Alkanes & Cycloalkanes

Looking down the C–C bond of ethane in a Newman projection, you see the front carbon’s three H atoms as three spokes of a wheel. The back carbon’s three H atoms are seen at the edges of the circle. Now rotate the back carbon. Two extremes emerge: eclipsed, where the back H atoms line up directly behind the front ones (0° offset), and staggered, where the back H atoms fall exactly between the front ones (60° offset). The eclipsed form is higher energy because the closely aligned electron clouds of the C–H bonds repel each other — this is called torsional strain. Think of it like two fans spinning in opposite directions: when the blades line up, they clash; when they’re offset, they pass each other freely.

The energy difference (~12 kJ/mol) is the activation barrier to rotation. At room temperature, thermal energy (~84 kJ/mol) easily clears the barrier → rotation is fast, ~10¹⁰ times/s.

Butane (CH₃–CH₂–CH₂–CH₃) is more interesting because the substituents being rotated are now methyl groups, not just H atoms. Methyl groups are much bulkier than H, so the energy differences between conformations are larger and more clinically relevant. Looking down the C2–C3 bond: when the two methyls are 180° apart (anti), they are as far apart as possible and the molecule is at its energy minimum. This anti conformation is why long-chain alkanes in solution and in the extended fatty-acid tails of phospholipids adopt a stretched-out zig-zag shape — each internal C–C bond sits in the anti arrangement to minimise energy.

The first four prefixes (meth-, eth-, prop-, but-) are irregular and must simply be memorised — they derive from the historical names of methanol, ether, propionic acid and butyric acid (butter fat). From C5 onwards the prefixes follow Greek and Latin number roots: penta (5), hexa (6), hepta (7), octa (8), nona (9), deca (10). Once you know these, every alkane name from methane to decane is trivially constructed: prefix + "-ane".

All alkane names end in the suffix "-ane". Alkenes use "-ene", alkynes "-yne".

An alkyl group is an alkane molecule that has lost one hydrogen atom — the open bond is where it attaches to the main chain as a side branch. The naming rule is simple: replace the “-ane” suffix with “-yl”. So methane → methyl (–CH₃), ethane → ethyl (–CH₂CH₃), and so on. The key complication is that propane and butane have more than one possible position to remove the H from, giving rise to isopropyl vs n-propyl, and the four butyl variants. These come up in naming problems every year.

The 1°/2°/3°/4° classification of carbons runs throughout all of organic chemistry — not just here. A primary (1°) carbon touches only one other carbon; a secondary (2°) touches two; a tertiary (3°) touches three; a quaternary (4°) touches four. The degree determines not just nomenclature but reactivity: tertiary radicals are more stable than primary ones (this chapter), tertiary carbocations are more stable than primary ones (future chapters), and the degree of the carbon bearing the leaving group in substitution reactions determines whether the reaction goes by SN1 or SN2 (future chapters). Learn it now and it pays dividends forever.

An alkyl group (R) is an alkane minus one H — what dangles off the main chain as a "substituent". Replace -ane with -yl:

This 1°/2°/3°/4° classification matters all through OC — it predicts radical stability (this chapter), alcohol classification (Ch 4) and SN1/SN2 selectivity in advanced courses.

Follow these six steps in order for every naming problem. The most common error students make is step 1 — they do not find the longest chain, they find a long chain. The longest chain may not be drawn horizontally across the page. Train yourself to trace every possible chain before committing to the parent. Once the parent is identified, the rest flows mechanically.

Other substituents you’ll encounter: –F fluoro-, –Cl chloro-, –Br bromo-, –I iodo-, –NO₂ nitro-, –NH₂ amino-.

Combustion is an alkane’s most important reaction for human civilization. When an alkane burns in excess oxygen, the C–C and C–H bonds break and the carbon atoms are completely oxidised to CO₂ while hydrogen becomes water. The energy released is the same energy that drives your car, heats your home, and powers the aircraft above your head. In medicine, incomplete combustion produces carbon monoxide (CO), which binds haemoglobin 200 times more tightly than O₂ — a fact that makes CO poisoning so dangerous. CO is odourless and displaces oxygen from haemoglobin without any warning sensation.

Complete combustion gives CO₂ + H₂O + heat. Three benchmark equations from the slide:

When an alkane is exposed to UV light or heated to ~300 °C in the presence of Cl₂ or Br₂, a C–H bond is replaced by a C–X bond. This radical halogenation proceeds through a chain reaction with three phases: initiation, propagation, and termination. The key insight is that the reaction is sustained without the need for UV or heat once it starts, because the chain propagates itself — each propagation step produces a new radical that continues the reaction. This same radical chain mechanism, when it runs inside the body’s lipid membranes (initiated by reactive oxygen species), causes oxidative damage to DNA, proteins, and cell membranes — the molecular basis of inflammation and aging.

A crucial selectivity principle emerges from radical halogenation: when a molecule has multiple types of C–H bonds, the reaction preferentially abstracts the H from the most substituted carbon. Tertiary C–H bonds (at 3° carbons) are broken most easily, primary C–H bonds least easily. This is because the resulting tertiary radical is the most stable, thanks to hyperconjugation — neighbouring C–H bond electrons donate into the half-filled p orbital of the radical carbon, stabilising it. More neighbouring C–H bonds (more substitution) means more hyperconjugation, means more stable radical.

An alkyl radical is itself sp² hybridised (nearly planar) with the unpaired electron in a p orbital perpendicular to the plane — the same geometry as a carbocation.

When two substituents are attached to different carbons of a ring, they can both be on the same face of the ring (cis) or on opposite faces (trans). Unlike the cis/trans of open-chain alkenes (where a π bond prevents rotation), the cis/trans on a ring is locked simply because you would have to break a bond to interconvert them — there is no free rotation of the ring backbone without first snapping one of the ring bonds. This makes ring cis/trans a true stereoisomerism, not just conformational difference. In drug design, the cis and trans isomers of a ring-containing drug can have completely different biological activities because they present their functional groups in different spatial orientations to the receptor.

Two substituents on different ring carbons can sit on the same face (cis) or on opposite faces (trans). Because ring C–C bonds cannot rotate freely, cis and trans ring isomers are non-interconvertible stereoisomers.

If both substituents are on the same carbon (e.g. 1-Br-1-Cl-cyclobutane), there is no cis/trans: only a single constitutional isomer.

Adolf von Baeyer made a bold but ultimately partly-wrong prediction in 1885. He assumed all cycloalkanes were flat (planar), then calculated how much the ring angle would have to deviate from the ideal tetrahedral 109.5°. For cyclopropane, a flat three-membered ring forces a 60° angle — a huge deviation — so cyclopropane should be very strained. For cyclohexane, a flat ring would require 120° bond angles. That’s only 10° from 109.5°, so Baeyer predicted cyclohexane would be essentially strain-free. So far so good. But Baeyer also predicted that large rings (7, 8, 9 members) would also be strained because their flat-ring angles would exceed 109.5°. Here he was wrong: large rings simply adopt non-planar conformations that restore the ideal 109.5° angles at every carbon. The key insight is that cyclohexane is strain-free not because its flat angle is close to 109.5°, but because it folds into a chair conformation.

• Cyclopropane: C–C–C angle 60° (real measured ~104° for the "bent bonds"), far from 109.5° → high bond-angle strain + all H's eclipsed → torsional strain. Hence cyclopropane is unusually reactive for a saturated compound.
• Cyclobutane: less strain than cyclopropane; slightly puckered.
• Cyclohexane: strain-free — adopts a non-planar chair with 111° bond angles and all bonds staggered.

The chair conformation of cyclohexane is one of the most important shapes in all of chemistry. Every carbon in the chair sits at a perfect 111° bond angle with all neighbouring bonds staggered — there is literally zero torsional strain and essentially zero angle strain. This is why cyclohexane is thermodynamically much more stable than the boat conformation (which is ~30 kJ/mol less stable due to eclipsing and steric clashing of the “flagpole” hydrogens). The two positions on each carbon in the chair have distinct names and distinct energies: axial positions point straight up or straight down (parallel to the ring axis), while equatorial positions splay outward around the equator of the ring.

Here is the rule that the exam tests every single year: a substituent prefers the equatorial position. In the axial position, a substituent points directly at the axial hydrogens on carbons 3 and 5 above it — a clash called 1,3-diaxial repulsion. In the equatorial position, it points away from the ring into space with nothing to clash against. The bigger the substituent, the more it wants equatorial. For tert-butyl (–C(CH₃)₃), the equatorial preference is so strong that essentially 100% of molecules adopt the equatorial chair at room temperature. This rigid-equatorial geometry is the basis of how glucose sits in its pyranose ring, and how cholesterol’s angular methyl groups sit in their axial positions on the steroid scaffold — these shapes determine biological function.

In the chair, each carbon carries one axial H (perpendicular to the average ring plane) and one equatorial H (sticking out around the equator at ~109.5° from axial). The 6 axial and 6 equatorial positions alternate around the ring.