1. Gases, Liquids and Solids
States of Matter
Gases, Liquids and Solids

2. Kinetic Molecular Theory of Liquids and Solids
Liquids and solids are condensed states.
Liquids
Have definite volume and assume the shape of its container
Much denser than gases
Are more difficult to compress than gases
Molecules
are close together
little of empty space between molecules
are held together by attractive intermolecular forces
are in constant motion
Move/slide one past another freely so liquid can flow
collision rate is higher than in gases

3. Kinetic Molecular Theory of Liquids and Solids
Almost uncompressible
Definite shape and volume
Most solids are denser than liquids
Particles
are more closely packed than in liquids
are arranged in highly organized order
Are always in constant motion
vibrate around fixed locations
Held together by strong intermolecular forces.

5. Structure of Solids Most of solids are crystalline.
Particles are arranged in an orderly, repeating, three- dimensional pattern.
The shape of the crystal depends on the arrangement of the particles within it.
The smallest group of particles within the crystal that retains the geometric shape of the crystal is called a unit cell.
A crystal lattice is a repeating pattern of unit cells.

7. Amorphous solids
Amorphous means without any defined form or structure.
Amorphous solids: Glass, plastic,

8. Allotropes (more then one form depending on structure) of Carbon
Diamond
Graphite

9. Intermolecular Forces (IMF)
Are attractive forces between molecules
Are much stronger in liquids and solids than in gases.
The more polar molecules are the stronger intermolecular forces are.
Strength of intermolecular forces determines the state of matter.
Energy is required to break IMF and change states of matter.
Melting and boiling points of any substance depend on strength of IMF in the substance.

10. Strength of Intermolecular Forces (IMF)
Metallic bond (in metals)
Ionic bond (in ionic compounds)
Hydrogen bond (in water and some other compounds)
Van der Waals forces (dipole-dipole, dispersion)
Interaction between partial charges in polar covalent molecules
The more polar molecules are the stronger intermolecular forces are.
Strength

11. Metallic Solids/Metallic Bond
Metallic solids are composed of metal atoms that are held together by metallic bonds.  
The electrons in metallic solids are delocalized.  
It means that valence electrons move freely between metallic cations.
A good picture of a metal solid is that of cation in a sea of electrons. 

12. Strength of the Intermolecular Forces
Phase Changes
Strength of the Intermolecular Forces

13. Phase Changes (state of matter changes)

14. State Changes Simulations

15. Phase Changes
Vaporization is the process in which liquid is transformed into a gas.
Condensation is the process by which a gas or a vapor becomes a liquid
Melting (fusion) is the process of solid becoming a liquid.
Freezing (solidification) is the process when a liquid becomes a solid.
Sublimation is the change of substance from a solid to a vapor without passing through the liquid state.
Deposition is the process by which a gas or a vapor turns straight into a solid.

16. Kinetic Molecular Theory and Phase Changes

17. What happens to IMF in the process of melting?

18. Melting As solid absorbs heat particles of solid vibrate faster
IMF are getting weaker (energy is required to weaken the IMFs)
The distance between particles increases
During the process of melting energy is spent to weaken IMF between particles and increase distance between particles.

19. Melting Point, Freezing Point
The melting point is the temperature at which solid turns into a liquid.
The freezing point is the temperature at which the liquid turns into a solid.
Tmeltng = T freezing
Melting point depends on the strength of IMF in a substance

20. Vaporization
Evaporation
Boiling

21. Vaporization (evaporation and boiling)
Vaporization is the process in which liquid is transformed into a gas.
How does this process happen?
What do we know about liquids and gases that will help us to analyze the process?
Liquids have quite strong IMFs between particles.
Gasses have not IMF between particles.
To leave the liquid phase and get to the gas phase particles need to break IMFs that holds them within the liquid

22. Vaporization (evaporation and boiling)
Molecules with sufficient kinetic energy
overcome intermolecular forces and
escape from the liquid into gas phase,
vapor over the surface of the liquid is created.
Vaporization rate depends on temperature.
the higher the temperature, the greater kinetic energy of particles, more particles are leaving the liquid.

23. Vaporization Rate Vaporization rate depends on temperature.
the higher the temperature,
the greater kinetic energy of particles,
more particles are leaving the liquid.

24. Vaporization (evaporation and boiling) as endothermic process
Particles with higher kinetic energy leave the liquid phase
Particles with lower kinetic energy stay in the liquid phase.
Temperature of the liquid decreases during vaporization. Why?
Constant supply of energy is needed to break IMFs and support vaporization.
Vaporization is an endothermic process (energy is consumed).

25. Evaporation Evaporation happens at the surface of a liquid,
at any temperature.

26. Equilibrium What is equilibrium?
Equilibrium - a state of balance between opposite processes.
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Some molecules evaporate,
Some molecules condense by returning to the liquid phase.
Equilibrium is established.
During condensation molecules strike the liquid surface and become trapped by intermolecular forces in the liquid.

28. Boiling
When a liquid is heated to a temperature when all molecules have enough kinetic energy to vaporize, the liquid begins to boil.
Boiling happens
at a boiling point,
Vapor forms within the liquid.
What is in the bubble?

29. Boiling At the boiling point bubbles of vapor form within the liquid.
The pressure within the bubble is due to the vapor pressure .
When vapor pressure in the bubble equals external pressure, the bubble rises to the surface of the liquid and bursts.

30. Vapor Pressure
Vapor pressure is a pressure of vapor particles exerted on the walls of a container.
An increase in the temperature of a liquid increases the vapor pressure.

31. Boiling Point, Vapor Pressure, and Atmospheric Pressure
Demo: Water boiling in a vacuum yc0
Demo observation conclusion:
Lower than normal atmospheric pressure – water boils at a lower than 100℃ temperature.
Higher than normal atmospheric pressure – water boils at at a higher than 100℃ temperature.
Boiling point of liquid depends on external/atmospheric pressure.

32. Analyze the diagram

33. Boiling Point (temp. of boiling)
The boiling point is the temperature at which the vapor pressure of a liquid is equal to the external pressure.

34. Normal Boiling Point
Normal boiling point is a boiling point when the external/atmospheric pressure is 1 atm.

35. Vapor Pressure vs. Temperature

36. Boiling Point Summary Boiling Point depends on
Strength of intermolecular forces in a substance
External/atmospheric pressure

37. Boiling: analyze the process from the point of IMF
Liquids have quite strong IMFs
Gasses have no IMFs
Particles of liquid should break IMFs to move from the liquid phase to the gas phase.
Energy is required to break IMFs.

38. Boiling Boiling is a process that takes time and energy.
Why do we need to supply energy all the time to support boiling? Try to think about the process considering kinetic energies of particles in a liquid.
When particles with higher kinetic energy leave the liquid
the temperature of the liquid decreases,
constant supply of energy is needed to break IMFs and support boiling temperature.
Boiling is an endothermic process.

39. Vaporization

40. Heating Curve of Water

43. Heating Curve (at 1 atm pressure)

45. Heating Curve
Heating Curve shows how the temperature is changing through the phase transition with energy (heat) change.

46. Solid Phase Only solid is present Heat is added.
Temperature of the solid is increasing
Kinetic energy of the particles in the solid is increasing.
Particles within the solid vibrate faster.

47. Melting Phase change: a solid turns into a liquid.
Solid and liquid are present at the same time.
Although the heat is added, the temperature does not change.
Energy is spent to weaken intermolecular forces of attraction between the particles.
Distances between particles are increasing.

48. Liquid Phase Only liquid is present Heat is added.
Temperature of the liquid is increasing.
Kinetic energy of the particles is increasing.
Particles move faster.

49. Boiling Phase change: liquid turns in gas (vapor)
Liquid and gas are present at the same time.
Heat is added, but the temperature does not change.
Energy is spent on breaking intermolecular forces of attraction between particles.
Distances between particles are increasing.

50. Gas Phase Only gas is present. Heat is added.
Temperature is increasing.
Kinetic energy of the particles is increasing.
Particles move faster.

52. Heating Curve of an Unknown Substance

55. Phase Diagram

56. Phase Diagram
A phase diagram summarizes the conditions at which a substance exists as a solid, liquid, or gas.

57. Each region on diagram represent a pure phase (solid, liquid, or gas/vapor)

58. Line between two regions indicates conditions under which two phases (states of matter) can exist in equilibrium (at the same time).

59. The point at which all three curves meet is called the triple point.

60. Triple Point

61. The triple point is the only condition under which all three phases (solid, liquid, and gas) exist at the same time.

62. The normal boiling point is the temperature at which the liquid boils when the external pressure is 1 atm.
The normal melting point is the temperature at which solid melts when the external pressure is 1 atm.

63. Critical point is temperature and pressure at which liquid and gas become undistinguishable: supercritical fluid.

64. Phase Diagram for CO2

65. The carbon dioxide is stored in large tanks as liquid carbon dioxide
The carbon dioxide is stored in large tanks as liquid carbon dioxide. Assuming we lived at sea level (1 atm), how could carbon dioxide be liquefied?

66. At 1 atmosphere and room temperature (250C), would you expect solid carbon dioxide to melt to the liquid phase, or sublime to the gas phase?

67. Albuquerque is 5,500 feet above sea level, which means the normal atmospheric pressure is less than 1 atm. In Albuquerque, will water freeze at a lower temperature or a higher temperature than at 1 atmosphere?

68. At a constant temperature, what would you do to cause this substance to change from the liquid phase to the solid phase?