States of Matter
Solid, Liquid or Gas??

The state of a pure substance (solid, liquid or
gas) depends on the strength of the attractive
forces between its particles.
In Solid State: The attractive forces are relatively strong and each particle remains
bonded to its adjacent particles.
As the solid is heated up, the average kinetic
energy (temperature) of its particles increases.
 Eventually, the particles have enough kinetic
energy (motion) to break the bonds with
neighbouring particles.
 At this particular temperature (its melting point),
the particles transition into liquid form.


At the melting point, all added energy is used to
disrupt the intermolecular forces between the
solid particles. This involves a change in
potential energy.
During this phase change (known as fusion), there will be no change in temperature
until the entire sample has melted.
In Liquid State: The particles can move past each other but they are constantly
breaking and forming bonds with neighbouring particles.
 As the liquid is heated up, their kinetic energy
increases until the particles have enough energy to break
all the intermolecular bonds and remain free of other
particles.
 At this particular temperature (its boiling point), the
particles transition into a gas.


At the boiling point, all added energy is used to
disrupt all the intermolecular forces between the
liquid particles. This involves a change in potential
energy.
During this phase change (known as vaporization), there will be no change in
temperature until the entire sample has transformed into a gas.

The particles in a solid are rigidly attached to one another and often form a
symmetrical structure. In a liquid, the particles are always in contact but can readily
move past one another. In a gas, the only contacts between particles are collisions.
The particles remain apart from one another after the collisions.
Melting and Boiling Points
When metals melt or boil,
metallic bonds must be broken.
Metals
When ionic compounds melt
or boil, ionic bonds must be
broken.
Ionic Compounds
When molecular compounds melt or
boil, intermolecular forces must be
overcome. The covalent bonds
between atoms within a molecule don’t
break when the compound melts or
boils.
Molecular Compounds
Substance
Melting
Point
(oC)
Boiling
Point
(oC)
Substance
Melting
Point
(oC)
Boiling
Point
(oC)
Substance
Melting
Point (oC)
Boiling
Point (oC)
Li(s)
+181
+1342
CsBr(s)
+636
+1300
H2(g)
-259
-253
Sn(s)
+232
+2602
NaI(s)
+661
+1304
Cl2(g)
-102
-34
Al(s)
+660
+2519
MgCl2(s)
+714
+1412
H2O(l)
0
+100
Ag(s)
+962
+2162
NaCl(s)
+801
+1465
C6H6O(l)
+41
+182
Cu(s)
+1085
+2562
MgO(s)
+2825
+3600
C6H12O6(s)
+146
decomposes
The melting points for metals and ionic compounds are much higher
than those of molecular compounds. The same trend is observed for the
boiling points.
Metallic bonds and ionic bonds are formed by the electrostatic
attractive forces between whole positive and negative charges, which
create very strong bonds.
Within the ionic compounds,
note that the magnesium oxide
(MgO) has much higher melting
and boiling points. Magnesium
(2+) and oxide (2-) have larger
charges than most of the ions in
the list, which have a 1+ or 1charge. Larger charges create a
much stronger attractive force
between ions.
The intermolecular forces between
molecules result from the interaction
between partial positive and negative
charges, which are much weaker
than ionic and metallic bonds.
Much less energy is required to
break the weaker bonds between
molecules in molecular compounds.
Consequently, they have much lower
melting and boiling points.
Although molecular compounds
have much lower melting and
boiling points, there is large
variability of melting and boiling
points among the molecular
compounds.
The Melting and Boiling Points of Diatomic Molecules
Substance
Number of electrons
in a molecule
Melting Point
(oC)
Boiling Point
(oC)
H2(g)
2
-259
-253
N2(g)
14
-210
-196
O2(g)
16
-219
-183
*Note: The substances
that have similar
numbers of electrons
also have similar
melting and boiling
points.
*Note: As the number
of electrons in these
molecules increases,
Cl2(g)
34
-102
-34
the melting and boiling
points also increase as
Br2(l)
70
-7.2
+59
a result of the stronger
London (dispersion)
I2(s)
106
+114
+184
forces.
 All of these molecules are symmetrical, non-polar molecules with a linear shape.
 Therefore, the only interactions between the molecules are LDFs (London forces)
F2(g)
18
-220
-188
The Melting and Boiling Points of Molecules of Similar Size
Substance
Number of electrons in a molecule
Melting Point
(oC)
Boiling Point
(oC)
H2O(l)
10
0
+100
H2S(g)
18
-86
-60
CO2(g)
22
(sublimation point) -78




In these compounds that have similar numbers of electrons, the previous trend is not
clearly observed.
Since carbon dioxide becomes a gas at temperatures lower than the boiling points of
the other two, it must have the weakest intermolecular interactions. This can be
explained as it is the only non-polar molecule (it has a linear shape AX2) in the list.
Therefore, although it may have the strongest of the London (dispersion) forces, it
does not also have the dipole-dipole interactions found in the other two molecules.
Water and hydrogen sulphide are both bent molecules (AX2E2) and are polar. Yet
why does water (with fewer electrons and therefore weaker London forces) actually
have much higher melting and boiling points?
Water molecules have stronger attractive forces between them because of the
hydrogen bonding that takes place between water molecules. Hydrogen bonding is
not evident between hydrogen sulphide molecules.

The following graph will help to illustrate the influence of hydrogen bonding.
(The dotted line shows what the boiling point of water would probably be if it had no
hydrogen bonds.)

The hydrides of the Group 14 elements are all symmetrical molecules with identical
bonded atoms. They are all tetrahedral (AX4) and non-polar. As a result, the only
intermolecular forces present are London (dispersion) forces and the trend of
increasingly larger molecules having increasing higher boiling points is observed with
no exception.

The hydrides of the Group 15, 16, and 17 elements are all symmetrical molecules but
they are all polar. Group 15 hydrides (AX3E) are all trigonal pyramidal, Group 16
hydrides (AX2E2) are all bent, and the Group 17 hydrides (AXE3) are all linear.

Furthermore, in the molecules of these hydrides in which the central atom is from
period 2 (N, O and F), the dipole-dipole interaction is particularly strong (hydrogen
bonding).
Chem 20
States of Matter Assignment
1. Predict which substances will have the higher boiling point in each of the following
pairs. Explain your predictions using bonding theories.
a. NH2Cl or PH2F
c. NH4Cl or CH3Br
e. CO2 or SO2
b. Ne or Xe
d. CH4 or C2H6
f. CH3OH or C2H5NH2
2. Compare the bonding and molecular polarity of SeO3(s) and SeO2(s).
3. List the following substances in order of increasing boiling points: C3H8, ZnO,
C7H15OH, and C5H12. Give a reason for your answer.
4. The alkanes are a family of non-polar molecular compounds containing carbon and
hydrogen. The data below contains the melting points and boiling points for alkanes
having 1 to 10 carbon atoms. Using graph paper (and a ruler!!|!), draw a graph
representing this data. Your graph should be a labelled line graph that plots
temperature (on the y-axis) vs. molecular electron count (on the x-axis).
Name
Formula
methane
Number of
Electrons per
molecule
Melting Point
(oC)
Boiling Point
(oC)
CH4
-182
-161
ethane
C2H6
-183
-89
propane
C3H8
-188
-42
butane
C4H10
-138
-1
pentane
C5H12
-130
+36
hexane
C6H14
-95
+69
heptane
C7H16
-91
+98
octane
C8H18
-57
+126
nonane
C9H20
-53
+151
decane
C10H22
-30
+174