Crystal structure and architecture
A fundamental feature of diamond is its spatial organization. Carbon atoms are in the $sp^3$-hybridization state, which determines the tetrahedral environment of each atom. The distance between neighboring atoms ($C-C$ bond length) is strictly fixed and equals $154 \text{ pm}$.
The crystal unit cell has a cubic shape with an edge length ($a$) of $356.7 \text{ pm}$. The atoms inside the lattice are distributed as follows:
- Occupy all the corners of the cube.
- Located at the centers of each face.
- Four atoms are embedded inside the cell (they occupy the centers of four of the eight possible small cubes, arranged in a checkerboard pattern).
Chemical composition and discovery
For a long time, the nature of this mineral remained a mystery until 1796, when chemist S. Tennant conducted a key experiment. The essence of the experiment was the controlled combustion of the gemstone.
It was found that burning diamond produces carbon dioxide ($CO_2$). The main discovery was that the volume of the released gas absolutely matched the amount of $CO_2$ obtained from burning the exact same mass of ordinary graphite. This led to an unambiguous conclusion: the mineral's structure consists exclusively of carbon atoms without any impurities of other elements.
Physical properties and identification method
Visually, the mineral consists of transparent crystals (usually of octahedral or cubic shape) characterized by very strong refraction. In jewelry, cut specimens are called diamonds. Among all substances known to science, diamond has the highest hardness, and its density is $3.513 \text{ g/cm}^3$.
In practice, it is important to be able to distinguish natural samples from artificial glass imitations. For this purpose, a specific chemical test is used with a concentrated aqueous solution of a complex salt — potassium tetralliodomercurate(II) ($K_2[HgI_4]$). Due to the significant difference in material densities, glass fakes always remain on the surface of the liquid, whereas the true carbon crystal sinks.
Thermodynamics and phase transitions
From a thermodynamic standpoint, under standard conditions graphite is a more stable form of carbon than diamond. Theoretically, the spontaneous transition of a transparent crystal into graphite is always possible, but at normal temperature, this process is kinetically hindered (the rate is practically zero). The situation changes upon strong heating: at $2000^\circ C$, the graphitization process proceeds very rapidly.
The reverse process — artificial synthesis — is also possible. Converting carbon into a diamond structure requires colossal pressure exceeding $2 \cdot 10^6 \text{ kPa}$. Currently, mass production of industrial diamonds is established using this technology. For comparison, the transition of coal to graphite requires a pressure of only hundreds of atmospheres.
Chemical properties
Being quite inert at room temperature, the material begins to show reactivity under harsh conditions. Main chemical reactions:
- Oxygen oxidation: In open air, the substance actively oxidizes and completely burns upon reaching a temperature of $900^\circ C$.
- Interaction with melts: The crystal lattice is effectively destroyed by strong oxidizers — molten alkali metal nitrates or carbonates.
Equations of oxidation reactions in melts:
- With sodium carbonate (with carbon monoxide release):
$$C + Na_2CO_3 = Na_2O + 2CO$$
- With sodium nitrate (with carbon dioxide formation):
$$C + 2NaNO_3 = CO_2 + 2NaNO_2$$
Place among other carbon modifications
In addition to diamond and graphite, carbon forms other shapes. Carbyne consists of linear molecules with atoms in the $sp$-hybridization state (its synthesis was predicted based on the theory of hybrid states). Fullerenes are complex polyhedra, and in nanotubes atoms are ordered along the cylinder surface.
Interestingly, free carbon (in the form of soot) can be reduced from carbon dioxide using active metals. For example, the combustion of magnesium in a $CO_2$ atmosphere is accompanied by the formation of black smoke according to the reaction: $$2Mg + CO_2 = 2MgO + C$$
A "heavy" solution of tetraiodomercurate will help distinguish a fake: light glass will stay afloat, while dense diamond (3.513 g/cm³) will sink to the bottom.
Class and exam questions
Why doesn't diamond turn into graphite right before our eyes?
Although thermodynamically graphite is indeed more stable under normal conditions, the transformation reaction is kinetically hindered. At room temperature, the rate of this process is practically zero, and it proceeds rapidly only upon heating up to 2000 °C.
How did S. Tennant prove that diamond is pure carbon?
In 1796, he burned a diamond and found that it produces carbon dioxide in exactly the same amount as burning a corresponding mass of ordinary graphite.
With which aggressive substances does diamond react?
It oxidizes in molten alkali metal nitrates and carbonates, forming alkali metal oxides or nitrites, with the release of carbon monoxide or carbon dioxide, respectively.
What else to study on this topic
- Other allotropic modifications of carbon: carbyne, fullerenes, nanotubes
- Thermodynamics of carbon phase transitions and kinetic control of reactions
- Industrial technologies for the synthesis of industrial diamonds at ultra-high pressures
- Reduction of carbon from carbon dioxide by active metals
Master the entire topic in the "Chemistry" course
This is a summary. The course covers everything tested in classes and exams.
- Full study guide on "Diamond" and 3,009 more pages of the subject
- Assignments and quizzes after every topic — just like in a colloquium
- AI tutor answers questions on the topic anytime
- 95 video lessons for the subject
Other topics in "Chapter 12. Group IVA Elements"
Prepared by WhiteCoat Pro faculty in accordance with the medical school curriculum. For educational purposes only.