Would you wear a graphite ring?

Diamonds and graphite are both composed of carbon atoms, so why do they exhibit such different properties? The answer lies in the structure differences between diamond and graphite.

In fact, the structure has a direct influence over the properties of materials. What do I mean by structure? The structure of solid materials can be described in terms how the atoms are bonded to each other (interatomic bonding) and the arrangement of atoms (crystallinity) in the material. Interatomic bonds are further divided into primary (chemical) bonds and secondary (van der Waals or physical) bonds. Primary bonds are much stronger and involve the outer electrons ( valence electrons) of the constituent atoms to create stable electron structures. There are three types of primary bonds: ionic, covalent, metallic.Ionic bonding is found in compounds composed of metallic and nonmetallic elements that are found on the two ends of the periodic table (e.g. NaCl). The metallic element gives up its valence electrons to the nonmetallic element; resulting in both elements with a stable configuration that is equal in all directions. Consequently this type of chemical bond is very stable, which is reflected in its high bonding energy (i.e. energy required to separate these two elements). Ionic bonding is typical of ceramic materials. Ceramics materials are typically hard, brittle, and electrically insulative–a direct consequence of the ionic bonding nature.

An example of Ionic bonding in magnesium oxide (MgO) where the arrows indicate non-directional bonding (modified from [1]).


Covalent bonding occurs when valence electrons are shared between adjacent atoms/molecules in a compound. They are found nonmetallic elemental molecules (e.g. H2, Cl2); general molecules (e.g. CH4, H2O); elemental solids (e.g. diamond, silicon). Although, covalent bonds are typically found in polymeric materials. Covalent bonds are directional meaning it may exist only in the direction between the atoms sharing the electrons. However, interatomic bonds can be partially ionic and partially covalent.The type of bonding in the compound depend on the position of the constituent atoms. The bond is more ionic when atoms are farther apart on the periodic table. Conversely, the bond is more covalent when the constituent atoms are closer together.

An example of covalent bonding in methane (CH4), note that the carbon and hydroen atoms each share 1 electron in this arrangement (modified from [1]).


Metallic bonding is found in compounds composed of elements from the first two columns of the periodic table as well as all elemental metals. This type of bonding is often referred to “sea of electrons” because the valence electrons are free to move throughout the metal and not bound to any atom. Thus metallic bonding is non-drectional. Due to the free electrons in metallic bonding, unlike the bound electrons in ionic or covalent bonding, metallic materials exhibit good conductivity (thermal/electrical) compared to ceramic or polymeric materials.

An example of metallic bonding showing a “sea of electrons” (modified from [1]).


Secondary bonds are typically found in most atoms/molecules; however, their effects are typically overshadowed by the primary bonds that are present. They exist as a result of atoms/molecules separating into positive and negative portions (i.e. to form dipoles).

A schematic diagram illustrating secondary bonding where positive and negative regions exist within one atom/molecule and the

secondary bond between atoms/molecules is indicated by the red arrows (modified from [1]).


Interatomic bonding reflects how atoms are bonded to each other. On the other hand, the arrangement of the atoms and the regularity of this arrangement can be described by the crystallinity of a material. Crystalline materials are composed of groups of atoms that form a repeating pattern in a long range order. All metals, most ceramics and some polymers have crystalline structures. In contrast, long-range atomic order is absent in noncrystalline materials (aka amorphous).


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Back to diamonds….
The crystal structure describes the how atoms, ions or molecules are arranged in a crystalline material. Crystalline materials are composed of groups of atoms in a repeating pattern, and each repeating unit is known as a unit cell.

d

Diamond and graphite have different crystal structures which are responsible for the properties observed. Diamond has a diamond cubic crystal structure unit cell (see below figure “a”). Every carbon atom bonds to four other carbon atoms with covalent bonds. This 3-dimensional arrangement of the covalent bonds produces a rigid structure, giving diamonds their hard properties. Furthermore the diamond structure results in a large amount of light that’s refracted, giving us sparkly diamond rings.

d

Graphite has a completely different crystal structure than diamond. Each unit cell is composed of layers of hexagonally arranged carbon atoms (see below “b”). Within each layer, each carbon atom is covalently bonded to three neighboring carbon atoms (i.e. in the same layer). On the other hand, each carbon atoms is bonded to a fourth carbon atom with secondary bonds in between layers. The weaker secondary bonds between layers result in the softer properties of graphite.

(a) diamond unit cell and (b) graphite unit cell where black arrows indicated strong covalent bonds and red arrows indicate weaker secondary bonds (modified from [2]).

In end, the reasons why we wear diamond jewelry and not and graphite jewelry all have to do the structure and the resulting properties.

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[1] W.D. Callister, Jr. and D.G. Rethwisch, Fundamentals of Materials Science and Engineering, NJ: John Wiley & Sons, Inc., 2008.

[2] P. Hofmann. (2009 Mar 14). Crystal Structures [Online]. Available: http://users-phys.au.dk/philip/pictures/physicsfigures/physicsfigures.html

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