1. Conceptual Chemistry
Unit 6 – States of Matter

2. Objective 1
Describe, at the molecular level, the difference between a gas, liquid, and solid phase.

3. Solids Definite shape Definite volume
Particles are vibrating and packed close together.
The particles do not
flow.

4. Crystalline Solids Particles are arranged in an organized pattern.
Example: Diamond

5. Amorphous Solids Particles are not organized in an orderly fashion.
Example: Glass

6. Liquids Indefinite shape Definite volume
Liquids will take the shape of a container, but they maintain the same volume.
Particles are touching and packed close together.
The higher energy allows
the particles to move
around each other.

7. Viscosity
A liquid’s resistance to flow.

8. Gases Indefinite shape Indefinite volume
Gases take the shape of a container. They also occupy the volume of the container no matter how big or small it is.
High energy motion

9. Plasma High energy matter A common example is the sun.
Super high energy gas particles that lost electrons.
Plasma is the most common form of matter in the Universe.

10. States of Matter Property Solid (s) Liquid (l) Gas (g)
Particle Spacing
Close
Great
Energy
Low
Medium
High
Motion
Shape
Definite
Indefinite
Volume

11. Objective 2
Describe states of matter using the kinetic molecular theory.

12. Kinetic Molecular Theory
can explain the behavior of matter in its different states.
Kinetic Molecular Theory: Explains the states of matter based on the concept that the particles in all forms of matter are in constant motion.
Kinetic Energy: Energy an object has due to its motion.

13. Kinetic Energy and Kelvin Temperature
Temperature: the average kinetic energy of the particles in a material
As particles are heated, they absorb energy, thus increasing their average kinetic energy and their temperature.
Motion stops at absolute zero (0 Kelvin).
Kelvin temperature scale reflects
the relationship between temperature
and average kinetic energy. It is
directly proportional.

14. Objective 3
Describe changes in states of matter with respect to kinetic energy and temperature.

15. Energy and Phase Changes
During a phase change, all energy goes to motion until phase change is done.
The temperature does not change until the phase change is done.

16. Melting
Solid  Liquid
Example 1
Example 2

17. Freezing
Liquid  Solid
Example 1

18. Evaporation/Boiling
Liquid  Gas
Example 1

19. Condensation
Gas  Liquid
Example

20. Sublimation
Solid  Gas
Example
Opposite of Sublimation? Deposition

21. Objective 4
Describe the different variables that define a gas.

22. Kinetic Theory of Gases
Gases are mostly empty space.
The molecules in a gas are separate, very small, and very far apart.

23. Kinetic Theory of Gases
Gas molecules are in constant, chaotic motion.
Collisions between gas molecules are elastic (there is no energy gain or loss).

24. Kinetic Theory of Gases
The average kinetic energy of gas molecules is directly proportional to the absolute temperature.
Gas pressure is caused by collisions of molecules with the walls of the container.

25. Gases doing all of these things!
Behavior of Gases
Gases have weight.
Gases take up space.
Gases exert pressure.
Gases fill their containers.
Gases doing all of these things!

26. Variables that Describe a Gas
Volume: measured in L, mL, cm3 (1 mL = 1 cm3)
Amount: measured in moles (mol), grams (g)
Temperature: measured in Kelvin (K)
K = ºC + 273
Pressure: measured in mm Hg, torr, atm, etc.
 P = F / A (force per unit area)

27. P = F /A Moderate Force (about 100 lbs) Small Area (0.0625 in2)
Enormous Pressure (1600 psi)

28. Large Surface Area (lots of nails)
Bed of Nails
Moderate Force
Small Pressure
P = F / A
Large Surface Area (lots of nails)

29. Units of Pressure 1 atm = 760 mm Hg 1 atm = 760 torr
1 atm = x 105 Pa
1 atm = kPa

30. Boyle’s Law As P, V and vice versa…. Inverse relationship
For a given number of molecules of gas at a constant temperature, the volume of the gas varies inversely with the pressure.
As P, V and vice versa….
Inverse relationship
P1V1 = P2V2

31. Boyle’s Law and Kinetic Molecular Theory
How does kinetic molecular theory explain Boyle’s Law?
Gas molecules are in constant, random motion.
Gas pressure is the result of molecules colliding with the walls of the container.
As the volume of a container becomes smaller, the collisions over a particular area of container wall increase…the gas pressure increases!

33. Pressure-Volume Calculations
Example: Consider the syringe. Initially, the gas occupies a volume of 8 mL and exerts a pressure of 1 atm.
What would the pressure of the gas become if its volume were increased to 10 mL?

34. Equation for Boyle’s Law
P1V1 = P2V2
where: P1 = initial pressure
V1 = initial volume
P2 = final pressure
V2 = final volume

35. P1V1 = P2V2 Using the same syringe example, just “plug in” the values:
(1 atm) (8 mL) = (P2) (10 mL)

36. P1V1 = P2V2
(1 atm) (8 mL) = P2
(10 mL)
P2 = 0.8 atm

37. P1V1 = P2V2 (1.2 atm)(12 L) = (3.6 atm)V2 V2 = 4.0 L
Example: A sample of gas occupies 12 L under a pressure of 1.2 atm. What would its volume be if the pressure were increased to 3.6 atm? (assume temp is constant)
P1V1 = P2V2
(1.2 atm)(12 L) = (3.6 atm)V2
V2 = 4.0 L

38. P1V1 = P2V2 (200 kPa)(28 L) = (P2)(17 L) P2 = 329 kPa
Example: A sample of gas occupies 28 L under a pressure of 200 kPa. If the volume is decreased to 17 L, what be the new pressure? (assume temp is constant)
P1V1 = P2V2
(200 kPa)(28 L) = (P2)(17 L)
P2 = 329 kPa

39. Temperature – Volume Relationships
What happens to matter when it is heated?
It EXPANDS.
What happens to matter when it is cooled?
It CONTRACTS.
Gas samples expand and shrink to a much greater extent than either solids or liquids.

40. Charles’ Law As T , V  and vice versa…. Direct relationship
The volume of a given number of molecules
is directly proportional to the
Kelvin temperature.
As T , V  and vice versa….
Direct relationship
Video Clip 1,
Clip 2

41. Temperature – Volume Relationship
Doubling the Kelvin temperature of a gas doubles its volume.
Reducing the Kelvin temperature by one half causes the gas volume to decrease by one half…
WHY KELVIN?
The Kelvin scale never reaches “zero” or has negative values.

42. Converting Kelvin To convert from Celsius to Kelvin: add 273.
Example: What is 110 ºC in Kelvin?
110 ºC = 383 K

43. Converting Kelvin To convert from Kelvin to Celsius: subtract 273.
Example: 555 K in Celsius?
555 K = 282 ºC

44. T2 = 746 K V1 = 117 mL; T1 = 100 + 273 = 373 K V2 = 234 mL; T2 = ???
Example: A sample of nitrogen gas occupies 117 mL at 100.°C. At what temperature would it occupy 234 mL if the pressure does not change?
V1 = 117 mL; T1 = = 373 K
V2 = 234 mL; T2 = ???
V1 / T1= V2 / T2
T2 = 746 K

45. Example: A sample of oxygen gas occupies 65 mL at 28. 8°C
Example: A sample of oxygen gas occupies 65 mL at 28.8°C. If the temperature is raised to 72.2°C, what will the new volume of the gas?
V1 = 65 mL; T1 = = K
V2 = ??? mL; T2 = = K
V1 / T1= V2 / T2
V2 = 74.3 mL

46. Temperature – Pressure Relationships
Picture a closed, rigid container of gas (such as a scuba tank) – the volume is CONSTANT.
What would happen to the
kinetic energy of the gas
molecules in the container if
you were to heat it up?
How would this affect pressure?
States of Matter Interactive
Egg in a Bottle:!
Video Clip

47. Temperature – Pressure Relationships
Raising the Kelvin temperature of the gas will cause an INCREASE in the gas pressure.
WHY?
With increasing temperature, the K.E. of the gas particles increases – they move faster!
They collide more often and with more energy with the walls of the container.