Carbon (Descriptive chemistry)

Covalent bonding — single, double, or triple bond

Diamond, graphite, graphene, fullerenes (e.g C₆₀ buckyballs), and carbon nanotubes.

Diamond has a 3D tetrahedral structure where each carbon is bonded to 4 others. It’s extremely hard, non-conductive, and has a high melting point.

Graphite has layers of carbon atoms arranged in hexagons. Each carbon is bonded to 3 others with delocalized electrons between layers → soft, slippery, and a good conductor of electricity (conducts electricity due to the delocalised 4th non-bonding C atom). Each sheet is held together by weak van der waals forces.

Acheson’s process. It requires electrical power. Sand (silicone dioxide) and coke (C) are maintained in a furnace for 24 hours. SiO2 + 3C = SiC + 2CO SiC = Si + C

When a substance has multiple allotropes and one is more stable than the others under standard conditions, such as diamonds degrading into graphite (under a catalyst, at very high temperature).

Charcoal is made by heating wood or sugar in a limited supply of Oxygen.

Coke is a by-product of the manufacturing of town-gas

Animal charcoal is made by burning animal bones and it consists of finely divided carbon on calcium phosphate.

It is obtained from burning hydrocarbons in the absence of Oxygen

It takes a lot of energy to force an atom to accept 4 electrons due to electron sheilding (electron weakens positive charge).

Lead, Germanium, Tin These elements are known to commonly exhibit a valence of 2 in their compounds.

The valence state The valence state determines how and why elements bond with others.

When 2s electrons become reluctant to bond so the other 2 electrons in np orbitals can bond (refers to group 14 heavy elements such as Tin,Lead and Germanium)

Catenation = ability to form chains with itself Decreases down group 14 because: – Bond strength C–C > Si–Si > Ge–Ge > Sn–Sn – Heavier atoms form weaker bonds → less stable chains – Higher bond enthalpy = more stable chains → carbon wins

High C–C bond energy (348 kJ/mol) Small atomic radius = strong overlap of orbitals No low-lying d-orbitals to destabilize bonding

C–H bond is stronger and shorter (C-C and C-H bonds have high bond energy) Si–H and Ge–H bonds are longer + weaker Larger atoms = poorer orbital overlap with hydrogen Carbon also doesn’t undergo orbital expansion

Heavier elements (Si, Ge) can use d-orbitals in bonding Leads to longer, weaker bonds → less stability Carbon lacks accessible d-orbitals → tighter, more stable bonds

Cyclohexane = saturated, single C–C bonds, no delocalization Benzene = planar, delocalized π system → extra stability Shows carbon’s versatility: can form both localized and delocalized structure

Stability decreases down the group:  C–F > C–Cl > C–Br > C–I Due to bond strength:  C–F (485 kJ/mol)»_space; C–Cl (327) > C–Br (285) > C–I (213) So fluorinated compounds are most thermally stable

Strongest single bond to carbon Gives compounds high thermal and chemical stability Found in Teflon, pharmaceuticals, refrigerants Resistant to metabolism — often makes drugs bioavailable & long-lasting

Halogens are electron-withdrawing → make C slightly δ⁺ Affects reactivity in nucleophilic substitution and elimination reactions More polar bonds = more reactive sites

Due to differences in intermolecular forces:  – CF₄ has weak London forces (small, non-polar molecule) → gas  – CI₄ has stronger London dispersion forces (large, polarizable iodine atoms) → solid Molecular size and electron cloud increase down the group → stronger van der Waals = higher melting/boiling points

The central carbon is shielded by 4 halogen atoms → steric hindrance Carbon is not electrophilic enough to attract water molecules Halogen atoms are not easily displaced without extreme conditions C–X bonds (especially C–F and C–Cl) are very strong and stable