The document discusses the three main states of matter - gases, liquids, and solids. It describes some key physical properties for each state, including how gases behave according to the kinetic molecular theory and gas laws, how liquids have properties like viscosity and surface tension, and how solids exist in defined crystalline or amorphous structures. The interactions between particles in each state, such as van der Waals forces, hydrogen bonding, and metallic bonding, help determine characteristic properties like boiling points.

1. Chapter 5
States of Matter
Gases, Liquids, and Solids
Denniston
Topping
Caret
7th
Edition
Copyright© The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

2. Changes in State
• Changes in state are considered to be
physical changes
• During a change of physical state many
other physical properties may also change
• This chapter focuses on the important
differences in physical properties among
– Gases
– Liquids
– Solids

3. Comparison of Physical Properties
of Gases, Liquids, and Solids

4. 5.1 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
– mass
We can systematically change
one of the properties and see
the effect on the others

5. 5.1TheGaseousState Measurement of Gases
• 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. 5.1TheGaseousState Barometer
• Measures atmospheric
pressure
– Invented by Evangelista Torricelli
• Common units of pressure
– atmosphere (atm)
– torr (in Torricelli’s honor)
– pascal (Pa) (in honor of Blaise
Pascal)
• 1 atm is equal to:
– 760 mmHg
– 760 torr
– 76 cmHg

7. 5.1TheGaseousState Kinetic Molecular Theory of Gases
1. Gases are made up of small atoms or
molecules that are in constant, random
motion
2. The distance of separation is very large
compared to the size of the individual
atoms or molecules
– Gas is mostly empty space
1. All gas particles behave independently
– No attractive or repulsive forces exist
between them

8. 5.1TheGaseousState
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
4. The average kinetic energy of the atoms or
molecules increases or decreases in
proportion to absolute temperature
– As temperature goes up, particle speed goes up
Kinetic Molecular Theory of Gases

9. 5.1TheGaseousState
Kinetic Molecular Theory of Gases
Explains the following statements:
• Gases are easily compressible – gas is mostly
empty space, room for particles to be pushed
together
• Gases will expand to fill any available volume
– move freely with sufficient energy to overcome
attractive forces
• Gases have low density – being mostly empty
space; gases have low mass per unit volume

10. 5.1TheGaseousState
• Gases readily diffuse through each other – they are
in continuous motion with paths readily available due to
large space between adjacent particles
• Gases exert pressure on their containers – pressure
results from collisions of gas particles with the container
walls
• Gases behave most ideally at low pressure and
high temperature
– Low pressure, average distance of separation is
greatest, minimizing interactive forces
– High temperature, rapid motion overcomes interactive
forces more easily

11. 5.1TheGaseousState
Ideal Gases vs. Real Gases
• In reality there is no such thing as an ideal
gas
– It is a useful model to explain gas behavior
• Nonpolar gases behave more ideally than
polar gases because attractive forces are
present in polar gases

12. Gas Diffusion5.1TheGaseousState
Ammonia (17.0 g/mol)
Ammonia diffused farther in same time, lighter moves faster
Hydrogen chloride (36.5 g/mol)

13. • Boyle’s law - volume of a gas varies
inversely with the pressure exerted by the
gas if the temperature and number of moles
are held constant
• The product of pressure (P) and volume (V)
is a constant
• Used to calculate
– Volume resulting from pressure change
– Pressure resulting from volume change
PV = k1
5.1TheGaseousState Boyle’s Law
PiVi = PfVf

14. 5.1TheGaseousState Application of Boyle’s Law
• Gas occupies 10.0 L at 1.00 atm pressure
• Product, PV = (10.0 L) (1.00 atm) = k1
• Double the pressure to 2.0 atm, decreases the
volume to 5.0 L
– (2.0 atm)(Vx) = (10.0 L)(1.00 atm)
– Vx = 5.0 L

15. 5.1TheGaseousState Boyle’s Law Practice
1. A 5.0 L sample of a gas at 25o
C 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?

16. • It is possible to relate gas volume and
temperature
• Charles’s law - volume of a gas varies
directly with the absolute temperature (K) if
pressure and number of moles of gas are
constant
• Ratio of volume (V) and temperature (T) is
a constant
2k
T
V
=
f
f
i
i
T
V
T
V
=
5.1TheGaseousState Charles’s Law

17. 5.1TheGaseousState Application of Charles’s Law
• If a gas occupies 10.0 L at 273 K with
V/T = k2
• Doubling temperature to 546 K, increases
volume to 20.0 L
10.0 L / 273 K = Vf / 546 K

18. 5.1TheGaseousState Practice with Charles’s Law
1. A 2.5 L sample of gas at 25o
C is heated to
50o
C at constant pressure. Will the volume
double?
2. What would be the volume?
3. What temperature would be required to
double the volume?

19. • If a sample of gas undergoes change
involving volume, pressure, and
temperature simultaneously, use the
combined gas law
• Derived from a combination of Boyle’s law
and Charles’s law
f
ff
i
ii
T
VP
T
VP
=
5.1TheGaseousState Combined Gas Law

20. • Calculate the volume of N2 resulting when
0.100 L of the gas is heated from 300. K to
350. K at 1.00 atm
• What do we know?
– Pi = 1.00 atm Pf = 1.00 atm
– Vi = 0.100 L Vf = ? L
– Ti = 300. K Tf = 350. K
• Vf= ViTf / Ti this is valid as Pi = Pf
• Vf= (0.100 L)(350. K) / 300. K = 0.117 L
• Note the decimal point in the temperature to indicate
significance
f
ff
i
ii
T
VP
T
VP
=
5.1TheGaseousState Using the Combined Gas Law

21. 5.1TheGaseousState Practice With the Combined
Gas Law
Calculate the temperature when a 0.50 L
sample of gas at 1.0 atm and 25o
C is
compressed to 0.05 L of gas at 5.0 atm.

22. • Avogadro’s law - equal volumes of any
ideal gas contain the same number of moles
if measured under the same conditions of
temperature and pressure
• Changes in conditions can be calculated by
rewriting the equation
3k
n
V
=
f
f
i
i
n
V
n
V
=
5.1TheGaseousState Avogadro’s Law

23. 5.1TheGaseousState Using Avogadro’s Law
• If 5.50 mol of CO occupy 20.6 L, how
many liters will 16.5 mol of CO occupy at
the same temperature and pressure?
• What do we know?
– Vi = 20.6 L Vf = ? L
– ni = 5.50 mol nf = 16.5 mol
– Vf= Vinf / ni = (20.6 L)(16.5 mol)
(5.50 mol)
= 61.8 L CO

24. 5.1TheGaseousState
• Molar volume - the volume occupied by 1
mol of any gas
• STP – Standard Temperature and Pressure
– T = 273 K (or 0o
C)
– P = 1 atm
• At STP the molar volume of any gas is
22.4 L
Molar Volume of a Gas

25. 5.1TheGaseousState
Gas Densities
• Density = mass / volume
• Calculate the density of 4.00 g He
– What is the mass of 1 mol of H2? 4.00 g
DensityHe = 4.00g / 22.4L
= 0.178 g/L at STP

26. • Combining:
– Boyle’s law (relating volume and pressure)
– Charles’s law (relating volume and temperature)
– Avogadro’s law (relating volume to the number of moles)
gives the Ideal Gas Law
• R is a constant, ideal gas constant
• R = 0.0821 L.
Atm/mol.
K
If units are P in atm, V in L, n in number of moles, T in K
PV=nRT
5.1TheGaseousState The Ideal Gas Law

27. =⋅
⋅
==
atm1
K273)
Kmol
atmL
6mol(0.08201
P
nRT
V 22.4 L
5.1TheGaseousState Calculating a Molar Volume
• Demonstrate molar volume of O2 gas
at STP

28. 5.1TheGaseousState
Practice Using the Ideal Gas Law
1. What is the volume of gas occupied by
5.0 g CH4 at 25o
C and 1 atm?
2. What is the mass of N2 required to
occupy 3.0 L at 100o
C and 700 mmHg?

29. • 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
• Total pressure of our atmosphere is equal to
the sum of the pressures of N2 and O2
– (principal components of air)
Pt=p1+p2+p3+...
22 ONair ppP +=
5.1TheGaseousState Dalton’s Law of Partial Pressures

30. 5.2 The Liquid State
• Liquids are practically incompressible
– Enables brake fluid to work in your car
• Viscosity - a measure of a liquid’s
resistance to flow
– A function of both attractive forces between
molecules and molecular geometry
– Flow occurs because the molecules can easily
slide past each other
• Glycerol - example of a very viscous liquid
– Viscosity decreases with increased temperature

31. 5.2TheLiquidState Surface Tension
• 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, for example – soap

32. 5.2TheLiquidState Vapor Pressure of a Liquid
• Place water in a sealed container
– Both liquid water and water vapor will exist in
the container
• How does this happen below the boiling
point?
– Temperature is too low for boiling conversion
• Kinetic theory - liquid molecules are in continuous
motion, with their average kinetic energy directly
proportional to the Kelvin temperature

33. energy + H2O(l) → H2O(g)
5.2TheLiquidState Temperature Dependence of
Liquid Vapor Pressure
• Average molecular kinetic
energy increases as does
temperature
• Some high energy
molecules have sufficient
energy to escape from the
liquid phase
• Even at cold temperatures,
some molecules can be
converted

34. H2O(g) → H2O(l) + energy
5.2TheLiquidState Movement From Gas Back to
Liquid
• Molecules in the vapor phase can lose
energy and be converted back to the
liquid phase
• Evaporation - the process of conversion
of liquid to gas at a temperature too low
to boil
• Condensation - conversion of gas to the
liquid state

35. 5.2TheLiquidState Liquid Water in Equilibrium
With Water Vapor
• 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

36. 5.2TheLiquidState Boiling Point
• 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
• Boiling point is dependant on the intermolecular
forces
– Polar molecules have higher b.p. than nonpolar
molecules

37. 5.2TheLiquidState Van der Waals Forces
• Physical properties of liquids are explained in
terms of their intermolecular forces
• Van der Waals forces are intermolecular forces
having 2 subtypes
– Dipole-dipole interactions
– Attractive forces between polar molecules
– London forces
– As electrons are in continuous motion, a nonpolar
molecule could have an instantaneous dipole

38. 5.2TheLiquidState London Forces
• Exist between all molecules
• The only attractive force between nonpolar
atoms or 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

39. 5.2TheLiquidState Hydrogen Bonding
• 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

40. 5.2TheLiquidState
Examples of Hydrogen Bonding
• Hydrogen bonding has an extremely important
influence on the behavior of many biological
systems
• H2O
• NH3
• HF

41. 5.3 The Solid State
• Particles highly organized, in a 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

42. 5.3TheSolidState
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 solids
• held together entirely by covalent bonds
• high m.p. and b.p.
• extremely hard
• diamond

43. 5.3TheSolidState
3.Molecular 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

44. 5.3TheSolidState
Four Types of Crystalline Solids