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States of Matter

States of Matter. Chapter 10. Kinetic Theory. Kinetic refers to motion. Kinetic energy is the energy of a moving object. Kinetic theory states that the tiny particles in all forms of matter are in constant motion.

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States of Matter

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  1. States of Matter Chapter 10

  2. Kinetic Theory • Kinetic refers to motion. • Kinetic energy is the energy of a moving object. • Kinetic theory states that the tiny particles in all forms of matter are in constant motion.

  3. A gas is composed of small particles that have no attractive or repulsive forces between them. • Move in constant random motion, but in straight paths.(Gases fill their containers regardless of shape or volume. Uncontained gases diffuse into space with out limit.) • During collisions kinetic energy is transferred without loss from one particle to another.

  4. Gas Pressure • Defined as the force exerted by a gas per unit surface area of an object. • Vacuum – empty space, no pressure (outer space) • Air exerts pressure on earth because of gravity. • Air pressure increases the higher you go.

  5. The SI unit of pressure is the pascal (Pa). • mmHg & atmosphere (atm) are other units of pressure. • 1 atm = 760 mmHg = 101.3 kPa

  6. Kinetic Energy & Temperature • When particles are heated they speed up therefore increasing temperature.

  7. Graphs • Horizontal axis (x-axis) represents the independent variable. • Vertical axis (y-axis) represents the dependent variable. • Range is the minimum and maximum values represented in the graph.

  8. Making Graphs • Choose appropriate ranges • The intervals should be convenient numbers (1, 5, 10) • Draw and label axis • Plot data • Connect data with smooth curve. (does not have to touch all data points)

  9. Nature of Liquids Liquids: 1. Particles that make up liquids are in motion. 2. Particles have an attractive force between them, but able to move. 3. This force is called intermolecular force.

  10. Evaporation: • conversion to a gas from a liquid that is not boiling 1. Vaporization: transition of a liquid to gas 2. When heat is added vaporization occurs faster. 3. Liquid – vapor (gas)

  11. Boiling Points: The temperature at which the vapor pressure of liquid is just equal to the external pressure. • Boiling points change when external pressures change. (less pressure = lower boiling points) 2. The temperature of a boiling liquid never rises above its boiling point.

  12. Nature of Solids Solids: • Particles are in fixed points; tend to vibrate in their spots. • Particles tend to be in highly organized patterns. • Melting point – the temperature at which the solid melts

  13. As heat is added the organization of particles begins to break down and solid melts. 5. Melting: Solid – liquid Freezing: Liquid – solid 6. Not all solids melt; some decompose

  14. Crystal Structures • The atoms, molecules or ions are arranged in an orderly, repeating 3-D pattern. • Unit Cell: smallest group of particles within a crystal; repeating group (simple cubic, body centered cubic & face centered cubic) .

  15. Crystal Structures 1. All crystals have a regular shape. 2. Crystals have sides or faces. 3. Crystals are classified in 7 groups

  16. Allotropes: • Two or more different molecular forms of the same element in the same physical state. • Ex. Carbon --- diamond - graphite

  17. Amorphous Solids: • Lack an ordered internal structure • Examples: Rubber, plastic, asphalt • Particles are randomly arranged.

  18. Change of State • Phase Diagram: gives the conditions of temperature & pressure at which a substance exists as a solid, liquid, & gas. • Triple Point: a set of conditions at which all three states can exist with one another.

  19. Sublimation: the change of a substance from a solid to a vapor without passing through the liquid state. Examples: Solid carbon dioxide (dry ice) Moth balls

  20. Add to Gas Pressure: • Barometers: devices commonly used to measure atmospheric pressure. This pressure depends on the weather.

  21. Gas Laws Chapter 12

  22. REMEMBER: • The SI unit of pressure is the pascal (Pa). • mmHg & atmosphere (atm) are other units of pressure. • 1 atm = 760 mmHg = 101.3 kPa

  23. Properties ofGases • Compressibility:measure of how much the volume of matter decreases under pressure. • Gases are easily compressed. • Gases expand to take the shape & volume of the container. • Gas particles are in constant motion.

  24. Variables that describe Gases • Pressure (P) in kPa • Volume (V) in L • Temperature (T) in K • Moles (n)

  25. Gas Pressure • Increase the # of gas particles; increases the gas pressure • Doubling the # of gas particles; doubles the pressure. • Once the pressure exceeds the strength of the container the container will rupture. (Direct Relationship)

  26. If you let gas out. Then the pressure drops. • When a sealed container of gas under pressure is opened. Gas inside moves from the region of higher pressure to the region of lower pressure.

  27. Volume: raise the pressure exerted by contained gas by reducing its volume. (Inverse Relationship)

  28. Temperature: 1. Raising the temperature of an enclosed gas increases gas pressure. 2. If the temperature doubles then the pressure doubles. (Direct Relationship)

  29. Standard Temperature & Pressure (STP) • Temperature – 273 K • Volume – 22.4 L • Pressure – 101.3 kpa = 1 atm = 760 mmHg = 760 torr

  30. Boyle’s Law • Pressure – Volume Relationship P1V1 = P2V2 • Example: A high – altitude balloon contains 30.0 L of helium gas at 103 kPa. What is the volume when the balloon rises to an altitude where the pressure is only 25.0 kPa? (Assume STP)

  31. P1 = 103 kPa V2 = ? L V1 = 30.0 L P2 = 25.0 kPa (103 kPa)(30.0 L) = (25.0 kPa)(V2) V2 = 124 L

  32. Charles’s Law Temperature – Volume Relationship V1 = V2 T1 = T2 **Temperature in Kelvin. Example: A balloon inflated in a room at 24 oC has a volume of 4.00 L. The balloon is then heated to a temperature of 58 oC. What is the new volume if the pressure remains constant?

  33. V1 = 4.00 L V2 = ? L T1 = 24 oC = 297 K T2 = 58 oC = 331 K (4.00L) = V2 (297 K) (331 K) V2 = 4.46 L

  34. Combined Gas Law Combines the 3 gas laws P1 V1 = P2 V2 T1 = T2 Example: The volume of a gas – filled balloon is 30.0 L at 40 oC and 153 kPa pressure. What volume will the balloon have at standard temperature and pressure (STP)?

  35. V1 = 30.0 L V2 = ? L T1 = 40 oC = 313 K T2 = 273 K P1 = 153 kPa P2 = 101.3 kPa (153 kPa)(30.0 L) = (101.3 kPa)(V2) ( 313 K ) (273 K ) V2 = 39.5 L

  36. Ideal Gas Law (Universal Gas Law) • Adding # of moles PV = nRT P = pressure V = volume n = # of moles R = ideal gas constant (8.31 L x kPa / K mole) T = temperature

  37. Example: You fill a rigid steel cylinder that has a volume of 20.0 L with nitrogen gas (N2) to a final pressure of 2.00 x 104kPa at 28 oC. How many moles of N2 does the cylinder contain?

  38. P = 2.00 x 104kPa n = ? moles of N2 V = 20.0 L T = 28 oC (2.00 x 104kPa)(20.0 L) = n(8.31)(301 K) N = 160 moles N2

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