States of Matter: Solids, Liquids, Gases, Plasma, and BEC

1. Chapter 3 States of Matter

2. Chapter 3 Sections • 3.1 Solids, Liquids, and Gases • 3.2 The Gas Laws • 3.3 Phase Changes

3. 3.1 Solids, Liquids, & GasesKey Concepts • How can shape and volume be used to classify materials? • How can kinetic theory and forces of attraction be used to explain the behavior of gases, liquids, and solids?

4. 3.1 Describing the States of Matter • There are five states of matter. • Solid, Liquid, Gas, Plasma & BEC • The state of matter depends on a materials temperature. • Each state has characteristics that are used to identify it. • Materials can be classified as solids, liquids, or gases based on whether their shapes and volumes are definite or variable.

5. 3.1 Solids • Solids have a definite shape. • Solids have a definite volume. • The particles of a solid are close together. • They do not have enough energy to move, they only vibrate.

6. 3.1 Liquids • Liquids have no definite shape. • Liquids do have a definite volume. • The particles of a liquid are close together. • They have enough energy to move. They slide past one another.

7. 3.1 Gases • Gases have no definite shape. • Gases have no definite volume. • The particles of a gas have enough energy to separate completely. • They are free to move in all directions.

8. 3.1 Plasmas • Plasmas are gas-like mixture of positively and negatively charged particles • Plasmas like gases have no definite shape nor definite volume. • As matter is heated to very high temperatures, the particles begin to collide very violently. • These violent collisions break the particles up into smaller particles that are charged.

9. 3.1 Bose-Einstein Condensate • Einstein read a paper by a physicist from India, Satyendra Bose and predicted a fifth state of matter. • At temperatures near –273OC, groups of atoms behave as one particle. • This is called a Bose-Einstein condensate or BEC.

10. 3.1 Kinetic Theory of Matter • Kinetic energy is the energy an object has due to its motion. • The kinetic theory of matter says that all particles of matter are in constant motion.

11. 3.1 Explaining the Behavior of Gases • Particles of gases are in constant motion. • These particles will move in straight line paths until they contact another particle. • Kinetic energy is transferred during those collisions. • There are forces of attraction among particles in all matter.

12. 3.1 Kinetic Theory of Gases • The constant motion of particle in a gas allows a gas to fill a container of any shape or size. • The kinetic theory as applied to gases has three main points: • Particles of gases are in constant, random motion. • The motion of one particle is unaffected by the motion of other particles unless the particles collide. • Forces of attraction among particles in a gas can be ignored under ordinary conditions.

13. 3.1 Explaining the Behavior of Liquids • Particles of liquids are in constant motion but much slower than gases. • The particles are packed more closely. This allows the forces of attraction to affect the motion of the particles. • A liquid takes the shape of its container because particles in a liquid can flow to new location. The volume of a liquid is constant because forces of attraction keep the particles close together.

14. 3.1 Explaining the Behavior of Solids • Particles of solids are in constant motion but they can’t move about. • The particles are packed even more closely. This allows the forces of attraction to lock the particles to one location. • Solids have a definite volume and shape because particles in a solid vibrate around fixed locations.

15. 3.2 The Gas Laws Key Concepts • What causes gas pressure in a closed container? • What factors affect gas pressure? • How are the temperature, volume, and pressure of a gas related?

16. 3.2 Pressure • Pressure is the result of a force distributed over an area. • P = (F/A) • Thepascal (Pa)is the SI unit of pressure. This is a very small unit, so kilopascals are used. • Collisions between particles of a gas and the walls of the container cause the pressure in closed container of gas.

17. 3.2 Factors That Affect Gas Pressure • Factors that affect the pressure of an enclosed gas are • its temperature, • its volume, • and the number of its particles.

18. 3.2 Temperature • As the temperature of a gas increases, the particles gain kinetic energy and move faster. • These faster particles collide with the container walls more frequently. • Their increased speed increases the force of the collisions. • Raising the temperature of a gas will increase its pressure if the volume of the gas and the number of particles are constant.

19. 3.2 Volume • As you decrease the volume of a container, the particles collide with the container walls more frequently. • Reducing the volume of a gas increases its pressure if the temperature of the gas and the number of particles are constant.

20. French physicist Jacques Charles collected data on the relationship between temperature and volume of gases. 3.2 Charles’s Law • The graphed data showed a straight line. • He extended his line to see where the volume would be 0. • Charles was able to predict absolute zero was at –273oC using his law.

21. 3.2 Charles’s Law • Gases expand when heated • According to Charles’ Law, the volume of a gas increases with increasing temperature, if the pressure stays the same. • If T then V • If T then V

22. Irish chemist Robert Boyle collected data on the relationship between pressure and volume of gases. The graphed data showed an inverse curve. 3.2 Boyle’s Law

23. 3.2 Boyle’s Law • The pressure of a gas depends on how often its particles strike the walls of the container. • According to Boyle’s Law, if you decrease the volume of a container of gas, the pressure of the gas will increase, if the temperature stays the same. • If V then P • If P then V

24. 3.2 The Combined Gas Law • The relationship described by Boyle’s law and Charles’s law can be described by a single law. • The combined gas law shows the relationship between the temperature, pressure, and volume of a gas, if we keep the number of particle the same.

25. 3.2 The Combined Gas Law

26. 3.2 The Combined Gas Law

27. 3.2 The Combined Gas Law

28. 3.3 Phase ChangesKey Concepts • 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?

29. States of matter are referred to as phases. A phase change is the reversible physical change that occurs when a substance changes from one state of matter to another. Melting, freezing, vaporization, condensation, sublimation, and deposition are six common phase changes. 3.3 Characteristics of Phase Changes

30. 3.3 Temperature and Phase Changes • The temperature of a substance does not change during a phase change. • The heating curve for napthalene (moth balls) demonstrates this fact.

31. 3.3 Energy and Phase Changes • Energy is either absorbed or released during a phase change. • During an endothermic change, the system absorbs energy from its surroundings. Melting is an example of an endothermic change. • How much energy is absorbed depends on the heat of fusion for that substance. • Ice has a heat of fusion of 334 joules/gram.

32. 3.3 Energy and Phase Changes • The heat of fusion works in the reverse way. • During an exothermic change, the system releases energy to its surroundings. Freezing is an example of an exothermic change. • Water releases 334 joules/gram of heat energy as it freezes.

33. 3.3 Melting and Freezing • The arrangement of molecules in water becomes less orderly as water melts and more orderly as water freezes. • The molecules in a solid are vibrating in place. They vibrate faster as more heat energy is gained by the solid. • Some of the molecules will gain enough energy to overcome the forces of attraction between them and move. • When all the molecules have broken free, melting is complete. Only then does the temperature begin to rise.

34. 3.3 Melting and Freezing • As the molecules of liquids begin to lose energy to its surroundings, the molecules begin to slow down. • Some of the molecules will slow down enough energy to become trapped by the forces of attraction between them and stop. • When all the molecules are firmly in place, freezing is complete. Only then does the temperature begin to fall. • Freezing does not always occur at low temperatures.

35. 3.3 Vaporization • The phase change in which a substance changes from a liquid to a gas is called vaporization. • Vaporization is an endothermic process. • The amount of energy absorbed depends on the substances heat of vaporization. • Water has a heat of vaporization of 2261 joules/gram. • There are two vaporization processes: • Boiling • Evaporation

36. 3.3 Evaporation • Evaporation takes place at the surface of a liquid and occurs at temperatures below that boiling point. • Evaporation is the process that changes a substance from a liquid to a gas at temperatures below the substance’s boiling point. • Molecule at the surface move fast enough to escape and become water vapor.

37. 3.3 Evaporation • When water is in a closed container, the water vapor is trapped above the liquid. • The pressure caused by the collisions of the vapor and the container walls is called vapor pressure. • As the temperature increases, so does this vapor pressure.

38. 3.3 Boiling • When water is heated the temperature and vapor pressure rise. • Water will begin to boil when the vapor pressure equals atmospheric pressure. • This occurs at the boiling point.

39. 3.3 Boiling • As the temperature of a liquid rises, the molecules move faster. • Some molecules reach the surface with enough energy to escape the forces of attraction. • The boiling point of a substance depends of the atmospheric pressure. • The lower the atmospheric pressure the lower the boiling point. • Have you ever seen high altitude directions on food containers?

40. 3.3 Condensation • Water vapor will cool when it comes into contact with surfaces. • Condensation is the phase change in which a substance changes from a gas to a liquid. • This process is an exothermic one. • Morning dew is a result of this process.

41. 3.3 Sublimation • Dry ice is the common name for the solid form of carbon dioxide. It will go directly from the solid phase to the gas phase. • Sublimation is the phase change in which a substance changes directly from a solid to a gas or vapor. • This process is an endothermic one.

42. 3.3 Deposition • When a gas changes directly into a solid without changing to a liquid, the phase change is called deposition. • This process is an exothermic one and is a reverse of sublimation. • Frost is an example of deposition.