Types of Solids
Molecular solids result due to weak intermolecular forces that attract the individual
molecules to each other. These solids require relatively little energy to become
liquids (or gases) and have the lowest melting points of all solids.
Research the properties of the carbon tetrahalides (CX4). Provide the state, electronic
shape, and melting point for each.
Table 1 Trends in the properties of the carbon tetrahalides
Formula
State at SATP
Electronic Shape
mp (oC)
CF4
gas
tetrahedral
-184
CCl4
liquid
tetrahedral
-23
CBr4
solid
tetrahedral
90
CI4
solid
tetrahedral
171
1. Explain the trend in melting points for the carbon tetrahalides.
Melting point of carbon tetrahalides increases down the group. They are nonpolar molecular compounds. Therefore
London dispersion forces are the only intermolecular forces present. As the number of electrons increase, so does the
number of London dispersion forces.
Ionic solids arise from the crystal lattice arrangement of positive
and negative ions that exert electrostatic attractions toward
oppositely charged ions – ionic bonds.
Ionic compounds tend to have much higher melting points due to
the strength of the ionic bonds, but this can vary.
Table 2 The chlorides of alkaline earth metals
Ionic compound
mp (oC)
magnesium chloride
714
calcium chloride
772
strontium chloride
874
barium chloride
962
Research the melting points for the compounds formed by alkaline earth metals with chloride. What is the trend that
arises for the melting point of chloride compounds in this group?
Melting point of alkaline earth metal compounds increases down a group. Electronegativity decreases down a group
creating a greater electronegativity difference when bonded. A greater electronegativity difference means it forms a
more stable bond.
2. Predict the melting point for radium chloride.
1000+
Metallic solids are held together by an arrangement of metal atoms sharing the electrons of their valence level. The
positive nuclei of each metal atom are attracted to the loosely held valence electrons through metallic bonding. View
images at http://www.ausetute.com.au/metallic.html.
Depending on the attraction of the nuclei to the valence
electrons, the strength of metallic bonds will vary, resulting in a
wide array of melting points for these solids.
Research the alkali metals to observe one of the trends noted in
the periodic table.
Table 3 Melting points of alkali metals
Alkali metal
lithium
sodium
potassium
rubidium
cesium
First ionization energy
(kJ/mol)
520
496
419
403
376
mp
(oC)
180
98
63
39
28
3. Explain the trend in the melting point for the
alkali metals.
Decreases down the group because ionization energy
decreases. Valence electrons are farther from the
nucleus, so the strength of the attractive force
decreases creating a weaker bond.
4. The only metal to exist in the liquid state, mercury has one of the broadest temperature ranges as a liquid of any
metal. Yet, mercury is unreactive. Research why mercury forms weak metallic bonds.
Mercury holds on to its valence electrons tightly because all valence orbitals are full ([Xe]4f145d106s2). This means that
electrons cannot flow freely, meaning there is a weak “electron sea”
Covalent network solids are the final type of solid using strong covalent bonds
to hold the atoms together in large structures (sometimes referred to as
macromolecules). The strength of these bonds creates the most stable solids
with the highest melting points.
The two allotropes of carbon, diamond and graphite, fit into this category.
Research the properties of these allotropes.
Table 4 The allotropes of carbon
Allotrope
VSEPR shape around
carbon
Hybridization of
carbon
Hardness
(Moh’s scale)
diamond
tetrahedral
sp3
10
graphite
trigonal planar
sp2
1-2
5. Explain the difference(s) in the intramolecular force within diamond and graphite.
The 4th carbon bond in graphite is longer and weaker, allowing the layers to slide. The tetrahedral conformation in
diamond makes it very strong.
6. Quartz also falls into this category of solids. Research the structure (formula, hybridization, bonding) and
properties (hardness, mp, solubility) of this covalent network.
SiO2, trigonal and hexagonal crystal structure, Moh’s scale = 7, mp = 1670˚C, insoluble, lustrous