MELTING



The
& BOILING POINTS
STRONGER
the forces (including both the
intramolecular and intermolecular forces) holding a
compound together, the HIGHER the melting & boiling
points.
Ionic bonds are STRONGER than covalent bonds, thus
making their melting points HIGHER
.
Covalent molecules which have STRONG
intermolecular forces (such as water, with its strong
hydrogen bonding) are more DIFFICULT to pull apart
from each other, which causes their melting points to
be HIGHERthan compounds without hydrogen bonding
or dipole-dipole forces (such as F2).
MELTING &BOILING POINTS OF
HALOGENS
in covalent molecules which experience only London
Dispersion Forces, recall that LARGER molecules
have STRONGER London Dispersion Forces;
 Consider the halogens:

Melting
Halogen Point (C)
Boiling
Point (C)
F2
-220
-188
Cl2
-101
-34
Br2
-7
58
I2
114
183
VOLATILITY
 Volatility:
is the tendency of a liquid at room temperature to
evaporate into a gas


In order to evaporate, the liquid molecules must
possess enough K.E. to overcome the forces
holding the molecules together.
Among the halogens, volatility decreases as you
go down the group, for exactly the same reasons
greater
that melting & boiling points increase:
mass (and thus greater
# of e-)  stronger
van der Waal forces.
CONDUCTIVITY


Recall that metallic solids have delocalized valence
electrons, which enables them to readily conduct
electricity.
By contrast, most covalent and ionic solids are nonconductors, because their valence electrons are not
delocalized.




In covalent molecules, the electrons are shared between just a
few atoms.
In ionic compounds, the valence electrons are completely lost or
gained to other atoms, to satisfy each atom’s octet rule.
Ionic substances become conductive in the liquid
state, however, because the ions can move (in an
ionic liquid, the ions carry the electric current).
Likewise, many ionic solutions will conduct
electricity, if ionic solid can be dissolved in water,
producing an aqueous solution
SOLUBILITY
 Recall
that “ like dissolves like ”.
In other words, solvents tend to dissolve solutes
with similar properties
.
 Polar solvents tend to dissolve polar
solutes.
 ex #1: NH3 dissolves in water to make household

cleaners
Ammonia & water are both polar, having
unshared e–s on the central atom of their
molecules.
Non-polar solvents tend to dissolve non-polar
solutes.
 ex: I2(s) dissolves in CCl4(l)

EX # 2: WATER DISSOLVES SODIUM
CHLORIDE BECAUSE
 The NaCl ions are attracted to the very polar
H2O molecules, which play “tug-of-war” with the
ionic bonds holding the Na+ and Cl– ions
 The NaCl dissolves in water because the pull of
hundreds of H2O molecules is stronger than the
ionic forces holding Na+ and Cl– ions together.

Consider also the solubilities of the following alcohols in water:
Alcohol Solubility in Water
(mol / 100g H2O)


name
alcohol
solubility
methanol
CH3OH
infinite
ethanol
C2H5OH
Infinite
propanol
C3H7OH
Infinite
butanol
C4H9OH
0.11
pentanol
C5H11OH
0.03
hexanol
C6H13OH
0.0058
heptanol
C7H15OH
0.0008
Alcohols
dissolve in water
as the – OH
group is able to
hydrogen bond
with H2O
molecules, and
be pulled in
solution
Thus, as the length of the alcohol molecule increases, water
molecules have less of a chance of hydrogen bonding with the
water
(more of the interactions with water occur at the non-polar C – C
and C – H groups, rather than the polar –OH group). This is why
larger alcohol molecules do not dissolve as readily in water.