Chapter 3 States of Matter

1. Chapter 3States of Matter Chemistry 1

2. Solids, Liquids, and Gases 3.1

4. Describe the States of Matter 3.1 • Use shape and volume as clues to which state • Solids – Definite shape and volume • Liquids- Definite volume, no definite shape • Gases – No definite shape or volume • Other States: • 99% of matter in the universe is plasma • Found on the sun and stars • Extremely high temperatures • Bose-Einstein Condensate (BEC) – 5th state of matter that exists at extremely low temperature where atoms behave as a single particle – predicted by Bose in 1920s and Einstein decades before • 1st produced at CU in 1995

5. Amorphous Solids • •Solid or liquid?? • •Do not have crystals! • •Can and will “flow” over time. • •Example: glass, wax, plastic • •Small temperature changes affect these (plastic on a cold day) • •Officially, these are liquids!!!

6. Kinetic Theory 3.1 • Energy an object has due to its motion • An object that moves has ke • The faster object moves, the greater the ke • All particles of matter are in constant motion

7. Explaining the Behavior of Gases 3.1 • Motion in Gases • Particles never at rest • Average speed of particles is 1600 km/hr • Atoms moves in straight line until collides with something • Collision – 1 atom loses ke and slows down; other atom gains ke and speeds up

8. Kinetic Theory of Gases 3.1 • Particles in a gas are in constant, random motion • Motion of 1 particle, unaffected unless particles collide • Forces of attraction ignored under normally conditions

9. Explaining the Behavior of Liquids 3.1 • Particles always moving • Move slower than a gas particle • Greater mass = slower speed • More closely packed particles • Force of attraction to keep particles close together

10. Explaining the Behavior of Solids 3.1 • Particles vibrate around a fixed location • Strong attraction among particles

11. The Gas Laws 3.2

13. Pressure 3.2 • The result of a force distributed ovan an area • SI unit = force/area • Force = Newtons (N) • Area = square meters (m2) • N/m2 = Pascal (Pa)

14. Factors that Affect Gas Pressure 3.2 • Temperature • Increase in temp ke increases  particles collide more  increase in pressure • Think of tires in a car • Volume • Reducing volume, increases pressure • Think of your lungs • Number of particles • Increasing particles, increases pressure

15. Charles’s Law 3.2 • Jacques Charles (1746-1823) • Investigated gases • Absolute Zero – 0 K • No scientist has produced a temp of 0K • Volume of gas directly proportional to temp in K if pressure and particles are constant • TEMP MUST BE IN KELVIN • V1/T1 = V2/T2

17. Boyle’s Law • Robert Boyle –Ireland • 1st to describe relationship of pressure and volume of gas • Volume of a gas is inversely proportional to its pressure if the temp and the number of particles are constant • P1V1=P2V2

19. Combined Gas Law 3.2 • (P1V1)/T1=(P2V2)/T2 http://www.nclark.net/GasLaws

20. Phase Changes 3.3

22. Characteristics of Phase Changes 3.3 • Reversible physical change that occurs when a substance changes from 1 state of matter to another • 6 common phase changes: melting freezing, vaporization condensation, sublimation, and deposition

24. Temperature and Phase Changes 3.3 • The temp of substance does not change during a phase change

25. Energy and Phase Changes 3.3 • Energy absorbed or released during a phase change • Endothermic – system absorbs energy • Ex: melting • Heat of Fusion – amount of energy absorbed from a substance • Fusion = melting • Exothermic – Releases energy • Ex: freezing

26. Melting and Freezing 3.3 • Molecules become less orderly when substance melts and more orderly when substance freezes • Melting – heat flows from air to ice, ice gains energy, molecules vibrate more quickly  molecules gain enough energy to move from fixed position • Increases average ke • Freezing – energy flows from water to air, water cools, average ke decreases, molecules drawn to an orderly position • Decreases average ke

27. Vaporization and Condensation 3.3 • Vaporization – liquid to a gas • Endothermic process (substance must absorb energy) • Heat of Vaporization – Amount of energy it takes to change from a liquid to a gas • 2 process: boiling and evaporation • Evaporation – Liquid to a gas at temp below boiling point • Molecules at the surface moving fast enough to escape the liquid • The greater the surface area, the faster the water evaporates • Vapor Pressure – pressure caused by the collisions of this vapor and the walls of the container – occurs in a closed container • Increases as the temp increases

28. Boiling • Vapor pressure and temp increases • When vapor pressure = atmospheric pressure, water boils • Water molecules move faster when heated, bubbles (boiling) are water vapor • Water vapor is less dense so bubbles rise • Bubbles burst and releases water vapor into the air • Boiling point depends on the atmospheric pressure • At high elevations, the atmospheric pressure is lower so boiling point is lower… so food takes longer to cook at this lower temp!

29. Condensation • Phase change in which a substance changes from a gas or vapor to a liquid • Exothermic process • Morning dew • Foggy glass after a shower

30. Sublimation and Deposition 3.3 • Sublimation – substance changes from a solid to a gas or vapor • Endothermic change • Ex: dry ice (form of carbon dioxide) • Deposition – gas or vapor changes directly into a solid • Exothermic change • Ex: frost to form on windows