Gases, Liquids and Solids

1
Gases, Liquids and Solids

• States of Matter, Chapter 10

2
Objectives

1. Describe the motion of gas particles according to
   the kinetic theory
2. Interpret gas pressure in terms of kinetic theory

3
Describing a Gas

• A gas expands to completely fill the container.
• A gas can flow.
• A gas can be compressed.
• Between the particles of a gas there is empty
  space.

4
Kinetic Molecular Theory

• In an attempt to understand how molecules or
  atoms of gases behave, we can refer to the
  kinetic molecular theory of gases.
• This theory describes what we refer to as an
  ideal gas, one that behaves according to a set of
  mathematical expressions of these ideas.

5
In the kinetic molecular theory

1. Gases consist of molecules whose volumes are
   negligible compared with the volume occupied by
   the gas.
2. Molecules of a gas are in constant, random
   motion.
3. The average kinetic energy of gas molecules is
   determined by the temperature of the gas.
4. Gas particles collide with each other and with
   the walls of the container without a loss of
   energy (the have perfectly elastic collisions).
5. We assume that the particles of an ideal gas
   exert no attractive or repulsive forces on each
   other.

6
Gas Pressure

• Gas pressure is the force exerted by a gas per
  unit surface area of an object.
• It is caused by gas particles colliding with an
  object.
• An empty space, with no particles and no
  pressure, is a vacuum.
• Atmospheric pressure results from air particles
  colliding with an object.

7
Atmospheric Pressure

• Changing weather conditions

8
Atmospheric Pressure

• Changing altitude

9
Measuring Pressure

• Barometer device that measures atmospheric
  pressure
• Invented by Evangelista Torricelli in 1643
• The pressure was measured by finding the height
  of the Hg column (using the unit mmHg).

10
Measuring pressure in a sample of gas

• A manometer measures the pressure of a gas in a
  container.

11
Units of Pressure

• Metric unit of pressure is the pascal (Pa).
• A more familiar units include the atmosphere
  (atm) and the millimeter of mercury (mmHg).
• Know this
• 1 atm 760.0 mm Hg 760.0 torr 101.3
  kilopascal (kPa)

12
Problem Convert 1.5 atm to the units of kPa and
mmHg.

• 1.5 atm x 101.3 kPa/1 atm
• 1.52 x 102 kPa
• 1.5 atm x 760 mmHg/1 atm
• 1.14 x 103 mmHg

13
Objectives

1. Describe the nature of a liquid in terms of the
   attractive forces between the particles
2. Differentiate between evaporation and boiling of
   a liquid

14
A Model for Liquids

• Liquid molecules are free to slide past each
  other, so liquids, like gases, can flow.
• Liquids take the shape of their container.
• Liquids are condensed, not compressible.

15
Vaporization

• Vaporization is the conversion of a liquid to a
  gas or vapor.
• If the conversion occurs at the surface of a
  liquid that is not boiling, the process is called
  evaporation.
• Evaporation is a cooling process. The particle
  with the highest kinetic energy tend to escape
  first, leaving cooler particles behind.

16
Vapor Pressure

• Vapor pressure that results from the vapor
  particles colliding with the walls of the
  container.
• It is a force due to a gas above the surface of a
  liquid.
• (Vapor is the gaseous state of a substance that
  is a liquid or solid at room temperature).

17
Phase changes with liquids

• Evaporation
• Liquid ? Vapor (gas)
• Condensation
• Vapor (gas) ? Liquid

18
Equilibrium vapor pressure

• When the rate of vaporization is equal to the
  rate of condensation, no net change in vapor
  pressure occurs.
• This is a dynamic equilibrium.

19
Temperature affects the rate of vaporization

• Increasing the temperature of a contained liquid
  increases the vapor pressure over the surface of
  a liquid.

20
Boiling

• Boiling occurs when the vapor pressure is equal
  to the external pressure.
• In Denver, where atmospheric pressure is low,
  boiling occurs at a lower temperature.
• In a pressure cooker, boiling occurs at a higher
  temperature.

21
Boiling point

• the temperature at which boiling occurs.
• Normal boiling point is the temperature at which
  boiling occurs at 1 atm pressure (normal
  pressure).
• Boiling is also a cooling process.

22
The Nature of Solids

• The particles in a solid tend to vibrate about
  fixed points. They are not free to flow.
• Most solids are highly organized, dense and
  incompressible.

23
Melting point.

• the temperature at which a solid turns into a
  liquid.
• Melting.
• Solid ? liquid
• Freezing
• Liquid ? solid

24
Objectives

1. Describe how the degree of organization of
   particles distinguishes solids from liquids and
   gases
2. Distinguish between a crystal lattice and a unit
   cell
3. Explain how allotropes of an element differ

25
Solids

• Crystal
• Unit cell
• Allotrope
• Several forms of an element (carbon)
• Amorphous solids (glass, soot)
• Supercooled solids
• Network solids (diamonds, graphite)

26
The Solid State Types of Solids

• Crystalline solids

27
A. The Solid State Types of Solids
28
The Solid State Types of Solids
29
Bonding in Solids
30
Changes of state

• Solid to liquid
• melting
• Liquid to solid
• freezing
• Liquid to gas
• boiling, evaporating
• Gas to liquid
• condensation
• Solid to gas
• sublimation
• Gas to solid
• deposition

31
Phase diagrams

• A phase diagram is a graph showing the
  relationship between pressure and temperature for
  solid, liquid and gas states of a substance.
• Triple point describes the set of conditions
  where all three phases, solid, liquid, and gas,
  can exist in equilibrium with each other.

32
Phase diagram for CO2
33
Phase diagram for water
34
Note

• The solid-liquid equilibrium line for CO2 has a
  negative slope, while the solid-liquid line for
  water has a positive slope.
• This shows that increasing the pressure on a
  solid sample of water can cause the solid to
  melt, while a solid sample of CO2 would remain a
  solid.
• This unique property of water allows a skater to
  glide over the surface of ice on a layer of water
  under their thin skate blades.

35
Heating curve

• A heating curve shows the relationship between
  temperature and time as heat is added to a
  system.
• Notice that as a phase change occurs, even though
  heat is being added, no change in temperature
  occurs.

36
Heating curve
37
Heating curve for water