Molecular Variety: Allotropes of Carbon and Phosphorus Quiz

Each carbon atom in diamond is covalently bonded to four other carbon atoms in a rigid tetrahedral framework. This three-dimensional arrangement involves sp3 hybridization, creating an extremely stable and hard structure. Because all valence electrons are tightly held in localized sigma bonds, diamond does not conduct electricity but possesses exceptional thermal conductivity and physical durability.

Graphite consists of planar hexagonal layers. While atoms within a layer are strongly bonded, the forces between the layers are weak Van der Waals attractions. This allows the sheets to slide over one another with minimal friction. This lubricating property is inherent to its structural organization and remains stable even under conditions where liquid oils might fail.

This is false because white phosphorus is actually the most reactive and unstable allotrope. It consists of P4 tetrahedra with 60-degree bond angles, creating intense ring strain. This strain makes it highly susceptible to oxidation, often leading to spontaneous combustion in air. Red phosphorus is polymeric and lacks this strain, making it much more chemically stable.

Fullerenes are unique because they exist as discrete molecules rather than infinite networks. The C60 molecule features a combination of pentagonal and hexagonal rings to create its curvature. Unlike diamond or graphite, these individual molecules can dissolve in non-polar solvents, which is a critical distinction when studying the solubility trends of various p-block elemental forms.

In graphite, each carbon atom uses only three of its four valence electrons for covalent bonding. The fourth electron remains in a p-orbital and becomes delocalized across the hexagonal layer. These mobile electrons move freely when an electrical potential is applied, whereas in diamond, all electrons are locked in single bonds, preventing any flow of charge.

Heating white phosphorus provides the energy required to break the strained P-P bonds in the P4 tetrahedra. The atoms then link up into long, stable chains to form red phosphorus. This transition from a molecular solid to a polymeric solid significantly alters the physical properties, resulting in a higher melting point and much lower chemical reactivity.

Black phosphorus is formed under high pressure and is the most stable form of the element. Its structure consists of wavy, puckered layers of phosphorus atoms. Similar to graphite, this layered arrangement gives it interesting electronic properties, including semiconductivity. Its stability is a result of the optimized bond angles and distances that minimize the internal energy of the crystal.

Graphene is a single layer of carbon atoms arranged in a honeycomb lattice. It is essentially one layer of graphite. Because of the strength of the sp2 covalent bonds and its unique thickness, it is one of the strongest materials known. Its two-dimensional nature leads to exceptional mechanical, thermal, and electronic properties used in advanced materials engineering.

Diamond hardness stems from its infinite three-dimensional lattice where every bond is a strong C-C covalent bond. There are no weak planes or forces within the crystal. Breaking or displacing any part of the structure requires overcoming the high dissociation energy of multiple covalent bonds simultaneously, which is why it remains the benchmark for hardness in inorganic chemistry.

Red phosphorus is used for matches because it is stable at room temperature and non-toxic compared to the white variety. When friction is applied, a small amount of red phosphorus converts to white phosphorus, which then ignites to start the flame. This controlled use of allotropic transformation is a practical application of the stability differences found in phosphorus chemistry.

Carbon nanotubes are essentially rolled-up sheets of graphene. Therefore, the carbon atoms maintain their sp2 hybridization. This allows for the formation of a hexagonal network with delocalized pi electrons. The resulting structure is incredibly strong and conductive, showcasing how the geometry of p-orbitals can be manipulated into different shapes like tubes, spheres, or sheets.

This is false. Diamond is significantly denser than graphite. In diamond, the atoms are closely packed in a rigid 3D lattice with a bond length of 154 pm. In graphite, while atoms are close within a layer, the large gaps between the layers (335 pm) result in a more open and less dense overall volume compared to the compact diamond crystal.

White phosphorus consists of non-polar P4 molecules. Based on the principle of like dissolves like, it is soluble in non-polar organic solvents such as carbon disulfide. Red and black phosphorus are polymeric or network solids; because they are held together by extensive covalent systems, they do not dissolve in common solvents, illustrating a key difference between molecular and polymeric allotropes.

Diamond and graphite are classic examples of giant covalent lattices where the bonding extends throughout the entire sample. Red phosphorus is also a polymeric network. Only white phosphorus exists as discrete P4 molecules. Distinguishing between molecular and network structures is essential for predicting properties like melting points, volatility, and solubility in different chemical substances within the p-block.

White phosphorus undergoes slow oxidation when exposed to moist air, a process that releases energy in the form of a faint green-white glow. This chemiluminescence is a famous characteristic of the P4 allotrope. It highlights the high reactivity of the phosphorus atoms when arranged in a strained tetrahedral molecular form, leading to a steady chemical reaction at the surface.