1. States of Matter: Gases, Liquids, and Solids
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Chapter 6
States of Matter:
Gases, Liquids, and Solids
Denniston Topping Caret
4th Edition

2. 6.1 The Gaseous State Ideal Gas Concept
Ideal gas - a model of the way that particles of a gas behave at the microscopic level.
We can measure the following of a gas:
temperature,
volume,
pressure and
quantity (mass)
We can systematically change one of the properties and see the effect on each of the others.

3. 6.1 The Gaseous State Measurement of Gases
The most important gas laws involve the relationship between
number of moles (n) of gas
volume (V)
temperature (T)
pressure (P)
Pressure - force per unit area.
Gas pressure is a result of force exerted by the collision of particles with the walls of the container.
6.1 The Gaseous State

4. 6.1 The Gaseous State Barometer - measures atmospheric pressure.
Invented by Evangelista Torricelli
A commonly used unit of pressure is the atmosphere (atm).
1 atm is equal to:
760 mmHg
760 torr
76 cmHg
6.1 The Gaseous State

5. 6.1 The Gaseous State Boyle’s Law 1
Boyle’s Law - volume of a gas is inversely proportional to pressure if the temperature and number of moles is held constant.
6.1 The Gaseous State
PV = k1 or
PiVi = PfVf

6. 6.1 The Gaseous State

7. Examples using Boyle's Law2
1. A 5.0 L sample of a gas at 25oC and 3.0 atm is compressed at constant temperature to a volume of 1.0 L. What is the new pressure?
2. A 3.5 L sample of a gas at 1.0 atm is expanded at constant temperature until the pressure is 0.10 atm. What is the volume of the gas?
6.1 The Gaseous State

8. 6.1 The Gaseous State 1 Charles’ Law
Charles’ Law - volume of a gas varies directly with the absolute temperature (K) if pressure and number of moles of gas are constant.
6.1 The Gaseous State or

9. 6.1 The Gaseous State

10. Examples Using Charles' Law2
1. A 2.5 L sample of gas at 25oC is heated to 50oC at constant pressure. Will the volume double?
2. What would be the volume in question 1?
3. What temperature would be required to double the volume in question 1?
6.1 The Gaseous State

11. 6.1 The Gaseous State Combined Gas Law 1
This law is used when a sample of gas undergoes change involving volume, pressure, and temperature simultaneously.
6.1 The Gaseous State

12. Example Using the Combined
Law 2
Calculate the temperature when a 0.50 L sample of gas at 1.0 atm and 25oC is compressed to 0.05 L of gas at 5.0 atm.
6.1 The Gaseous State

13. 6.1 The Gaseous State Avogadro’s Law 1
Avogadro’s Law - equal volumes of an ideal gas contain the same number of moles if measured under the same conditions of temperature and pressure.
6.1 The Gaseous State or

14. Example Using Avogaro's Law2
Assuming no change in temperature and pressure, how many moles of gas would be needed to double the volume occupied by 0.50 moles of gas?
6.1 The Gaseous State

15. 6.1 The Gaseous State Molar Volume of a Gas
Molar Volume - the volume occupied by 1 mol of any gas.
STP - Standard Temperature and Pressure
T = 273 K (or 0oC)
P = 1 atm
At STP the molar volume of a gas is 22.4 L
We will learn to calculate the volume later.
6.1 The Gaseous State

16. 6.1 The Gaseous State Gas Densities We know: density = mass/volume
Let’s calculate the density of H2.
What is the mass of 1 mol of H2?
2.0 g
What is the volume of 1 mol of H2?
22.4 L
Density = 2.0 g/22.4 L = g/L
6.1 The Gaseous State

17. 6.1 The Gaseous State 1 The Ideal Gas Law
Combining Boyle’s Law, Charles’ Law and Avogadro’s Law gives the Ideal Gas Law.
6.1 The Gaseous State
PV=nRT
R (ideal gas constant) = L.Atm/mol.K

18. Let’s calculate the molar volume at STP using the ideal gas law:
PV = nRT
What would be the pressure?
1 atm
What would be the temperature?
273 K
What would be the number of moles?
1 mol 2
6.1 The Gaseous State
22.4 L

19. 6.1 The Gaseous State Examples using The Ideal Gas Law
1. What is the volume of gas occupied by 5.0 g CH4 at 25oC and 1 atm?
2. What is the mass of N2 required to occupy 3.0 L at 100oC and 700 mmHg?
6.1 The Gaseous State

20. 6.1 The Gaseous State 1 Dalton’s Law of Partial Pressures
Dalton’s Law - a mixture of gases exerts a pressure that is the sum of the pressures that each gas would exert if it were present alone under the same conditions.
6.1 The Gaseous State
Pt=p1+p2+p3+...
For example, the total pressure of our atmosphere is equal to the sum of the pressures of N2 and O2.

21. 6.1 The Gaseous State 3 Kinetic Molecular Theory of Gases
Provides an explanation of the behavior of gases that we have studied in this chapter. Summary follows:
1. Gases are made up of small atoms or molecules that are in constant and random motion.
2. The distance of separation is very large compared to the size of the atoms or molecules.
the gas is mostly empty space.
6.1 The Gaseous State

22. 6.1 The Gaseous State 3. All gas particles behave independently.
No attractive or repulsive forces exist between them.
4. Gas particles collide with each other and with the walls of the container without losing energy.
The energy is transferred from one atom or molecule to another.
5. The average kinetic energy of the atoms or molecules is proportional to absolute temperature.
K.E. = 1/2mv2 so as temperature goes up, the speed of the particles goes up.
6.1 The Gaseous State

23. 4
How does the Kinetic Molecular Theory of Gases explain the following statements?
Gases are easily compressible.
Gases will expand to fill any available volume.
Gases have low density.
6.1 The Gaseous State

24. Remember: pressure is a force per unit area resulting from collision of gas particles with the walls of the container.
If pressure remains constant why does volume increase with temperature?
6.1 The Gaseous State

25. Gases behave most ideally at low pressure and high temperatures.
Ideal Gases Vs. Real Gases
6.1 The Gaseous State
In reality there is no such thing as an ideal gas. Instead this is a useful model to explain gas behavior.
Non-polar gases behave more ideally than polar gases because attractive forces are present in polar gases.

26. 6.2 The Liquid State 5 Liquids are practically incompressible.
Enables brake fluid to work in your car
Viscosity - a measure of a liquids resistance to flow.
Flow occurs because the molecules can easily slide past each other.
Glycerol - example of a very viscous liquid.
Viscosity decreases with increased temperature.

27. Surface tension - a measure of the attractive forces exerted among molecules at the surface of a liquid.
Surface molecules are surrounded and attracted by fewer liquid molecules than those below.
Net attractive forces on surface molecules pull them downward.
Results in “beading”
Surfactant - substance added which decreases the surface tension
example: soap
6.2 The Liquid State

28. 6.2 The Liquid State Vapor Pressure of a Liquid 6
What happens when you put water in a sealed container?
Both liquid water and water vapor will exist in the container.
How does this happen below the boiling point?
Kinetic Theory - Liquid molecules are in continuous motion, with their average kinetic energy directly proportional to the Kelvin temperature.
6.2 The Liquid State

29. The green line represents the minimum energy required to break the intermolecular attractions.
Even at the cold temp, some molecules can be converted.
6.2 The Liquid State
energy + H2O(l)  H2O(g)

30. Once there are molecules in the vapor phase, they can be converted back to the liquid phase
H2O(g)  H2O(l) + energy
6.2 The Liquid State
evaporation - the process of conversion of liquid to gas, at a temperature too low to boil
condensation - conversion of the gas to the liquid state.

31. 6.2 The Liquid State
When the rate of evaporation equals the rate of condensation, the system is at equilibrium.
Vapor pressure of a liquid - the pressure exerted by the vapor at equilibrium
H2O(g) H2O(l)

32. Boiling point - the temperature at which the vapor pressure of the liquid becomes equal to the atmospheric pressure.
Normal boiling point - temperature at which the vapor pressure of the liquid is equal to 1 atm.
What happens when you go to a mountain where the atmospheric pressure is lower than 1 atm?
The boiling point lowers.
6.2 The Liquid State

33. Boiling point is dependant on the intermolecular forces
Polar molecules have higher b.p. than nonpolar molecules.
6.2 The Liquid State

34. 6.2 The Liquid State Van der Waals Forces 7
Van der Waals Forces are types of intermolecular forces.
Consists of:
Dipole-dipole interactions (section 4.5)
London forces
6.2 The Liquid State

35. 6.2 The Liquid State London forces: Electrons are in constant motion.
exist between all molecules,
is the only attractive force between nonpolar atoms or molecules,
and dipole-dipole attractions occur between polar molecules.
Electrons are in constant motion.
Electrons can be, in an instant, arranged in such a way that they have a dipole. (Instantaneous dipole)
The temporary dipole interacts with other temporary dipoles to cause attraction.
6.2 The Liquid State

36. 6.2 The Liquid State Hydrogen Bonding 8 Hydrogen bonding:
not considered a Van der Waals Force
is a special type of dipole-dipole attraction
is a very strong intermolecular attraction causing higher than expected b.p. and m.p.
Requirement for hydrogen bonding:
molecules have hydrogen directly bonded to O, N, or F
6.2 The Liquid State

37. Examples of hydrogen bonding:
H2O
NH3 HF
6.2 The Liquid State

38. 6.3 The Solid State 9
Particles highly organized and well defined fashion
Fixed shape and volume
Properties of Solids:
incompressible
m.p. depends on strength of attractive force between particles
Crystalline solid - regular repeating structure
Amorphous solid - no organized structure.

39. 6.3 The Solid State Types of Crystalline Solids 1. Ionic Solids
held together by electrostatic forces
high m.p. and b.p.
hard and brittle
if dissolves in water, electrolytes
NaCl
2. Covalent Solid
Held together entirely by covalent bonds
extremely hard
Diamond
6.3 The Solid State

40. 6.3 The Solid State 3. Molecular solids 4. Metallic solids
molecules are held together with intermolecular forces
often soft
low m.p.
often volatile
ice
4. Metallic solids
metal atoms held together with metal bonds
metal bonds
overlap of orbitals of metal atoms
overlap causes regions of high electron density where electrons are extremely mobile - conducts electricity
6.3 The Solid State

41. 6.3 The Solid State

42. The End
Chapter 6