Martha
Pamato
Travel into a Diamond
Diamond is a truly remarkable material, it is the hardest naturally occurring mineral, it is an excellent thermal conductor and electrical insulator and has a fantastic array of optical properties making it one of the most valuable materials on Earth!
Take a trip inside a diamond and find out what makes this mineral so special!
Diamond vs. graphite
Despite its famous durability, diamond is not thermodynamically stable at surface P/T conditions, graphite is!
In diamond, each C atom is strongly (covalently) bonded to four others. The hardness of diamond is due to the arrangement of C atoms in cubic symmetry.
Only 3 valence e that form sp bonds, 4 e delocalized across sheet
-
th
-
2
sheets linked via Van der Waals forces
All 4 valence e participate in C-C bonding via sp orbital hybridization
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3
In graphite, each C atom is bonded to three others forming sheets. These sheets are weakly bonded predisposing graphite to cleaving (breaking) along these sheets.
graphite
diamond
Dive into the structure of diamond and graphite above!
rotate, zoom-in and out, or expand to explore
Diamond crystallography
The unit-cell of diamond contains 8 C atoms arranged in cubic symmetry
Bravais lattice: fcc
Space group: Fd3m
Point group: Fm3m
Atomic density n = 8/a
Density of surface atoms
(100): 2/a
(110): 4/√2a
(111): 4/√3a
atoms
3
2
2
2
a = 3.567 Å
b = 3.567 Å
c = 3.567 Å
α = 90.00 º
β = 90.00 º
​ɣ = 90.00 º
Volume = 45.15 ų
The crystal structure of diamond can be simply described as two interpenetrating face-centered cubic lattices displaced with respect to one another along the body-diagonal of the cell by 1/4 the length of the diagonal. This displacement length is equal to the C–C bond length which is ~ 1.54 Å
Visualize the unit-cell of diamond in 3-dimensions below!