States of Matter

1. States of Matter

2. 3-1 Solids, Liquids and gases • Materials can be classified as solids liquids, or gases, based on whether their shapes and volumes are definite or variable. • What property could you use to distinguish the liquid or gas from the solids in a level? • Why is the air bubble above the liquid in the tube? • When might the tube in the middle of the level be used? • When might the tube on the right be used?

3. Describing the states of matter • Solids – definite shape and volume • Liquids – definite volume but not a definite shape • Gases - neither a definite shape nor volume • Other States of Matter – plasma Class Participation Opportunity: complete a biography of a scientist report on Satyendra Bose or Albert Einstein being certain to tie in their connection to other states of matter Reflection: Use textbook pages 69-70 to organize this information in a scientific way including examples(table, graphic organizer, foldable)

4. Kinetic theory • Kinetic – in Greek, “to move” • Kinetic energy is the energy an object has due to its motion. • Faster an object moves = greater its kinetic energy • Particles inside the object and around the object are moving too even though they can’t be seen Kinetic theory of matter says that all particles of mater are in constant motion.

5. Explaining the behavior of gases • Motion in Gases • Unlike billiard balls = gas particles are never at rest • Like billiard balls = each particle moves in a straight line Even though there are forces of attraction among all particles in matter, since the particles in gases are moving so fast, the attractions are too weak to from moving off course. • When one atom collides with another it transfers its kinetic energy to the other • Kinetic Theory of Gases The constant motion of particles in a gas allows a gas to fill a container of any shape or size.

6. Explaining the behavior of liquids A liquid takes the shape of its container because particles in a liquid can flow to new locations. The volume of a liquid is constant because forces of attractions keep the particles close together. How is the movement of students in a crowded hallway like the behavior of liquids? • Why don’t particles in a solid move as freely as particles in a gas at room temperature?

7. Explaining the behavior of solids Solids have a definite volume and shape because particles in a solid vibrate around fixed locations. What stayed the same and what changed between the two photographs pictured here? Demonstration with Marbles What state of matter does this model represent? How are the particles in a solid like the marbles in this model? How well does this part of the model represent the behavior of particles in a solid?

8. 3-2 the gas laws What causes gas pressure in a closed container? What factors affect gas pressure? How are the temperature, volume, and pressure of a gas related?

9. Pressure Pressure is the result of force distributed over an area. The SI unit of pressure is derived from force and area. Force is measured in newtons (N). Area is measured in square meters (m²). Pressure is N/m² but the SI unit for pressure is the pascal which is written (Pa).

10. Factors the affect Gas pressure Temperature: Raising the temperature of a gas will increase its pressure if the volume of a gas and the number of particles are constant. Volume: Reducing the volume of a gas increases its pressure if the temperature and the number of particles are constant. Number of Particles: increasing the number of particles will increase the pressure if the temperature and volume are constant.

11. Charles’ Law • Volume of a gas is directly proportional to its temperature in kelvins if the pressure and the number of particles of the gas are constant. • Research French physicist Jacques Charles (1746-1823) for class participation credit.

12. Boyle’s law • The volume of a gas is inversely proportional to its pressure if the temperature and the number of particles are constant.

13. Compare and Contrast Charles’ Law with Boyle’s Law

14. The combined gas law Describes the relationship among temperature, volume, and pressure of a gas when the number of particles is constant.

16. Class Participation Opportunity Research hot-air balloons. Be sure to connect what you learn to what we’re studying.

17. 3-3 Phase changes What are six common phase changes? What happens to a substance’s temperature and a system’s energy during a phase change? How does the arrangement of water molecules change during melting and freezing? How are evaporation and boiling different?

18. Characteristics of phase changes Phase changes are reversible physical changes occurring when substances change from one state of matter to another. When at least two states of the same substance are presents, scientists describe each different state as a phase. Melting, freezing, vaporization, condensation, sublimation, and deposition are six common phase changes.

19. Review What do the phase changes with red arrows have in common? What do the phase changes with blue arrows have in common?

20. Temperature and Phase Changes • Temperature during a phase change does not change. • In what state is the naphthalene at 20ºC? • In what state is the naphthalene at 90ºC? • What are the melting and freezing points of naphthalene? • What would the heating curve for naphthalene look like if the graph were extended beyond a temperature of 100ºC?

21. Energy and Phase Changes • During a phase change, energy is transferred between a substance and its surroundings. • It is either absorbed or released.

22. Melting and freezing In water, the arrangement of molecules becomes less orderly as water melts and more orderly as water freezes. • Melting - in ice, molecules keep in a fixed position - As the ice melt, heat flows from the air to the ice and as the ice gains energy, the molecules vibrate more quickly. - melting point of water is 0ºC • Freezing- water placed in a freezer = kinetic energy of its molecules decreases - When all molecules are in an orderly arrangement, freezing is complete.

23. Vaporization Liquid to a gas = Vaporization Endothermic process: liquid must absorb energy to change Heat of vaporization = 1 gram of water gaining 2261 joules of energy when it vaporizes - varies from substance to substance A substance changes from a liquid to a gas to a liquid over and over again. During these phase changes, energy flows from the inside to the outside.

24. Two Vaporization Processes • Evaporation-at the surface of a liquid and occurs at temperatures below the boiling point Discuss vapor pressure in a closed container • Boiling-water temp. and vapor pressure increase When vapor pressure and atmospheric pressure are equal , the boiling point is reached. At 100ºC some molecules have enough energy to overcome the attraction of the other molecules and since water vapor is less dense, the gas bubbles rise to the surface.

25. Condensation Phase change where gas or vapor changes into liquid • an exothermic process Examples: • “the cloud” on the bathroom mirror after a hot shower • Dew on grass in the morning

26. Sublimation Sublimation = phase change from solid to gas without changing to a fluid first endothermic Example: dry ice Mosquito trap baited with dry ice

27. Deposition Deposition = gas or vapor changes directly into a solid without forming a liquid first exothermic change Example: frost on windows