Crystal Structures
Particle Diagram
Type of
crystal
Dia
mo
nd
Gr
ap
hit
e
Ice
Melting and
Boiling Points
Solid
Liquid
Gas
Electrical
Conductivity
Hardness
/Brittleness
Solubility
Density
Crystal Structures
Particle Diagram
Type of
crystal
Iod
ine
Ma
gn
esi
um
So
diu
m
Chl
ori
de
Melting and
Boiling Points
Solid
Liquid
Gas
Electrical
Conductivity
Hardness
/Brittleness
Solubility
Density
Particle Diagram
Type of crystal
Macromolecular
(FCC – face
centered cubic)
Dia
mo
nd
Melting and
Boiling Points
High - giant
macromolecular,
lots of strong
covalent bonds.
Mp: 4000
(usually
sublimes)
Bp: -
Macromolecular
(Hexagonal)
Gr
ap
hit
e
High (but not as
high as
diamond) - giant
macromolecular,
lots of strong
covalent bonds.
Weak VdW
forces between
layers.
Solid
Liquid
Liquid carbon
– achievable
at only 100
atm and 4500
K.
Gas
Gaseous
monatomic
carbon
Usually
sublimes.
Sublimes
rather than
melting.
Gaseous
monatomic
carbon
Electrical
Conductivity
Hardness
/Brittleness
Poor – all 4
valence
electrons in
carbon taken up
in bonding pairs
– so no free
charges.
Very hard (each C
bonded to 4 others).
Strong covalent
lattice – rigid
tetrahedral structure.
Good – one
non-bonding
electron per
carbon atom.
Delocalised
electrons can
move through
layers and carry
charge.
Soft – each C
bonded to 3 others
– arranged in
hexagonal layers –
can slide.
Poor – no
charged
particles.
Hard – tetrahedral
H-bonding structure.
Brittle – cubic planes
have weaknesses
which when hit can
break apart.
Brittle – weak VdW
forces between
layers easily broken
–o layers break off
easily.
Solubility
Density
Insoluble in
water/polar
solvents – no
charges.
Insoluble in
non-polar
solvents –
covalent
bonds too
strong.
Insoluble in
water/polar
solvents – no
charges.
Insoluble in
non-polar
solvents –
covalent
bonds too
strong.
Mp: 3600
(sublimes)
Bp: Molecular
(Hexagonal)
Ice
Low – weak
intermolecular
forces do not
need a lot of
energy to break.
Higher mp/bp
than other
molecular
substances due
to H-bonding.
Mp: 0
Bp: 100
Brittle – H-bonds
easily broken as
they are
intermolecular
forces (albeit the
strongest IM force).
Soluble in
polar solvents
– water is a
polar
molecule.
Immiscible in
non-polar
solvents –
water is a
polar
molecule so
no stabilising
interactions
can form.
0.917 gcm-3
Ice less dense than
liquid water –
regular orientation
of H bonds causes
tetrahedral structure
and molecules to be
spaced far apart.
In water, H-bonds
constantly breaking
and reforming –
less rigid structure.
Particle Diagram
Type of
crystal
Iod
ine
Melting and
Boiling Points
Molecular –
held together by
weak VdW
forces.
Low – weak
VdW forces
between
molecules.
(FCC – face
centered cubic)
Mp: 113.7
Bp: 184.3
Metallic
High – strong
metallic bonds
throughout
structure –
strong
electrostatic
attraction
between
positive metal
nuclei and
delocalised
electrons.
(HCP –
hexagonal close
packed)
Ma
gn
esi
um
Solid
Liquid
Gas
Electrical
Conductivity
Poor – no
charged
particles.
Hardness
/Brittleness
Soft and brittle –
weak VdW forces
between molecules
easily broken.
Solubility
Sparingly soluble in
water/polar solvents –
iodine is non-polar.
Soluble in non-polar
solvents – forms weak
intermolecular
attractions (VdW).
Metallic bonds
not fully broken.
Gaseous metal
atoms.
Good – lots of
delocalised
electrons can
move throughout
structure and
carry charge.
Soft – regular
layered arrangement
allows layers to slide
over each other,
making it malleable
(not brittle).
Poor solubility in both
polar and non-polar
solvents – metallic
bonds too strong
(always stronger than
solvent to solute
attractions).
Delocalised
electrons prevent
fragmentation – can
move around lattice.
Mp: 650
Bp: 1091
Ionic
So
diu
m
Ch
lori
de
(FCC – face
centered cubic)
High – lots of
strong ionic
bonds (in all
directions) –
strong
electrostatic
attraction
Mp: 801
Bp: 1465
Conducts when
molten/dissolved
– lattice broken,
ions free to
move.
Don’t conduct
when solid – ions
locked in lattice.
Hard - strong
electrostatic
attraction in all
directions.
Brittle – when layers
of alternating
charges distorted,
like charges repel.
Breaking the lattice
into fragments.
Soluble in water – water
is polar solvent and
stabilises separated
ions. Polar water
molecules surround the
ions.
Insoluble in non-polar
solvents (no dipole
interactions like in polar
solvents). Electrostatic
attraction of ionic
bonding too strong.
De
nsi
ty
Crystal Structure
Challenge Q
Answer
Diamond
Silicon and silicon dioxide have
similar structure and
properties to diamond.
Compare with explanations.
Crystalline silicon has the same structure (tetrahedral) as diamond.
Silicon dioxide also has the same structure (tetrahedral) only some silicon atoms are now
replaced by oxygen atoms. Each silicon bridged to neighbours by an oxygen atom.
Both have high mp – around 1700 for silicon dioxide – due to strong covalent silicon-oxygen
bonds. Hard like diamond – rigid structure, bonds difficult to break. Again it is insoluble in
water and organic solvents – no attractions throughout the structure that would be more
favourable than breaking the strong covalent bonds.
Both do not conduct – lack of charged particles.
Silicon has bigger atoms than carbon – so the bonds are longer and weaker – meaning silicon
has a lower melting point (1410) than diamond.
Silicon dioxide also has a much longer bond length than diamond – so its mp is also lower.
Graphite
Graphene has similar structure
to graphite. Compare structure
and properties with
explanations.
Graphene is a single sheet of graphite. Carbon atoms are still hexagonally bonded. Very high
strength to mass ratio – no layers mean the structure cannot slide – so the softness of
graphite is gone. Very strong due to strong covalent bonds.
Can still conduct electricity – as delocalised electrons can still move throughout the structure
and carry charge. One electron from each carbon does not partake in bonding pairs.
Extremely thin as it is a single layer of graphite.
Ice
The structure of ice allows it to
float on liquid water, which is
unusual for a solid. Explain how
this is possible.
Ice is less dense than liquid water. Water is a polar molecule (because oxygen is so
electronegative, and much more electronegative than hydrogen). The oxygen atoms also
have 2 lone pairs of electrons.
The delta positive H is attracted strongly to the LP on the delta negative O – forming a H
bond. In liquid water, the hydrogen bonds are being continually broken and reformed. In
solid ice, the hydrogen bonds are fixed.
Each water molecule forms 4 hydrogen bonds, forming a tetrahedral geometry shape – to
maximise the distance between the negative oxygen atoms. The lower temperature of ice
means there is less energy to h bonds cannot be continually broken and reformed. At least
18 phases (packing geometries) of ice.
Crystal Structure
Challenge Q
Answer
Iodine
Chlorine and bromine have the
same structure as iodine. Compare
structures and properties with
explanations.
Iodine is larger than both bromine and chlorine atoms. In order of increasing size:
Chlorine → bromine → iodine.
The bond length of iodine is therefore the greatest as it has the greatest atomic radius.
Increase in atomic radius means the VdW forces between molecules are greater, leading to an
increase in melting point with increasing atomic size.
VdW forces are induced dipoles caused by the random movement of electrons within the
structure – the more electrons, the more possible dipoles can be created.
All dissolve in water to some extent, but all are sparingly soluble. They are still soluble in
non–polar solvents such as hexane, as both are non-polar molecules.
Bond enthalpy decreases from chlorine → bromine → iodine as the bond length increases due
to increasing atomic radius.
Magnesium
Sodium and aluminium are metals in
the same period as Mg. Compare
the structures and properties with
explanations.
Magnesium has 2 electrons in valence shell – so in its structure it has Mg 2+ ions and 2
delocalised electrons per ion in the electron cloud. As the elements are in the same period,
they will have the same shielding.
Mg2+ has higher charge and a smaller ionic radius than Na – which forms a Na+ ion and each
ion gives one electron to the electron cloud.
Magnesium is 12 –packed, meaning each Mg atom is in contact with 12 other atoms.
However, sodium is less efficiently packed, only 8-coordination. In an 8-coordination, no
atoms touch within the layer – only above and below. This less efficient packing also
contributes to weaker metallic bonds meaning that sodium has a lower mp than Mg.
Na is softer than Mg as the metallic bonding is easier to break.
Al has 3 valence electrons, forming a Al3+ ion and each ion gives 3 electrons to the electron
cloud. Aluminium ions are also much smaller than Mg2+ ions, meaning that Al has a higher
melting point due to stronger metallic bonds.
Conductivity increases from Na→ Mg→ Al as more delocalised electrons per atom.
Sodium chloride
Magnesium oxide also has this
structure. Compare structure and
properties with explanations.
Magnesium oxide is made up of Mg2+ and O2- ions. Both negative and positive ions have
higher charges than in sodium chloride (Na+ and Cl-).
This means the ionic bonding in magnesium oxide will be stronger and magnesium oxide will
have a higher melting point.
Magnesium oxide is more brittle as the ions have greater charges, so will repel more.
MgO is less soluble than NaCl – the charges are greater, meaning the bonds are stronger,
meaning it is less favourable to form attractive interactions with solvents in MgO.