Carbon And its Compounds-B

CloseHelium is a noble gas with a completely filled outer shell (2 electrons). It is chemically inert and does not form bonds with any element, including carbon. Oxygen, nitrogen, and hydrogen readily form covalent compounds with carbon.

CloseUnsaturated organic compounds have at least one carbon-carbon double bond (alkenes) or triple bond (alkynes). These multiple bonds allow addition reactions. Saturated compounds (alkanes) contain only single bonds.

CloseHydrogen has one electron in its outer shell (1s¹). It needs one more electron to achieve the stable duplet configuration of helium, so its valency is 1. It forms one bond (e.g., H–H, H–Cl, H–O–H).

CloseCovalent bonds are strong intramolecular forces that hold atoms together within a molecule (e.g., H–H bond energy 436 kJ/mol). The forces between molecules (intermolecular) are weak, which is why covalent compounds have low melting points.

CloseCatenation is the unique ability of an element to form covalent bonds with itself, creating long chains, branched chains, or rings. Carbon shows maximum catenation due to its small size and strong C–C bond strength.

CloseGraphite has a layered structure where each carbon atom is bonded to three others in a hexagonal plane. These planar layers are stacked one above the other with weak van der Waals forces between them, allowing layers to slide easily.

CloseSilicon can form Si–Si bonds, but these are much weaker than C–C bonds (226 kJ/mol vs. 348 kJ/mol). Silicon compounds (silanes) are highly reactive, especially with oxygen and water, unlike stable carbon compounds.

CloseCarbon’s small atomic radius allows its valence electrons (4) to be held tightly by the nucleus, enabling strong, stable covalent bonds. Larger atoms like silicon have more diffuse orbitals, leading to weaker bonds and less stable compounds.

CloseMethane (CH₄) has a central carbon atom sharing one pair of electrons with each of four hydrogen atoms. Carbon’s tetravalency (four bonds) is fully satisfied by four hydrogens in a tetrahedral geometry.

CloseThe carbon-carbon single bond has a bond energy of about 348 kJ/mol, and C=C double bond about 614 kJ/mol. This strength, combined with carbon’s small size, allows long stable chains—the basis of organic chemistry.

CloseStrong C–C and C–H bonds enable carbon to form millions of stable compounds with varied structures (chains, rings, branches, multiple bonds). Abundance alone is insufficient; silicon is abundant but forms far fewer stable compounds.

CloseIn graphite, each carbon uses three of its four valence electrons for bonding, leaving one electron delocalized across each hexagonal layer. These mobile electrons can move freely when a voltage is applied, conducting electricity.

CloseFullerenes are distinct allotropes of pure carbon (like diamond and graphite). They consist of carbon atoms arranged in closed cages or tubes (e.g., C₆₀ buckyball, carbon nanotubes), discovered in 1985.

CloseMethane (CH₄) is the simplest hydrocarbon, containing carbon and hydrogen. It is the primary component of natural gas. It contains no oxygen, nitrogen, or sulphur.

CloseNitrogen has five valence electrons. In NH₃, it shares three electrons (one with each hydrogen atom) to complete its octet (eight electrons). This is covalent bonding, not electron gain (which would form N³⁻, rare) or loss.

CloseIn graphite, each carbon forms three sigma bonds with three neighboring carbons in a hexagonal planar sheet. The fourth valence electron is delocalized above and below the plane.

CloseSynthetic diamonds are made by subjecting graphite or pure carbon to extremely high pressure (50,000–100,000 atm) and high temperature (1,500–2,000°C), mimicking natural diamond formation conditions deep in the Earth.

CloseIn diamond, each carbon atom uses all four valence electrons to form four strong single covalent bonds with four neighboring carbon atoms in a tetrahedral 3D network. This gives diamond its extreme hardness.

CloseIn N₂, each nitrogen atom has five valence electrons. To form the triple bond, each atom contributes three electrons (sharing three pairs). The remaining two electrons on each atom remain as a lone pair.

CloseSynthetic diamonds have the same crystal structure, chemical composition (pure carbon), hardness, and optical properties as natural diamonds. Experts need specialized equipment to tell them apart; chemically they are identical.

CloseGraphite feels slippery because the weak van der Waals forces between its hexagonal layers allow them to slide over each other easily. This property makes it useful as a lubricant and in pencil “lead.”

CloseThe first fullerene discovered (1985, Kroto, Smalley, Curl) was buckminsterfullerene C₆₀, a hollow cage of 60 carbon atoms arranged in 12 pentagons and 20 hexagons, resembling a soccer ball.

CloseDiamond has a perfect tetrahedral 3D covalent network with all bonds equally strong. It scores 10 on the Mohs hardness scale, making it the hardest naturally occurring substance known.

CloseA covalent bond is defined as the mutual sharing of one or more pairs of electrons between two atoms. This contrasts with ionic bonds (electron transfer) and metallic bonds (electron sea).

CloseAmmonia has one nitrogen atom covalently bonded to three hydrogen atoms. Nitrogen contributes three electrons (one per H), and each hydrogen contributes one electron, forming three single covalent bonds.

CloseIn graphite, carbon atoms are arranged in planar hexagonal (benzene-like) rings that form continuous sheets. Each hexagon has alternating single/double bonds due to delocalization of the fourth electron.

CloseCarbon dioxide contains one carbon atom double-bonded to two oxygen atoms (O=C=O). Carbon completes its octet (four bonds), and each oxygen completes its octet (two bonds plus two lone pairs).

CloseBiogas (from organic waste decomposition) contains about 50–70% methane. CNG (Compressed Natural Gas) is mainly methane (85–95%). LPG is propane/butane, coal gas contains H₂ & CO, producer gas contains CO & N₂.

CloseThe most stable form of elemental sulfur at room temperature is S₈, a puckered ring of eight sulfur atoms. This is different from carbon, oxygen (O₂), and hydrogen (H₂).

CloseBuckminsterfullerene (C₆₀) has a shape exactly like a soccer ball (football) with 20 hexagons and 12 pentagons. It is not a perfect sphere (cube) or simple ring.

CloseIn N₂, the two nitrogen atoms share three pairs of electrons (six electrons total), forming a very strong triple bond (N≡N) with bond energy 941 kJ/mol, making N₂ chemically inert.

CloseIn NH₃, there are three N–H single covalent bonds. Nitrogen also has one lone pair of electrons. There are no double or triple bonds, and the bonds are covalent, not ionic.

CloseCarbon’s tetravalency (four bonds) allows it to bond with many elements, while catenation allows it to bond with itself to form chains, branches, and rings. Together, these give millions of carbon compounds.

CloseCarbon has four valence electrons (2s²2p²). It needs four more electrons to complete its octet. It achieves this by forming four covalent bonds, giving it a valency of 4.

CloseBetween separate covalent molecules, forces are London dispersion, dipole-dipole, or hydrogen bonds—all weaker than covalent or ionic bonds. This explains low melting/boiling points and volatility.

CloseTo melt or boil a covalent compound, only the weak forces between molecules need to be overcome, not the strong covalent bonds within molecules. Little energy is required, so melting/boiling points are low.

CloseSaturated hydrocarbons (alkanes) contain only carbon-carbon single bonds. Each carbon is bonded to the maximum possible number of hydrogens. Double or triple bonds indicate unsaturation.

CloseAllotropes are different structural forms of the same element in the same physical state. Diamond, graphite, graphene, and fullerenes are allotropes of carbon. Isomers are different compounds with the same molecular formula.

CloseCarbon’s small size and 6 protons create a strong positive nucleus that pulls shared electron pairs tightly. This strong electrostatic attraction gives C–C and C–H bonds high bond energies (348–435 kJ/mol).

CloseMethane has one carbon atom and four hydrogen atoms. CH₂ is methylene (unstable), C₂H₆ is ethane, and CH₃ is a methyl group (not a stable compound alone).

CloseBoth are pure carbon (same valency, atomic number). Diamond has sp³ hybridized carbons in a 3D tetrahedral network; graphite has sp² hybridized carbons in planar hexagonal layers with delocalized electrons.

CloseOver 3 million organic compounds (carbon-containing) have been identified and characterized. All other elements combined account for about 100,000 compounds. This highlights carbon’s unique versatility.

CloseCarbon forms strong, stable covalent bonds with itself and other elements (H, O, N). This stability allows complex organic molecules to exist without decomposing easily at normal temperatures.

CloseBuckminsterfullerene (C₆₀) was named after architect Richard Buckminster Fuller, who designed geodesic domes that resemble the C₆₀ cage structure. The discoverers saw the similarity and used his name.

CloseCovalent compounds do not dissociate into ions in solution or melt. Without mobile ions or free electrons (except graphite), they cannot conduct electricity. Being solid or low density is not the reason.

CloseTetravalent means it can form four covalent bonds. Carbon has exactly four valence electrons (2s²2p²) in its outer shell, so it can share these four electrons with other atoms.

CloseLarger atoms have valence electrons farther from the nucleus, shielded by more inner shells. The nucleus attracts shared electron pairs less strongly, leading to weaker covalent bonds (e.g., Si–Si weaker than C–C).

CloseDiamond has a rigid, three-dimensional (3D) tetrahedral network where each carbon is bonded to four others. This 3D structure is responsible for diamond’s extreme hardness and high thermal conductivity.

CloseA triple bond involves sharing of three pairs of electrons (six electrons total). For example, in N₂ (N≡N) or C₂H₂ (acetylene, HC≡CH). One shared pair = single bond, two pairs = double bond.