1. melting & boiling points
The the forces (including both the intramolecular and intermolecular forces) holding a compound together, the the melting & boiling points.
Ionic bonds are than covalent bonds, thus making their melting points
Covalent molecules which have intermolecular forces (such as water, with its strong hydrogen bonding) are more to pull apart from each other, which causes their melting points to be than compounds without hydrogen bonding or dipole-dipole forces (such as F2).
STRONGER
HIGHER
STRONGER
HIGHER
STRONG
DIFFICULT
HIGHER

2. MELTING &BOILING POINTS OF HALOGENS
in covalent molecules which experience only London Dispersion Forces, recall that molecules have London Dispersion Forces;
Consider the halogens:
LARGER
STRONGER
Halogen
Melting Point (C)
Boiling Point (C) F2
-220
-188
Cl2
-101
-34
Br2 -7 58 I2
114
183

3. volatility Volatility:
In order to evaporate, the liquid molecules must possess to overcome the forces holding the molecules together.
Among the halogens, volatility as you go the group, for exactly the same reasons that melting & boiling points increase: mass (and thus # of e-)  van der Waal forces.
is the tendency of a liquid at room temperature to evaporate into a gas
enough K.E.
decreases
down
greater
greater
stronger

4. conductivity
Recall that metallic solids have delocalized valence electrons, which enables them to readily conduct electricity.
By contrast, most covalent and ionic solids are non- conductors, 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

5. solubility Recall that “ dissolves ”.
In other words, solvents tend to dissolve solutes with
Polar solvents tend to dissolve solutes.
ex #1: NH3 dissolves in to make household cleaners
Non-polar solvents tend to dissolve non-polar solutes.
ex: I2(s) dissolves in CCl4(l)
like
like
similar properties
polar
water
Ammonia & water are both polar, having unshared e–s on the central atom of their molecules.

6. 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.

7. Alcohol Solubility in Water
Consider also the solubilities of the following alcohols in water:
Thus, as the length of the alcohol molecule increases, water molecules have 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
alcohol molecules dissolve as readily in water.
Alcohol Solubility in Water
(mol / 100g H2O)
name
alcohol
solubility
methanol
CH3OH
infinite
ethanol
C2H5OH
Infinite
propanol
C3H7OH
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
less
larger
do not