This is part of preliminary HSC Chemistry course under the topic of Bonding.

HSC Chemistry Syllabus

    • Investigate elements that possess the physical property of allotropy

        Allotropy: What Are Allotropes?

          What is Allotropy?

          Allotropy refers to the phenomenon in which an element can exist in more than one distinct crystalline form, known as allotropes. Allotropes are manifestations of a single element in the same physical state (solid, liquid, or gas) that showcase significantly different physical properties such as density, hardness, electrical conductivity, and colour. Elements such as carbon, sulfur, phosphorous, and oxygen exhibit allotropy.Allotropes typically emerge due to the differences in stability at specific temperature and pressure conditions. In other words, certain forms of an element may be more stable under a particular set of environmental conditions, leading to the formation of different allotropes.

          Allotropes of Carbon 


          Diamond is one of the allotropes of carbon. Diamond has a robust covalent network structure which comprises a limitless matrix of carbon atoms. The specific alignment of the carbon atoms' layers plays a pivotal role in defining the unique properties of diamond. 
          Structure of diamond 
          In a diamond, each carbon atom is covalently bonded to four other carbon atoms at the vertices of a tetrahedron. This strategic formation imparts extraordinary strength to the structure, making it nearly impervious to breakage. Consequently, diamonds possess a remarkable hardness, elevated melting points, high boiling points, and display low reactivity.Interestingly, the tetrahedral configuration of carbon atoms creates gaps within its structure, allowing light to propagate within with least resistance. Diamond's refractive quality is enhanced when diamonds are cut in a way that allows light to enter and then be reflected in several directions.

          Properties of Diamond
          Property Reasoning
          Hard The three-dimensional lattice structure formed from a repetition of tetrahedral structures connected by strong covalent bonds between the carbon atoms make it incredibly difficult to break
          High M.P and B.P Because of the strength of the intermolecular bonds (bonds within molecules) between the carbon atoms, there must be lots of energy required to break them
          Transparency The carbon atoms are arranged in a very orderly fashion formed by tetrahedral structures with large spaces between them. Because of this, light is able to pass between the atoms to give diamond its colourless appearance and high light-refractive index.
          Not electrically conductive Since diamond is a covalent network substance, this means that the bonds between the carbon atoms are covalent bonds and the valence electrons of the carbon atoms are not free to move.
          High Density The structure of the diamond shows that the atoms are tightly bound in a strong three dimensional lattice structure.
          Insoluble Diamond is extremely hard and unreactive due to the covalent network structure and a lack of free valence electrons. Thus there is no substance that will be able to chemically react with diamond.
          Excellent Thermal Conductivity Since the carbon atoms in diamond are strongly bonded via covalent bonds which are precisely aligned, the diamond is known as an ideal crystal. The atoms in the crystal lattices of solids vibrate which allow for thermal conduction but since the carbon atoms in diamonds are so ideally constructed, they do not interact with one another.
            Uses of Diamond
          Use Relevant property
          Jewellery Lustrous, High light refractive index, hard, scratch resistant, transparent
          Industrial cutting tools and drills Because diamond is the hardest substance on earth and has a high melting point, this means that it can cut through anything without melting
          Heat sinks Because Of the high thermal conductivity of diamond, it is able to quickly extract heat from any sensitive areas such as computer chips.


          Structure of graphite 
          Graphite's structure significantly different to that of diamond, which in turn substantially influences its distinct physical and chemical properties. Graphite also adopts a covalent network structure consisting of layers of carbon held together by dispersion forces. In each layer, each carbon atom in graphite is covalently bonded to three other carbon atoms, unlike the four in diamond. This allows the carbon atoms in each layer to form six-membered carbon rings.Intriguingly, one electron from each carbon atom is not engaged in a covalent bond. These surplus valence electrons combine to form a sea of delocalised electrons. This sea of electrons grants graphite its electrical conductivity, a distinctive property of graphite that is not seen in diamond.

          Properties of Graphite 

          Property Reasoning
          Good Lubricant Because the carbon atoms in graphite are connected in hexagonal rings which stack in sheets which are held together by weak dispersion forces, the layers can easily slip past each other to give graphite its slippery nature. This is the same reason why graphite is used in pencils as this property allows the substance to mark paper when pressure is applied.
          Extremely Soft For the same reason graphite is a good lubricant, the weak dispersion forces between the layers means that the layers of carbon easily slide over each other to make it soft and greasy.
          Medium Density Graphite’s density is less than that of diamond. This is due to the structural layout. The layers of the graphite are separated by large distances since the forces which hold the layers together are weak and unable to bind them together tightly.
          High M.P and B.P Graphite can still be considered as a covalent network substance despite not having any bonding in the vertical direction between layers. The carbon atoms are connected via strong covalent bonds which extend throughout the horizontal lattice. These intramolecular bonds (bonds between multiple molecules) are hard to break and thus more emergy is required to break them.
          Good Electrical Conductor As previously explained, because each carbon atom in graphite is bound to only three other atoms, there is one single electron that is not covalently bound for each of them. These extra valence electrons form a sea of delocalized electrons which helps to make it an excellent electrical conductor. However electricity is only conducted along the plane of the layers so at 90 degrees it is unable to conduct.  

            Uses of Graphite 

          Use Relevant property
          Lead Pencils Soft and slippery with layers easily separated
          Electrodes Good electrical conductivity
          Polishes and paints Soft, slippery nature, metallic luster
          Lubricant in machines Slippery nature and high melting point
          Dry cell battery Good electrical conductivity, high melting point


          Allotropes of Oxygen

          Both allotropes of oxygen are covalent molecules.

          Structure of dioxygen and ozone

          Dioxygen (O₂)

          Also known as oxygen gas. This is the most common allotrope of oxygen and the form that is essential for the respiration of most life forms on Earth. It consists of two oxygen atoms bonded together, making up about 21% of the Earth's atmosphere. Dioxygen is a colourless, odourless, and tasteless gas at room temperature.

          Ozone (O₃)

          Ozone is a less stable allotrope of oxygen consisting of three oxygen atoms. It has a distinctive sharp smell and is pale blue in color. Ozone is present in the Earth's stratosphere, where it forms the ozone layer that protects living organisms by blocking harmful ultraviolet radiation from the sun. At ground level, ozone is a harmful pollutant and a significant component of smog.

          Allotropes of Phosphorus

          Structure of phosphorus allotropes

          White Phosphorus (P₄)

          White phosphorus consists of P₄ tetrahedra, where each phosphorus atom is bonded to three other phosphorus atoms. It is a highly reactive, waxy, white solid that glows in the dark when exposed to oxygen (chemiluminescence) and is stored under water to prevent spontaneous combustion. White phosphorus is used in the production of phosphoric acid, certain phosphorus compounds, and has applications in military incendiaries.

          Red Phosphorus

          Red phosphorus is an amorphous network of phosphorus atoms. It is more stable and less reactive than white phosphorus. Red phosphorus is used in safety matches, fireworks, smoke bombs, and as a flame retardant in plastics and textiles. Unlike white phosphorus, it does not glow in the dark and is not as dangerous.

          Black Phosphorus

          Black phosphorus has a layered structure similar to graphite, with individual layers of phosphorus atoms connected in a hexagonal pattern. This allotrope is the least reactive form of phosphorus and has good electrical conductivity. Its properties have gained interest for potential applications in electronics and materials science, particularly in semiconductors and in the emerging field of phosphorene, a single layer of black phosphorus.