Theory of Matter The kinetic energy (velocity) of the particles increases as temperature is increased. 5. The particles in a system transfer energy by means of elastic collisions (collisions in which all energy transfer results in motion, not deformation).

  8. Kinetic Molecular Theory of Matter No kinetic energy is lost in elastic collisions. Kinetic energy is transformed to work and/or heat in inelastic collisions. Kinetic energy is a disruptive force between particles; it makes them more independent of each other. Potential energy is a cohesive or attractive force between particles.

  9. Comparison of States

  10. Some Definitions: Density is the ratio of mass to volume. Compressibility is a measure of volume change resulting from a pressure change. Thermal Expansion is a measure of volume change resulting from temperature change.

  11. Properties of States Solids Definite Volume, Definite Shape Density is High: 1.0 – 20 g/cm3 Compressibility is Low Thermal Expansion is Low: 0.01% per C

  12. The Solid State In solids, cohesive forces predominate over kinetic energy. Particles are usually held together in a regular array, and vibrate about fixed positions. Electrostatic attractions between particles keep them close together in fixed positions. The particles fill 50-70% of the space available, the rest is “void volume.”

  13. Properties of States Liquids Definite Volume, Indefinite Shape Density is Fairly High: 0.5 – 15 g/mL Compressibility is Low Thermal Expansion is Fairly Low: 0.1% per C

  14. The Liquid State In liquids, cohesive forces are balanced by kinetic energy. Particles move freely about each other but do not separate. Electrostatic attractions between particles keep them close together, but able to move rel-ative to one another. The particles fill about 50% of the available space.

  15. Properties of States Gases Indefinite Volume, Indefinite Shape Density is Low: 0.2 – 10 g/L Compressibility is High Thermal Expansion is Moderate: 0.3% per C

  16. The Gaseous State In gases, cohesive forces are overcome by kinetic energy. Particles move indepen-dently of each other. Electrostatic attractions between particles are very weak, do not hold them together. The particles fill about 1% of the space available.

  17. Changes in State Most substances can exist in any phase, solid, liquid, or gas. The phase at which the substance exists depends on its temperature and the applied pressure. A phase diagram is a graph showing the phase behavior of a given substance.

  18. Phase Diagram for Water

  19. Phase Diagram for CO2

  20. Changes of State Exothermic: Endothermic Release Heat Energy Absorb Heat Energy H is negative H is positive

  21. Changes of State Freezing pointthe temperature at which the liquid and solid phases of a substance are in equilibrium Boiling point the temperature at which the liquid and vapor phases of a substance are in equilibrium. Equilibrium is a state in which two oppos- ing processes occur at equal rates.

  22. Energy and Heat Energy is the capacity to do work. Energy can exist in different forms, and can change in form: Heat Light Electrical Mechanical

  23. Heat Energy First Law of Thermodynamics Energy is neither created nor destroyed, just changed in form and/or transferred.

  24. Energy Units The joule (J) is the base unit for energy or work (force x distance). 1 J = 1 kgm2/sec2 4.18 J = 1 calorie 1 calorie is the amount of heat energy required to raise the temperature of 1 g of water by 1C.

  25. Specific Heat Specific heat is the amount of heat energy required to raise the temperature of 1.00 gram of a substance by 1.00C. It takes 4.18 J to raise the temperature of 1.00 g of water by 1.00C. 4.18 J/gC is the specific heat of water.

  26. Specific Heat 4.18 J/gC very high! Metals have specific heats of 0.1 to 1.0 J/gC; They conduct heat. Most nonmetallic materials are insulators, with specific heats of 1 to 2 J/gC.

  27. Specific Heat How much heat is absorbed if I heat 100 g of water from 25.0C (room temperature) to 100.0C (boiling point of water)? How much heat is given off if I cool 1.00 lb (454 g) of iron metal from 100.0C to 25.0C? Specific heat of iron is 0.444 J/gC.

  28. Specific Heat What is the specific heat of a rock? Its mass is 125 g. I heat it to 100.0C in a boiling water bath, then drop it into 100.0 g of water that's at 20.0C. The water temperature rises to 30.0C.

  29. Energy and Changes of State A heating or cooling curve is a graph showing the amount of energy required to change the temperature or phase of a given amount of a substance.

  30. Heating and Cooling Curves

  31. Energy and Changes of State In an endothermic phase change: heat energy is absorbed by a substance its particles gain kinetic energy forces between particles are overcome The substance goes into a less ordered state. It melts, boils, or sublimes.

  32. Energy and Changes of State In an exothermic phase change: heat energy is given off by a substance its particles lose kinetic energy forces between particles can act The substance goes into a more ordered state. It freezes, condenses, or deposits.

  33. Energy and Changes of State During an endothermic phase change, all the energy being supplied to the substance is used to disrupt forces between particles. No temperature change is observed. The temperature of a substance that is present in two phases will remain constant until all of one phase is consumed.

  34. Energy and Changes of State During an exothermic phase change, all the energy being released by the substance is allowing intermolecular forces to bring particles to a more ordered state. No temperature change is observed. The temperature of a substance that is present in two phases will remain constant until all of one phase is consumed.

  35. Heats of Fusion and Vaporization The amount of heat required to cause a phase change in a given material is a physical prop-erty of that material. Heat of fusion, Hf = heat to convert one gram of substance from solid to liquid at its melting point. Units: J/g Heat of vaporization, Hv = heat to convert one gram of substance from liquid to gas at its boiling point. Units: J/g

  36. Heats of Fusion and Vaporization How much energy will be consumed when 150 g of water at 100C is boiled to steam, also at 100C? Hvwater = 2260 J/g How much energy will be released if 250 g of water freezes to ice, all at 0C? Hfwater = 334 J/g

  37. Heating and Cooling Curve Calculations How much energy will be consumed when 200 g of ice at -5.0C is converted to steam at 120.0C? Specific Heats, J/gC Ice: 2.09 Water: 4.18 Steam: 2.03

  38. Evaporation and Boiling What is really going on when a substance goes from the liquid to the gas phase or vice-versa? In evaporation, some particles (atoms or molecules) have enough kinetic energy to overcome cohesive forces and escape from the surface of the liquid.

  39. Evaporation and Boiling What is really going on when a substance goes from the liquid to the gas phase or vice-versa? In boiling, many particles have enough kinetic energy to enter the gas phase. Bubbles of gas form in the bulk liquid.

  40. Evaporation and Boiling

  41. Evaporation and Boiling

  42. Evaporation and Boiling If a liquid is put in a closed container, molec-ules of the liquid will escape into the gas (or vapor) phase. The amount of vapor will depend on: Temperature Higher T increases amount of gas Cohesive forces between molecules Stronger forces decrease amount of gas

  43. Evaporation and Boiling Device for measuring the vapor pressure of a liquid.

  44. Evaporation and Boiling

  45. Evaporation and Boiling A liquid boils when its vapor pressure equals the external pressure on the liquid. The normal boiling point of a liquid is the temperature at which it boils under atmospheric pressure. The boiling point of a liquid will increase or decrease with changes in applied pressure.

  46. Evaporation and Boiling Evaporation is conversion from liquid to vapor at temperatures below the boiling point. Vapor is usually used for the gas phase of compounds that are liquid at room temper-ature and pressure. Gas is the term used for compounds that are not liquid at room temperature and pressure.

  47. Evaporation and Boiling

  48. Evaporation and Boiling

  49. Intermolecular Forces Intermolecular forces are attractive forces that act between molecules (also atoms and ions). There are three types: London Dispersion Forces Dipole – Dipole Interactions Hydrogen Bonds Intermolecular forces are hierarchical and additive.

  50. Dipole-dipole forces are electrostatic forces that occur between polar molecules.