1. THE NATURE OF GASES, LIQUIDS, SOILDS AND CHANGES OF STATES
States of Matter
THE NATURE OF
GASES, LIQUIDS, SOILDS
AND
CHANGES OF STATES

2. Kinetic Energy
Energy of motion

3. (Motion)
(or atoms)

4. Gas particles travel in straight line ,and change their direction when they collide with other molecules.
Collisions are perfectly elastic, Total Kinetic energy remains constant
Fill the containers regardless of shape and volume
Uncontained gases diffuse in the space

5. Gas Particles
Gas particles are always in motion except for the very specific condition known as Absolute Zero.

6. Gas Particles Move Randomly

7. At the same temperature, lighter gases move on average faster than heavier gases.

8. Gas Pressure
Gas pressure is a gauge of the number and force of collisions between gas particles and the walls of the container that holds them.
The SI unit for pressure is the pascal (Pa),
atmospheres (atm),
millimeters of mercury (mmHg),
and torr.
1 atm= 760 mm Hg = kPa

9. Gas Pressure 760 mmHg…. 760 torr….. 1.00 atm…. 101,325 Pa….
kPa…..
Is the typical atmospheric pressure at sea level.
On the top of Mount Everest the atmospheric pressure : ~ 33.7 KPa

10. Sample Problem 10-1 A gas sample is at a pressure of 1.50 atm.
Convert this to kilopascals
mm of Hg
1.50 atm x kPa =152 kPa
1 atm
x 760 mm Hg =1140 mm Hg
1.50 atm

11. Gas Pressure (P)
Pressure refers to the force the gas produces on the walls of the container that it occupies.
The phenomenon of pressure is really a force applied over a surface area.

12. Barometer
A barometer uses the height of a column of mercury to measure gas pressure in millimeters of mercury or torr (1 mmHg = 1 torr). The mercury is pushed up the tube from the dish until the pressure at the bottom of the tube (due to the mass of the mercury) is balanced by the atmospheric pressure.

13. Barometer: Gas Pressure
1 atm is the pressure required to support 760 mm Hg in a barrometer at 25 deg C = kPa

14. Manometers

15. Diffusion

17. Effusion
The escape of a gas through a tiny pore or pinhole in its container is called EFFUSION.

18. Kinetic Energy and Kelvin Temperature
According to the kinetic theory of gases, at absolute zero the molecules of a gas would not move.
More advanced theories show that even at 0 K a very slight movement will persist.

19. Kinetic Energy and Kelvin Temperature
When the substance is heated, the particles of the substance absorb energy and stored it as Potential Energy
The remaining absorbed energy speeds up the particles ie. Increases the average kinetic energy
This results in increase in temperature

21. Increase in average KE of the particles causes the temperature to rise
At lower temperature the particles have lower KE
At the absolute zero ( degree C) the motion of the particles theoretically ceases.
Kelvin temperature of the substance is directly proportional to the KE
At any given temp. the particles of all substances regardless of physical state have same KE

22. Many molecules have intermediate Kinetic energy
Boltzman Distribution Curve
Many molecules have intermediate Kinetic energy
Percent of molecules
Few molecules have high Kinetic Energy
Kinetic energy

23. Normal Boiling Point
The temperature at which the vapor pressure of the liquid is just equal to the external pressure
Changing external pressure changes boiling point
Low pressure = Lower Boiling Point
High pressure= Higher Boiling Point
Ie. Water boiling in Denver

24. Boiling

28. A Model Of Liquids Liquids are incompressible.
Liquids maintain their volumes regardless of the sizes and shapes of the containers in which they are kept.
Adopt the shape of the containers in which they are kept.
Liquids diffuse very slowly.
Liquids evaporate from open containers.

29. Liquids can diffuse slowly

30. Liquid particles are in motion/ slide past one another/ flow
Presence of intermolecular forces
Higher intermolecular forces = higher boiling point
Intermolecular forces reduce the space between the particles resulting in more denser liquids than gases
Particles vibrate and spin as they move / have low KE / can not escape into gaseous state (only few can escape)

31. Evaporation
The molecules of a liquid do not all have the same kinetic energy.
At a higher temperature, there will be a number of molecules which have a kinetic energy, E, which high enough to overcome the attractive forces between the molecules of the liquid. These molecules will escape from the surface of the liquid as vapor, a process known as evaporation.
In the process, the liquid cools, and heat from the surroundings has to be supplied in order to maintain the evaporation.

32. Evaporation / KE

33. Evaporation

34. Evaporation /Condensation
Forming the bonds
Breaking the bonds

35. Vapor and liquid in equilibrium
n container A, the liquid is evaporating.  Some of the molecules have enough kinetic energy to escape (turn to a gas) by pushing against the pressure of the atmosphere.  Container B shows the flask is saturated.  When new molecules of liquid are vaporized, the gas cannot hold additional molecules, therefore some of the molecules condense back to liquid.

36. Vapor Pressure & Boiling Point
 Vapor pressure is the pressure exerted by a liquid in equilibrium with its pure liquid phase at a given temperature.
The vapor pressure of a liquid is dependent only upon the nature of the liquid and the temperature.
Different liquids at any temperature have different vapor pressures.
The vapor pressure of every liquid increases as the temperature is raised.
The normal boiling point of any pure substance is the temperature at which the vapor pressure of that substance is equal to 1 atmosphere (760 mm Hg).

37. Water

38. Surface Area: the surface area of the solid or liquid in contact with the gas has no effect on the vapor pressure.

39. The rate of boiling is limited by the rate of heat transfer into the liquid.
* Evaporation takes place more slowly than boiling at any temperature between the melting point and boiling point, and only from the surface, and results in the liquid becoming cooler due to loss of higher kinetic energy particles.

40. Condensing (gas to liquid)
On cooling, gas particles lose kinetic energy and eventually become attracted together to form a liquid.
There is an increase in order as the particles are much closer together and can form clumps of molecules.
The process requires heat to be lost to the surroundings ie heat given out, so condensation is exothermic (ΔH -ve).
This is why steam has such a scalding effect, its not just hot, but you get extra heat transfer to your skin due to the exothermic condensation on your surface!

43. Solid fixed volume and shape at a particular temperature
greatest forces of attraction are between the particles
pack together as tightly as possible
and ordered arrangement.

44. Solid particles vibrate about their position in the structure.
With increase in temperature, the particles vibrate faster

45. Solid
Solids have the greatest density.
Solids cannot flow freely
Do not take shape of the container
fixed surface and volume
extremely difficult to compress

46. Diffusion is almost impossible in solids
Solids will expand a little on heating but nothing like as much as liquids because of the greater particle attraction restricting the expansion and contraction occurs on cooling.
The expansion is caused by the increased energy of particle vibration, forcing them further apart causing an increase in volume and corresponding decrease in density.
Diffusion is almost impossible in solids

47. Solid

48. Melting
solid is heated the particles vibrate more strongly as they gain kinetic energy and the particle attractive forces are weakened.
melting point
The particles become free to move around and lose their ordered arrangement.
Energy is needed to overcome the attractive forces.
So heat is taken in from the surroundings and melting is an endothermic process (ΔH +ve).

49. Freezing On cooling, liquid particles lose kinetic energy.
freezing point the forces of attraction are sufficient to remove any remaining freedom
and the particles come together to form the ordered solid arrangement.
freezing is an exothermic process (ΔH -ve).

57. Amorphous solids Lack ordered internal arrangement
Plastic, rubber, asphalt, glasses
Atoms are randomly arranged
Glasses: super-cooled liquid, soften when heated
Inorganic compounds cooled to a rigid state without crystallizing
Intermediate between crystalline and free flowing liquids

58. Glass- blowing
When glass breaks, it has irregular angles and edges

59. Glass

60. Changes of State

61. Cooling curve: Changes of State
A cooling curve summarizes the changes:
gas ==> liquid ==> solid
Cooling curve:
the temperature stays constant during the state changes of condensing Tc and freezing Tf.

62. Cooling Curve all the energy removed on cooling at these temperatures
weakens the inter-particle forces without temperature fall.

64. Heating Curve
A graph of the temperature of 1 kg of water, originally ice at -50° C, as heat is added to it.

66. Heating curve:
the temperature stays constant during the state changes of melting at Tm and boiling at Tb.
This is because all the energy absorbed in heating at these temperatures goes into weakening the inter-particle forces without temperature rise.
solid ==> liquid ==> gas

67. Phases

68. Phase Diagrams : graph of solid, liquid, gas
The stability of solid, liquid and gas phases depends on the temperature and the pressure.
The three phases are in equilibrium at the triple point.
The gas and liquid phases are separated by a phase transition only below the temperature of the critical point.

69. Phase Diagram

70. Water Triple point =0.016 deg C 0.61 kPa (0.0060 atm)
Decrease in P ; lowers the bp
raises the mp

71. Sublimation

72. Sublimation:
This is when a solid, on heating, directly changes into a gas without melting,
AND the gas on cooling re-forms a solid directly without condensing to a liquid.
usually involve just a physical change BUT its not always that simple!

73. Sublimation
Freeze-dried coffee: freezing freshly brewed coffee and removing water vapor with vacuum pump
Dry ice for frozen food
Solid air freshener
Mothballs
Separation and purification of organic compounds