The document discusses different states of matter and intermolecular forces. It explains that gases have weak interactions between molecules while liquids and solids have strong interactions. Solids exist in crystalline or amorphous forms, with crystalline solids having a highly ordered particle arrangement. Covalent network solids like diamonds have strong covalent bonds between all atoms while molecular solids like graphite are held together by van der Waals forces. Ionic solids consist of positive and negative ions arranged in a crystal lattice. Metals have delocalized valence electrons that provide metallic bonding throughout the solid.

1. Slide 1
Intermolecular forces are responsible for the different
phases of matter [gas, liquid, solid] –
• gases have only weak interactions between molecules,
• liquids and solids have strong interactions.
States of Matter

2. Slide 2
States of Matter
The fundamental difference between states of matter is
the distance between particles.

3. Slide 3
States of Matter
Because in the solid and liquid states particles are closer
together, we refer to them as condensed phases.

4. Slide 4
The States of Matter
• The state a substance is in at a
particular temperature and
pressure depends on opposing
factors:
The kinetic energy of the
particles
The strength of the attractions
between the particles

5. Slide 5
Hydrogen Bonding
• The dipole-dipole interactions experienced
when H is bonded to N, O, or F are unusually
strong.
• We call these interactions hydrogen bonds.

6. Slide 6
Hydrogen Bonding
Hydrogen Bond:
Molecules in which H is bound to the very electronegative
O, N, or F.

7. Slide 7
Hydrogen Bonding
CH4 110 K
SiH4 160 K
GeH4 175 K
SnH4 215 K
H2O 373 K
H2S 215 K
H2Se 225 K
H2Te 270 K
Effect of Hydrogen Bonding on Boiling Point:Effect of Hydrogen Bonding on Boiling Point:

8. Slide 8
Hydrogen Bonding
• The nonpolar series (SnH4
to CH4) follow the
expected trend.
• The polar series follows
the trend from H2Te
through H2S, but water is
quite an anomaly.

9. Slide 9
Intermolecular Forces
Application of HydrogenApplication of Hydrogen
Bonding!!Bonding!!
Cellulose molecules are
present in the trunk of the
tree. They form strong
hydrogen bonds between O
and H.

10. Slide 10
Phase Changes

11. Slide 11
Phase Changes - Entropy

12. Slide 12
Phase Changes
• Molar Heat of Fusion (∆Hfus):
The energy required to melt one mole of solid (kJ).
• Molar Heat of Vaporization (∆Hvap):
The energy (kJ) required to vaporize one mole of liquid.
• Molar Heat of Sublimation (∆Hsub):
The energy (kJ) required to sublime one mole of solid.
∆Hsub=∆ Hfus + ∆Hvap

13. Slide 13
Phase Changes

14. Slide 14
Gases
Which elements exist normally as gases?

15. Slide 15
Gases
Gases are –
• compressible
• expandable
• form homogeneous mixtures
We can use a model in which –
• molecules move rapidly
• respond quickly to changes in volume
• the volume of the molecules is very small (0.1%)
compared with the total volume occupied by a gas

16. Slide 16
Gases and Pressure
The pressure of the atmosphere …

17. Slide 17
Gases and Pressure
The units of pressure that you need to know:
* atmosphere (atm)
* Pa (N/m2
, 101,325 Pa = 1 atm)
* Torr (760 Torr = 1 atm)
* bar (1.01325 bar = 1 atm)
* mm Hg (760 mm Hg = 1 atm)
* [lb/in2
(14.696 lb/in2
= 1 atm)]

18. Slide 18
The Ideal Gas Law
Starting with the Ideal Gas Law we can work back to a
series of extremely important relationships that describe
the behaviour of a gas under different conditions –
• Boyle’s Law
• Charles’ Law
• Avogadro’s Principle …

19. Slide 19
Boyle’s Law
• Pressure–Volume Law (Boyle’s Law):
The volume of a fixed
amount of gas maintained
at constant temperature is
inversely proportional to the
gas pressure –
P1V1 = P2V2
Pressure
1
Volume ∝

20. Slide 20
Charles’ Law
Temperature–Volume Law (Charles’ Law):
V ∝ T
The volume of a fixed
amount of gas at constant
pressure is directly
proportional to the Kelvin
temperature of the gas –
V1 = V2
T1 T2

21. Slide 21
Avogadro’s Principle
The Volume–Amount Law (Avogadro’s Principle):
V ∝ n
At constant pressure and
temperature, the volume of
a gas is directly
proportional to the number
of moles of gas present –
V1 = V2
n1 n2

22. Slide 22
The Ideal Gas Law
• Ideal gases obey the following equation -
• The gas constant R = 0.08206 L·atm·K–1
·mol–1
or …. = 8.31451 J.K-1
.mol-1
TRnVP ⋅⋅=⋅
when pressure
is in Pascals

23. Slide 23
Ideal Gases
Standard Temperature and Pressure (STP):
“1 mole of an ideal gas occupies 22.414 L at STP”
STP conditions are –
273.15 K (0 o
C) and 1 atm pressure

24. Slide 24
Ideal Gases
The molar volumes of real gases do differ from 22.41 L …
but not by that much -

25. Slide 25
Vapor Pressure
• At any temperature, some molecules in a liquid have
enough energy to escape.
• As the temperature rises, the fraction of molecules that
have enough energy to escape increases.

26. Slide 26
Vapor Pressure
As more molecules
escape the liquid, the
pressure they exert
increases.

27. Slide 27
Vapor Pressure
The liquid and vapor
reach a state of dynamic
equilibrium: liquid
molecules evaporate and
vapor molecules
condense at the same
rate.

28. Slide 28
Vapor Pressure
• The boiling point of a liquid
is the temperature at
which its vapor pressure
equals atmospheric
pressure.
• The normal boiling point is
the temperature at which
its vapor pressure is 760
torr.

29. Slide 29
Gas Stoichiometry
Problem:
Carbonate bearing rocks (like limestone, CaCO3) react with
dilute acids such as HCl to produce carbon dioxide -
CaCO3(s) + 2 HCl(aq) CaCl2(aq) + CO2(g) + H2O(l)
How many grams of CO2 are produced from complete
reaction of 33.7 g of limestone? What is the volume of the
CO2 at RTP?

30. Slide 30
Gas Stoichiometry
Problem:
Assuming no change in temperature and pressure, calculate
the volume of O2 (in liters) required for the complete
combustion of 14.9 dm3
of butane (C4H10):
2 C4H10(g) + 13 O2(g) → 8 CO2(g) + 10 H2O(l)

31. Slide 31
Gas Stoichiometry
Problem:
Hydrogen gas, H2, can be prepared by allowing zinc metal to
react with aqueous HCl. How many dm3
of H2 can be
prepared at 742 mm Hg and 15o
C if 25.5 g of zinc (Mr =
65.4 g/mol) is allowed to react?
Zn(s) + 2 HCl(aq) → H2(g) + ZnCl2(aq)

32. Slide 32
Crystalline Silica (SiO2)
« Quartz »
Amorphous Silica [glass]

33. Slide 33
The surface of Nickel
[scanning tunneling microscopy]
Galena [PbS]
Crystalline solids …

34. Slide 34
Ionic Solids …

35. Slide 35
Sodium chloride … “rock salt”
• Na+
ions are the smaller spheres (102 pm) that sit in the “holes” between
the larger Cl-
ions (181 pm)
• the Cl-
ions form a “loose” cubic structure
• counting up the number of Na+
ions and Cl-
ions in the unit cell
gives us 4 of each (ions on a corner count for 1/8 ; ions on a face for ½ ;
ions on an edge for ¼ ; and the ion in the center for 1)
• other examples of the rock salt structure are KBr, RbI, MgO, CaO, and AgCl

36. Slide 36
Solids
Crystalline—particles are
in a highly ordered
arrangement.
Silicon dioxide

37. Slide 37
Covalent-Network and
Molecular Solids
• Diamonds are an example of a covalent-network solid in
which atoms are covalently bonded to each other.
They tend to be hard and have high melting points.

38. Slide 38
Covalent-Network and
Molecular Solids
• Graphite is an example of a molecular solid in which
atoms are held together with van der Waals forces.
They tend to be softer and have lower melting points.

39. Slide 39
Metallic Solids
• Metals are not covalently
bonded, but the attractions
between atoms are too strong
to be van der Waals forces.
• In metals, valence electrons
are delocalized throughout the
solid.

40. Slide 40