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Cyclohexane is the cyclic alkane ${\mathrm{C}}_6{\mathrm{H}}_{12}$. Its preferred chair conformation allows near-tetrahedral bond angles and largely staggered neighboring bonds.

Ring flips exchange axial and equatorial positions. Consequently a substituent can retain its up/down orientation while changing the steric interactions that determine conformational preference and reaction accessibility.

Cyclopentane is the cyclic alkane ${\mathrm{C}}_5{\mathrm{H}}_{10}$. Its five-membered ring puckers rather than remaining completely planar, reducing unfavorable eclipsing interactions among neighboring bonds.

Envelope-like conformations interconvert through ring motion. This illustrates why simply drawing a regular polygon does not specify the lowest-energy geometry of a saturated ring.

Ethane is the alkane ${\mathrm{C}}_2{\mathrm{H}}_6$, containing two tetrahedral carbon centers joined by a single bond. Rotation about that bond connects staggered and eclipsed conformations.

The staggered arrangement is lower in energy, so ethane provides a basic example of torsional preferences in a molecule whose connectivity remains unchanged. The rotational barrier is finite: a single-bond drawing does not mean every torsion angle has the same energy.

Hexane is an alkane with formula ${\mathrm{C}}_6{\mathrm{H}}_{14}$; the unqualified laboratory name often refers to the straight-chain isomer, n-hexane. Other constitutional isomers share the same formula.

Its mostly nonpolar molecular surface explains its use as a solvent for many hydrophobic substances. Commercial mixtures called hexanes need not be a chemically pure single isomer, so the material name and molecular identity should not be treated as interchangeable.

Methane is $\mathrm{C}{\mathrm{H}}_4$, the simplest alkane. Its carbon forms four approximately tetrahedral carbon-hydrogen bonds, giving a symmetric molecule without a permanent dipole.

It is a major component of natural gas and can be oxidized to carbon dioxide and water. Its infrared-active vibrations also make it relevant to atmospheric radiation. The molecular formula specifies a single carbon center, unlike longer alkanes that can have multiple carbon-skeleton isomers.

Transfer is favourable when new water–solute interactions compensate the cavity and reorganization costs of accommodating the solute. The Gibbs energy also includes the loss of solute–alkane interactions; it is not simply a count of hydrogen bonds. isobutane cannot compensate those costs with strong directional interactions, explaining its positive transfer Gibbs energy. acetone and both amines can accept hydrogen bonds from water and have favourable electrostatic solvation, making their transfer moderately favourable. isopropylamine can additionally donate to water, consistent with a somewhat greater preference for water than trimethylamine.

acetamide combines a strong carbonyl acceptor with N–H donors and has the most favourable hydration in the group, accounting for its much more negative transfer Gibbs energy. Its pure solid lattice is not an extra sublimation penalty to insert here: the stated process starts from a dilute alkane solution and ends in dilute water, rather than dissolving the pure crystal. The nonmonotonic relationship between hydrogen-bond acceptor parameter and transfer Gibbs energy is expected because donor strength, nonpolar surface, solvent cavity formation, and all other solvation contributions also enter.

For matched standard states, $\Delta G_{\mathrm{tr}}^{\circ }=-RT\ln K_{\mathrm{tr}}$, with $K_{\mathrm{tr}}$ the water-to-alkane activity ratio. As an illustrative 298 K calculation, the hydrocarbon value implies $K_{\mathrm{tr}}\simeq 7.0\times 10^{-4}$, whereas the amide value implies approximately $3.6\times 10^4$. These are standard-state partition preferences, not absolute aqueous solubilities. amine protonation could introduce an additional pH-dependent contribution to an experimentally measured distribution coefficient, but no such speciation information is supplied and it is not needed for the molecular comparison.

The wick supplies liquid wax by capillary action; wax vapour reacts with atmospheric oxygen, releasing the heat that keeps the supply going. The wick is a delivery system, rather than the principal fuel.

First, an external flame warms the wick and adjacent wax. Melting produces a small liquid reservoir, and capillary action draws it into the spaces between the wick fibres. Near the hot tip, evaporation and some pyrolysis supply gaseous fuel. Melting and evaporation are physical changes; pyrolysis and combustion change the chemical substances present. Once established, the candle flame transfers enough heat back to the wax to maintain this sequence.^[ref]

The candle flame is a diffusion flame. Its relatively dark inner region contains newly supplied fuel with limited access to oxygen. In the yellow luminous region, incomplete conversion of the hydrocarbons produces small soot particles, which become incandescent. Farther out, mixing supplies enough oxygen to consume most of the remaining fuel, carbon monoxide and soot. Thus the yellow light is largely thermal emission from hot particles, rather than a yellow gas.^[ref]

For a representative alkane component of paraffin wax, complete combustion has the chemical equationReal candle materials are mixtures, and some, such as fatty materials, already contain oxygen; the displayed chemical equation is an illustrative wax model. The principal products of complete combustion are carbon dioxide and water vapour. Convection carries the hot gases upward and brings fresh air into the flame. A cold collector can reveal condensed water; collecting the gases and passing them into limewater gives a calcium carbonate cloudiness from carbon dioxide.

Removing the oxygen supply, cooling the reaction zone, or cutting off wax delivery extinguishes the flame. The smoking wick immediately afterwards can still emit combustible vapour; without sufficient heat at the mixing zone, that vapour need not sustain a flame.