This document provides an overview of states of matter, gas laws, intermolecular forces, liquids, solids, and phase changes. It discusses the key properties and behaviors of gases, liquids, and solids, including gas laws, vapor pressure, boiling point, factors that affect boiling point, types of solids, and phase changes. The document also provides examples and practice problems to illustrate these concepts.

1. GASES, LIQUIDS, AND SOLIDS
CHAPTER 5
CHEM 103
• states of matter
• gas laws
• intermolecular forces
• liquids
• surface tension
• vapor pressure
• factors affecting boiling point
• solids
• phase changes

2. The states of matter

3. Which phase of matter has negligible interactions
between molecules?
1. gas
2. liquid
3. solid
4. More than one answer is correct.

4. PRESSURE
Pressure is defined as force per unit area and is expressed as:
1 atm = 760 mm Hg = 760 torr = 101,325 Pa
MERCURY BAROMETER MANOMETER

5. Gas Laws
PV = constant or P1V1 = P2V 2
BoyleBoyle’s law:’s law: for a fixed mass of gas at a constant temperature, the
volume is inversely proportional to the pressure

6. Gaseous nitrogen in an air bag has a pressure of
745 mm Hg and a volume of 65.0 L. If the gas is transferred to a
25.0 L bag at the same temperature what is the new pressure?
1. 287 mm Hg
2. 48425 mm Hg
3. 1940 mm Hg
4. 1210625 mm Hg

7. CharlesCharles’s Law:’s Law: the volume of a fixed mass of gas at a constant
pressure is directly proportional to the temperature in Kelvins (K).
V
T
V1
T1
V2
T2
= a constant or =

8. P
T
P1
T1
P2
T2
= a constant or =
Gay-LussacGay-Lussac’s Law’s Law:: for a fixed mass of gas at constant volume,
the pressure is directly proportional to the temperature in kelvins (K).

9. Boyle’s law, Charles’s law and Gay-
Lussac’s law can be combined into one law
called the combined gas lawcombined gas law..

10. A gas occupies 3.00 L at 2.00 atm. Calculate
its volume when the pressure is 10.15 atm at
the same temperature.
because the temperature is constant T1 = T2
P1 = 2.00 atm V1 = 3.00 LInitial:
Final: P2 =10.15 atm V2 = ?

11. At a temperature of 273 K
(0 °C) and a pressure of
1 atm (STP conditions), one
mol of any gas occupies a
volume of 22.4 L
22.4 L
Avogadro’s Law
Equal volumes of all gases
at the same T and P
contain the same number
of moles.

12. Ideal Gas Law
Avogadro’s law allows us to write a gas law that is valid not
only for any P, V, and T but also for any mass of gas.
Ideal gas law:Ideal gas law:
PV = nRT
P = pressure of the gas in atmospheres (atm)
V = volume of the gas in liters (L)
n = moles of the gas (mol)
T = temperature in kelvins (K)
R = ideal gas constantideal gas constant (a constant for all gases)

13. We find the value of RR by using the fact that 1.00 mol of
any gas at STP occupies 22.4 L
1.00 mol of CH4 gas occupies 20.0 L at 1.00 atm. What is
the temperature of the gas in kelvins?
PVR =
nT
= (1.00 atm)(22.4 L)
(1.00 mol)(273 K)
= 0.0821
L•atm
mol•K

14. How many moles of Nitrogen gas are in a 65 L
cylinder with a pressure of 1.09 atm at 298 K?
R= 0.082057 L atm/K mol
1. 2.
3. 4.
2.9

15. Gas Laws
DaltonDalton’s law of partial pressures:’s law of partial pressures: the total pressure, PT, of
a mixture of gases is the sum of the partial pressures of
each individual gas:
TTo a tank containing N2 at 2.0 atm and O2 at 1.0 atm we add an unknown
quantity of CO2 until the total pressure in the tank is 4.6 atm. What is the
partial pressure of CO2?

16. What is the partial pressure of argon, if the total pressure of
the gas mixture is 5.8 atm and the partial pressures of the
other gases are 2.7 atm?
1. 2.1 atm
2. 3.1 atm
3. 8.5 atm
4. 16 atm

17. Kinetic Molecular Theory
Assumptions:
 Gases consist of particles moving in random directions with various
speeds.
 Gas particles have no volume.
 Gas particles have no attraction between them.
 The average kinetic energy of gas particles is proportional to the
temperature in kelvins.
 Molecular collisions are elastic
 Molecules collide with the walls of their container; these collisions
constitute the pressure of the gas.

18. IDEAL vs REAL
• Ideal gas:Ideal gas: follows the assumptions of the kinetic molecular theory.
• Real gasesReal gases: atoms or molecules do occupy some volume and there are
forces of attraction between their atoms or molecules.
In reality there is no ideal gas, however at STP conditions most real gases
behave close enough to ideal.
Vm= real volume of 1 mol
V°m= ideal volume of 1 mol

19. Which of the following is NOT true about
real gases?
1. A gas particle will move faster as it is
heated.
2. Gas particles have no volume.
3. Molecular collisions with walls of the
container constitute the pressure of the
gas.
4. There are forces of attraction between
gas molecules.

20. Intermolecular Forces
The strength of attractive forces between molecules
determines whether any sample of matter is a gas, liquid,
or solid.
Near STP, the forces of attraction between molecules of most gases are so small that
they can be ignored. When T decreases and/or P increases, the forces of attraction
become important to the point that they cause condensation and ultimately solidification.
 London forces
 Dipole-dipole interactions
 Hydrogen bonding

21. London Dispersion Forces
London dispersion forces are the attraction
between temporary induced dipoles.

22. London Dispersion Forces
– Exist between all atoms and molecules.
– They are the only forces of attraction between
nonpolar molecules.
– This force increases as the mass and number of
electrons increases.
– Even though these forces are very weak, they
contribute significantly to the attractive forces
between large molecules.

23. Dipole-Dipole Interactions
Polar molecules contain permanent partial charges
resulting in attractions between opposite partial charges
between molecules. Dispersion forces are experienced in
addition to dipole forces.
Butane is a nonpolar molecule the
only interactions between butane
molecules are London forces.
Acetone is a polar molecule; its molecules are held
together in the liquid state by dipole-dipole
interactions.

24. Hydrogen Bonds
Force of attraction between the partial positive charge on a
hydrogen bonded to an atom of high electronegativity, most
commonly O, N, or F and the partial negative charge
(unshared electrons) on a nearby O, N or F.

25. Table 6-2, p. 171
δ−

26. Which of the following is the strongest?
1. London Dispersion Forces
2. Dipole-Dipole Interactions
3. Hydrogen Bonds
4. Covalent Bonds

27. Liquids
– As pressure increases in a real gas, its molecules come closer and
closer with the result that attractions between molecules become
important.
– Forces of attraction among gas molecules condense it into a liquid.
– In liquids, there is very little space between molecules; liquids are
difficult to compress.
– The density of liquids is much greater than that of gases because
the same mass occupies a much smaller volume.
– The position of molecules in a liquid is random and there is irregular
space between them into which other molecules can slide; this
causes liquids to be fluid.

28. Surface Tension
The layer on the surface of a liquid produced by uneven intermolecular
attractions at its surface. Surface tension is directly related to strength
of the intermolecular attraction between molecules.

29. VAPOR PRESSURE
When the equilibrium liquid  vapor is reached, the pressure produced by
the vapor on the liquid remains constant as well as the number of vapor
molecules per unit volume.
The pressure of vapor in equilibrium with a liquid is called vapor pressure
at that particular temperature. Vapor pressure is temperature dependent.
At a given temperature each liquid has a characteristic vapor pressure
depending on the strength of the intermolecular forces.
Note: As long as liquid and vapor are present the pressure exerted by the vapor is independent of
the volume of the container.

30. LIQUID-VAPOR
EQUILIBRIUM
Vaporization is the process in which a
liquid is converted to a gas (referred as
vapor). If the container is open
evaporation proceeds until all the liquid
is gone. However if the container is
closed a dynamic equilibrium is
established between the liquid and the
vapor in which the rate of condensation
equals the rate of vaporization.
VAPOR
LIQUID
BROMINE
CLOSED CONTAINER
SAFETY CONTAINER

31. BOILING POINT
The temperature at which a liquid boils
depends on the pressure above it. A
liquid boils at a temperature at which
its vapor pressure is equal to the
pressure above its surface. If this
pressure is 1 atm (760 mm Hg) the
temperature is referred to as the
normal boiling point. The normal
boiling point of water is 100 °C.
The boiling point of a liquid can be
reduced by applying vacuum.

32. If a substance has a high vapor pressure at room
temperature, it is also likely to have a _______ boiling point
because the vapor pressure of a liquid must _______ the
atmospheric pressure in order for boiling to occur.
1. high; equal
2. high; surpass
3. low; equal
4. low; be lower than

33. Boling Point
Temperature at which:
vapor pressure of a liquid = atmospheric pressure.

34. Factors Affecting Boiling Point
1) Intermolecular forces
Water (H20)
MW = 18 amu
Boiling point = 100° C
Methane (CH4)
MW = 16 amu
Boiling point = -164° C
2) Number of sites for intermolecular interaction
Methane (CH4)
MW = 16 amu
Boiling point = -164° C
Weak London
forces
Weak London forces
plus
Strong hydrogen
bonds
Hexane (C6H14)
MW = 86 amu
Boiling point = 69° C
Weak London
forces
Stronger
London
forces
The more electrons in the molecule and the larger the area = stronger London forces

35. 3) Molecular shape
Pentane (C5H12)
MW = 72 amu
Boiling point = 36.2 ° C
2,2-Dimethylpropane (C5H12)
MW = 72 amu
Boiling point = 9.5° C
The larger the area the
stronger the London forces

36. Solids
Crystallization (solidificationCrystallization (solidification): When liquids are
cooled, their molecules come close together,
attractive forces between them become strong,
random motion stops, and a solid is formed.

37. SOLIDS
Virtually all substances that are gases or liquids at 25 C and 1 atm are molecular.
Nonmolecular solids are classified in three groups (b, c, and d below).
X = nonmetal M+
= metal X-
= anion e-
= electron
DIAMOND TABLE SALT IRONSOLID I2 or WATER

38. IONIC SOLIDS
The geometry of ionic crystals is more complex than that of metals. Packing
can be visualized in terms of the cell units. Anions (larger size) are typically
located at the corners and cations (smaller size) are distributed in the holes in
between.

39. Liquid water
Solid water
MOLECULAR SOLIDS

40. Protein = chain of amino acids
Carbohydrate = chain of sugars
POLYMERIC SOLIDS

41. NETWORK SOLIDS
Diamonds have a melting point above 3500 °C. In diamonds, each carbon
atom forms single bonds with four other carbon atoms arranged tetrahedrally
around it. The three dimensional covalent bonding contributes to the diamond
unusual hardness.

42. NETWORK SOLIDS
Graphite also has a melting point above 3500 °C. In graphite carbons are
arranged in layers of hexagonal pattern. Each carbon is covalently bonded to
three others (two by single bonds and one by double bond). The forces
between adjacent layers of graphite are of the dispersion type (and weak).
Layers

43. “Soot” is a dark powdery deposit of unburned fuel residues, usually
composed mainly of amorphous carbon.
AMORPHOUS SOLIDS
Combustion pipe clogged with soot

44. “Buckyball”
Carbon structure containing 60 carbons
“Nanotube”
Carbon tubular structure
INTERESTING RECENTLY DISCOVERED FORMS OF CARBON
1996 Nobel prize

45. Phase Changes
Phase change:Phase change: a change from one
physical state (gas, liquid, or solid)
to another.
Phase:Phase: any part of a system that looks
uniform throughout: ice, water, steam.

47. Phase Diagram

48. Which of the following processes takes
the most energy?
1. Melting 50 g ice at 0°C
2. Boiling 50 g water at 100°C
3. Heating 50 g water from 0°C to 10°C
4. Heating 50 g steam from 100°C to 110°C