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Diamonds are crystalline polymorphs or allotropes of carbon. It is a metastable allotrope with a unique crystal structure called the diamond cubic. Metastable means that diamond is not the least stable allotrope or solid form of carbon and stabilizes into the most chemically stable crystalline form, graphite, at extremely slow (near negligible rates).
While used commonly as a precious stone, the diamond’s incredible hardness, extremely high chemical & thermal inertness, and abrasion resistance make it highly useful in different applications.
The chief reasons behind their immense popularity across all industries are the amazing hardness, abrasiveness, chemical, thermal & electrical inertness, and natural lustre of diamonds. The structure and bonding of the carbon atoms/molecules lend all such properties.
This metastable allotrope of carbon has a cubic crystal structure with face-centred cubic lattices. The unit cell of the cubic lattice has four covalently bonded carbon atoms arranged in a tetrahedral fashion.
The face-centred cubic lattice-based crystal structure of diamonds looks something like this à
As can be seen from the visualisation, four carbon atoms inside the volume of the cubic structure are connected to four other atoms.
Six carbon atoms are found in the face of each cube, each having two bonds. Of the atoms at the vertices, four have four bonds, three with the atoms of neighbouring vertices and one with atoms inside the volume.
This abundance of covalent bonds among the carbon atoms makes diamonds so hard and inert. Every carbon atom shares its electrons with four other carbon atoms through covalent bonds of equal strength. A single carbon atom has six electrons, 2 in the K shell/s-orbital and 4 in the valence L shell/2 in the s-orbital and 2 in the p-orbital. The electronic configuration of a carbon state in the ground state is something like this à
However, the electronic configuration of a carbon atom does not account for the tetrahedral nature of its crystalline structure or the extreme toughness & inertness of a diamond molecule.
The next section analyses the crystalline structure of diamonds minutely.
The basic atomic arrangement in diamonds is something like this à
Every carbon atom connects with four other carbon atoms. And this tetrahedral atomic arrangement repeats itself in the three-dimensional space of a diamond. This strong, rigid three-dimensional structure forms a repeating network with all the other carbon atoms.
The basic crystal structure of diamonds has two basis atoms and what might seem to be two intersecting or interpenetrating face-centered cubic lattices. The lattices are so arranged and located that four other carbon atoms surround every carbon atom in each of the interpenetrating lattices.
The face-centered lattice structure has a two-atomic basis. There are two carbon atoms at every lattice point; one is located at the lattice point, and the other is shifted by a (1/4, 1/4, ¼) vector from it. Hence, the number of atoms in the unit cell of a diamond cubic lattice goes up from 4 to 8. Thus, a tetrahedral structure forms with each carbon atom surrounded by four other equidistant neighbours.
From the above, the cubic diamond structure combines two FCC sublattices that interpenetrate with one another along the diagonal of the cubic unit cell. We can also note carbon atoms’ basic tetrahedral bonding arrangement in diamond.
The total number of atoms in each unit cubic cell is eight, four for each lattice. For such a simple and symmetric structure to form, every carbon atom must exhibit similar and symmetric properties. And it is the tetrahedral arrangement that provides optimal stability to the fundamental molecular arrangement of diamonds.
It is the covalent bonds of the carbon atoms that lend it its unique crystalline arrangement as well as physical properties. Let’s dissect the nature and features of diamond carbon covalent bonds.
The monoatomic carbon molecules in diamond are sp3 hybridized (more on this later). Each carbon atom has a strong covalent bond with four neighboring carbon atoms. Because of these covalent bonds and this dense interpenetrated lattice network, diamond possesses both high melting & boiling points.
The covalent bonds hold the electrons firmly, preventing them from being easily displaced by applying heat or electricity. There are no free electrons present that can be excited and made to leave the crystal structure.
So, what covalent bonds make diamonds so tough and inert?
For diamond crystals to possess a stable tetrahedral symmetry, the electronic configuration of carbon must be altered. A carbon atom has single covalent bonds with four other carbon atoms; hence they undergo sp3 hybridization in the L or outermost shell. This hybridization or alteration leads to the formation of hybrid atomic orbitals in the L shell of a carbon atom in the ground state.
The L or 2nd shell is the outermost shell of carbon atoms. During hybridization, a single electron from the 2s orbital get lifted to the 2p orbital. This leads to a hybrid sp3 orbital formation as the 2s and 2p orbital merge. It is this hybridization that lends a tetrahedral symmetry to the fundamental molecular structure of a diamond. The sp3 orbitals of every carbon atom bond with the sp3 orbitals of four other carbon atoms. Each atom undergoes covalent bonding with another by sharing one electron each and forms a regular tetrahedral structure with equal angles of 109 degrees 28 minutes.
But how does all of this go down? How do covalent bonds form?
So, what is a chemical bond? What definition is the most accurate?
Research methodologies and techniques of modern chemistry define chemical bonds as overlapping two atoms’ electron waves. There are two types of chemical bonds, covalent and ionic or electrovalent.
In the case of covalent chemical bonds, the overlapping can be seen as two atoms sharing valence electrons. Ionic or electrovalent bonds involve the transfer of electrons from a negatively charged/electron-rich ion to a positively charged/electron-deficient ion. Covalent bonds do not involve any transfer of electrons but sharing of electrons.
Any atom that partakes in covalent bonding (or even ionic bonding) looks to acquire a stable electronic configuration in its valence shell. If you are familiar with the Periodic Table, you must know about the column of noble gases. When undergoing any bond, every element in the periodic table looks to acquire a stable electronic configuration in its valence shell, the same as that of the valence shell of the noble or inert gas in its row.
The carbon atoms in diamonds undergo sp3 hybridization before forming covalent bonds with other carbon atoms. Each bond is around 154 picometres long, the same as the carbon-carbon bonds in organic compounds such as ethane, propane, and other kinds of alkanes.
Hybridization is the scientific term for the phenomenon where atomic orbitals fuse to form metastable hybrid orbitals. The sp3 hybridization of carbon atoms plays a central role in developing the tetrahedral structure of diamond molecules and imparting unmatched physical properties.
Sigma and pi bonds are two types of covalent bonds. While the sigma bond is formed by axial or end-to-end overlapping of half-filled orbitals, pi bonds form via lateral or side-wise overlapping of half-filed orbitals.
Electron clouds of sigma bonds are cylindrical and symmetrical along the nuclear axis of the bonded atoms. Pi bonds exhibit elongation above and below the atomic or nuclear plane.
Covalent bonds are not as strong as their ionic counterparts. However, that does not mean that covalent bonds are not strong and stable enough. The number of covalent bonds in the interpenetrating FCC cubic lattices makes diamonds so strong and resilient.
Here’s a quick look at the myriad physical properties of diamond à
Materials science defines durability as the ability of any substance to resist deterioration caused by physical, chemical, and biological agents on it. In the case of diamonds, the gem industry identifies three parameters to define its durability – hardness, toughness, and stability.
Diamonds are some of the most highly lustrous substances on the planet. Properly cut diamonds reflect light without any distortion. The refractive index of a diamond (when measured in sodium light of wavelength 589.3 nanometre) is 2.417.
The value of light dispersion is around 0.044, evident from the fire manifested from cut diamonds.
Some diamonds also exhibit fluorescence when exposed to certain colours at certain intensities, especially at the long-wave ultraviolet portion of the visible spectrum. When cooled to very low temperatures, certain kinds of diamonds also possess a particular absorption spectrum. Analysing the absorption spectrums of diamonds through spectrometers enables gem companies to distinguish between natural, artificial, and enhanced diamonds.
Diamond is a good conductor of heat, and this is because of the multiple strong covalent bonds in the unit crystal cells. Monocrystalline synthetic diamonds showcase 2000-2500 W-m/m2-K thermal conductivity values, five times higher than the next best thermal conductor, copper.
The unique arrangement of covalently bonded carbon atoms makes diamond the most thermally conductive substance known to man.
Pure diamonds are not a good conductor of electricity. However, natural blue diamonds with boron impurities can act as semiconductors. Fabrication and large-scale production of boron-doped diamond semiconductor chips is an ongoing field of research in the materials science domain and semiconductor industry.
Boron-doped synthetic diamonds have also been found to act as a type-II superconductor at a superconducting transition temperature of 4 Kelvin.
Well, that was quite an extensive discussion on diamonds’ myriad physical and chemical properties. Discussions would be incomplete, however, without looking at the properties of diamonds’ less-lustrous but no less important counterpart allotrope, graphite. Let’s also look at amorphous carbon, another highly useful carbon allotrope, along with graphite.
Graphite is the other amazing allotrope of crystalline carbon. It is extremely soft and brittle but is also extremely resistant to heat & almost any kind of intense chemical reaction. At the same time, graphite is a good conductor of electricity due to carbon atoms in graphite forming three bonds instead of four. The remaining electron is what makes graphite a good conductor.
The graphite crystal structure comprises meshes of inter-bonded carbon atoms arranged in stacks upon one another. The mesh layers are quite spread out or squashed, and the interacting force between each layer is very low. (Van Der Waal’s force)
Each carbon atom in graphite forms three bonds, which leaves a fourth unbounded electron at the bonding level. With all the unbounded electrons from all over the multiple sheets, the conductivity of graphite becomes quite high.
Graphite’s layered structure and weak interacting force between the layers make them very soft and brittle. Graphite is metallic, while diamond is lustrous and brilliant. Diamond has high thermal conductivity but zero electrical conductivity; conversely, graphite is an excellent thermal insulator with very high electrical conductivity.
As the term suggests, amorphous carbon is an allotrope without long-range crystalline order. Amorphous carbon is a term used to denote any carbon material without any crystalline order and having localized pi electrons. However, short-range crystalline order does exist in them.
Coal, coke, soot, charcoal, carbon black, etc., are all different amorphous carbons. In truth, the short-range crystalline order imparts polycrystallinity and nano-crystallinity.
The structural diversity of amorphous carbon materials stems from the varying ratio of sp3 to sp2 hybridizations. Most amorphous carbon materials possess dangling pi bonds that render them unstable. Amorphous carbons with more sp3 hybridized carbons are known as diamond-line carbons.
Amorphous carbons do not have long-range crystalline patterns like diamonds. This makes them easily breIntroduction to Diamondsakable and, along with freely-dangling pi bonds, more reactive than diamonds.
Industrial applications of amorphous carbon are many. They are used in energy production, semiconductor manufacturing, solar cell manufacturing, production of carbon black, gun powder, printing ink, paint, water filters, electrodes, oil refining, and much more.
Let’s wrap things up with a look back at Diamond and some of its amazing and diverse applications.
Diamonds are most famous for being incredibly precious stones. But, recent advances in technology and deeper analysis of its properties have made diamonds a key part of several other industries.
Diamonds are now used extensively in ultra-precision cutting & polishing tools. Nano-scale machining using diamond-tipped tools is a key aspect of nanotech and nanostructure fabrication.
Check this link out for some awesome insights.
Monocrystalline synthetic diamonds exhibit the highest magnitude of thermal conductivity among any known natural or manufactured material. Ranging between 2000-2500 W-m/m2-K, it is five times higher than copper.
This makes diamonds the perfect material for creating heat sinks. Diamond heat sinks are used in the semiconductor industry to control overheating of manufactured semiconductor materials. Diamond heat sinks find applications in electronics manufacturing and several specialized engineering applications.
Boron-doped synthetic diamonds possess excellent semiconductor properties. A wide bandgap, high dielectric strength & breakdown field, and high thermal conductivity make them ideal, albeit extremely costly, semiconductors.
The development of diamond field-effect transistors is an ongoing field of research. FETs for high voltage applications have been designed and even developed though it could be years before we get a feasible and economical device in the market.
There have been substantial recent developments in the utilization of diamonds in healthcare. Using nanodiamonds in prosthetics, sensing & imaging, optical implants, and drug delivery are major sub-fields witnessing remarkable R&D.
Diamonds may become integral in quantum computing. Researchers are experimenting with doped synthetic diamonds with unpaired electrons, using quantum superposition to spin the electron up and down. Such dualistic behavior can be represented as a qubit/quantum bit in quantum information processing systems. Find out more here.
And that brings us to the end of this article on diamonds, their binding & crystallographic structure, applications, and their fellow allotropes. I hope it was an informative read for everyone.
Chemistry is an astonishing subject with its challenges. There are heavy interconnections among nearly every topic and sub-topic. Work hard and study minutely to excel. And if need be, seek expert academic assistance from MyAssignmenthelp.com.
The atomic structure is tetrahedral, with each carbon atom forming string sp3 hybridised sigma bonds with four other carbon atoms.
They are arranged tetrahedrally in a three-sided pyramid structure. The structure is well-defined and very rigid, with each bond at an angle of 109 degrees 28 minutes to the other.
sp3 hybridised single sigma covalent bonds exist between each carbon atom in diamonds.
Each carbon atom has four strong sp3 hybridised sigma bonds. Along with interpenetrating FCC lattices, this makes diamonds so hard to break, melt, or react with other materials.
Diamond has a tetrahedral atomic structure and a unique cubic lattice of two inter-penetrating face-centred cubic lattices. The C-C bonds are strong sp3 hybridised sigma covalent bonds.
Graphite, the other major crystalline allotrope of carbon, has sp2 hybridised covalent bonds with a free electron in every bond. Carbon atoms form mesh-like layers stacked upon one another and held together by weak Van Der Waal’s force.