Carbon and Its Compounds — Karnataka (SSLC) Class 10 Science Solutions (Free)

Carbon forms the basis of an extraordinary number of known compounds — far more than any other single element — because of five properties working together: 1. Tetravalency: carbon has exactly 4 electrons in its outer shell, letting it form 4 covalent bonds simultaneously, with hydrogen, oxygen, nitrogen, sulphur, or other carbon atoms. 2. Catenation: carbon atoms bond to EACH OTHER unusually readily and stably, forming long chains, branches, and rings — far more extensively than any other element can. 3. Polymorphism (allotropy): carbon exists in multiple pure structural forms — diamond (rigid 3D network), graphite (layered sheets), and buckminsterfullerene (hollow cage structures) — each with dramatically different physical properties despite being made of identical atoms. 4. Small atomic size: allows carbon's bonds (especially C–C) to be unusually strong, since smaller atoms allow closer, more effective orbital overlap when bonding. 5. Ability to form diverse compounds: the combination of the above lets carbon build millions of distinct organic compounds, from the simplest gas (methane) to the complex molecules (DNA, proteins) that life itself is built from. Together, these are exactly why organic chemistry — the entire branch of chemistry devoted to carbon compounds — exists as its own separate field, distinct from the chemistry of every other element combined.

Covalent bonding occurs when two atoms SHARE a pair of electrons to achieve a stable outer-shell configuration, rather than one atom fully giving up an electron to the other. Examples: C–H and C–C bonds in methane, O=O in oxygen gas, and the C=O bonds in carbon dioxide. Ionic bonding, by contrast, occurs when one atom fully TRANSFERS an electron to another, creating oppositely charged ions that then attract each other electrostatically. Example: in NaCl, sodium transfers an electron to chlorine, forming Na⁺ and Cl⁻ ions held together by the attraction between opposite charges. Key differences in behaviour: - Covalent compounds have relatively low melting and boiling points, and generally do NOT conduct electricity, since their electrons are held within localised, shared bonds rather than free to move as charge carriers. They can be non-polar (electrons shared evenly, as in symmetrical molecules) or polar (electrons shared unevenly, pulled toward the more electronegative atom). - Ionic compounds have high melting points (the electrostatic attraction between ions throughout the whole crystal lattice is strong), and DO conduct electricity — but only when molten or dissolved in water, where the ions become free to move and carry charge; a solid ionic crystal does not conduct, since its ions are locked in a fixed lattice. Both bonding types are ultimately driven by the same underlying goal — atoms achieving a stable, complete outer electron shell (the octet rule) — they simply achieve it by different mechanisms (sharing versus transferring).

Hydrocarbons are organic compounds made up of ONLY carbon and hydrogen atoms — the simplest category of organic compounds, and the building block from which more complex organic compounds (with other elements attached) are derived. Saturated hydrocarbons (alkanes): contain only single (C–C) bonds between carbon atoms, meaning every carbon has as many hydrogen atoms attached as it can possibly hold. General formula: CₙH₂ₙ₊₂. Example: methane, CH₄ (n=1); ethane, C₂H₆ (n=2). Unsaturated hydrocarbons: contain at least one double or triple bond between carbon atoms, meaning fewer hydrogen atoms are attached than the maximum possible. - Alkenes (one C=C double bond): general formula CₙH₂ₙ. Example: ethene, C₂H₄. - Alkynes (one C≡C triple bond): general formula CₙH₂ₙ₋₂. Example: ethyne (acetylene), C₂H₂. Why "unsaturated" compounds are more reactive: a double or triple bond represents extra, less stable electron density that is relatively easy for other reagents to attack and break open — this is why alkenes and alkynes readily undergo addition reactions (adding new atoms across the double/triple bond), while saturated alkanes, having no such multiple bond, are comparatively unreactive and mainly only undergo combustion or substitution reactions.

A functional group is a specific atom or group of atoms within a molecule that determines how that molecule chemically behaves — molecules sharing the same functional group tend to undergo similar chemical reactions, regardless of how different the rest of their carbon skeleton is. 1. Hydroxyl group (–OH): defines alcohols. Example: ethanol, C₂H₅OH — the alcohol found in alcoholic beverages, also used as a solvent and biofuel. 2. Carboxyl group (–COOH): defines carboxylic acids. Example: acetic acid, CH₃COOH — the acid responsible for vinegar's sour taste and characteristic smell. 3. Aldehyde group (–CHO): defines aldehydes. Example: formaldehyde, HCHO — used as a preservative and in the manufacture of certain plastics. Why functional groups matter: the identity of the functional group, not just the number of carbon atoms, largely determines a compound's physical properties (solubility in water, boiling point) and chemical reactivity. Two compounds built on the exact same carbon skeleton but carrying different functional groups (for example, ethanol with –OH versus acetic acid with –COOH, both based on a 2-carbon chain) behave completely differently — one is a mild solvent, the other a corrosive acid — which is exactly why functional groups, not just molecular formulas, are the primary way organic chemists classify and predict compound behaviour.

Isomerism is the phenomenon where two or more compounds share the exact same molecular formula, but have their atoms arranged in a different structure — because the structural arrangement differs, the compounds have different physical and sometimes chemical properties, despite technically being made of identical atoms in identical numbers. Example — C₄H₁₀ has two isomers: 1. n-Butane: CH₃–CH₂–CH₂–CH₃, a straight (unbranched) chain. Boiling point: about −0.5°C. 2. Isobutane (2-methylpropane): CH₃–CH(CH₃)–CH₃, a branched chain. Boiling point: about −12°C. Why the boiling points differ: Boiling point depends on how strongly neighbouring molecules attract each other (via weak intermolecular forces) — the more surface area two molecules have in close contact, the stronger this attraction, and the more energy (higher temperature) is needed to separate them into a gas. n-Butane's straight-chain shape allows neighbouring molecules to line up alongside each other with substantial surface contact, creating relatively stronger intermolecular attraction and therefore a HIGHER boiling point. Isobutane's branched, more compact, roughly spherical shape reduces the surface area available for close contact between molecules, weakening the intermolecular attraction and giving it a LOWER boiling point — this pattern (branched isomers boiling at a lower temperature than their straight-chain counterparts) holds generally across many hydrocarbon isomer pairs.

Saturated hydrocarbons (alkanes): general formula CₙH₂ₙ₊₂. Example: propane, C₃H₈ (n=3, giving 2(3)+2 = 8 hydrogens). Unsaturated hydrocarbons: - Alkenes (one C=C double bond): general formula CₙH₂ₙ. Example: propene, C₃H₆. - Alkynes (one C≡C triple bond): general formula CₙH₂ₙ₋₂. Example: propyne, C₃H₄. Why the hydrogen count differs: Each carbon atom always forms exactly 4 bonds in total (tetravalency never changes). In a saturated alkane, every one of carbon's bonds not used for the C–C backbone is used to attach a hydrogen atom — the maximum possible. Forming a double bond between two carbons uses up TWO of the bonding positions on each of those carbons for that one double bond (instead of one), which means each carbon involved has one fewer position left over for a hydrogen atom — removing 2 hydrogens overall compared to the fully saturated formula. A triple bond uses three bonding positions on each carbon involved for that single bond, removing 4 hydrogens overall compared to the saturated formula. This is precisely the mechanical reason CₙH₂ₙ₊₂ (saturated) steps down to CₙH₂ₙ (one double bond) and then CₙH₂ₙ₋₂ (one triple bond).

Combustion is the rapid oxidation of a hydrocarbon in the presence of oxygen, releasing energy as heat (and often light) — the exothermic reaction that makes hydrocarbon fuels useful for heating, cooking, and powering engines. Complete combustion of methane: CH₄ + 2O₂ → CO₂ + 2H₂O + heat Complete combustion of ethane: 2C₂H₆ + 7O₂ → 4CO₂ + 6H₂O + heat Complete vs incomplete combustion: Complete combustion occurs when there is a SUFFICIENT supply of oxygen, allowing every carbon atom in the fuel to be fully oxidised all the way to carbon dioxide, and every hydrogen atom to water — this is the clean-burning, most energy-efficient outcome. Incomplete combustion occurs when oxygen supply is LIMITED. With not enough oxygen available to fully oxidise every carbon atom, the reaction instead produces carbon monoxide (CO) — a toxic, colourless, odourless gas — and/or unburned carbon particles (soot), alongside less heat energy than complete combustion would release. This is exactly why gas appliances (like a stove or heater) need proper ventilation: a restricted air supply causes incomplete combustion, silently producing dangerous carbon monoxide even though the flame may still look like it's burning normally.

Balanced equation for the complete combustion of propane: C₃H₈ + 5O₂ → 3CO₂ + 4H₂O Verification of balance: Carbon: 3 (left) = 3 (right, in 3CO₂) ✓ Hydrogen: 8 (left) = 8 (right, in 4H₂O) ✓ Oxygen: 5 × 2 = 10 (left) = (3 × 2) + 4 = 6 + 4 = 10 (right) ✓ Oxygen needed for 2 moles of propane: From the balanced equation, 1 mole of C₃H₈ requires 5 moles of O₂. For 2 moles of C₃H₈: 2 × 5 = 10 moles of O₂ are required. This kind of mole-ratio calculation, read directly off a correctly balanced equation, is exactly how the amount of air/oxygen needed for efficient, complete combustion is calculated in real burner and engine design.