Alkanes & Cycloalkanes

Alkanes & Cycloalkanes — IUPAC Nomenclature

Part 1 of 7 — Nomenclature

Alkanes are saturated hydrocarbons: every carbon is $sp^3$-hybridized and bonded to the maximum number of hydrogens, so the general formula for an acyclic (open-chain) alkane is $C_nH_{2n+2}$. A single ring removes two hydrogens, giving cycloalkanes the formula $C_nH_{2n}$.

A reliable, unambiguous name is not optional in organic chemistry — "hexane" and "2,2-dimethylbutane" share the formula $C_6H_{14}$ but are entirely different compounds (constitutional isomers). The IUPAC system encodes the exact connectivity of a molecule into its name, so that any chemist can redraw the structure from the name alone.

The unbranched (straight-chain) parents you must know cold:

The suffix -ane signals an alkane. The same numerical prefixes (meth-, eth-, prop-, …) reappear throughout organic chemistry attached to other suffixes (-ene, -yne, -ol), so memorizing them now pays off for the entire year.

The Five-Step IUPAC Algorithm

Naming a branched alkane is a deterministic procedure. Follow the steps in order and the name falls out:

1. Find the parent chain. Identify the longest continuous chain of carbons. This may not be the chain drawn horizontally — you must trace through bends. If two chains are tied in length, choose the one with more substituents.

2. Number the chain. Number the carbons consecutively, starting from the end that gives the lowest set of locants to the substituents (compare the first point of difference).

3. Identify and name the substituents. A substituent formed by removing one H from an alkane is an alkyl group, named by changing -ane to -yl: methyl ($-CH_3$), ethyl ($-CH_2CH_3$), propyl, etc.

4. Assemble the name alphabetically. List substituents in alphabetical order, each preceded by its locant. Use multiplying prefixes (di-, tri-, tetra-) for repeated groups — but these prefixes are ignored when alphabetizing (e.g. "diethyl" alphabetizes under e).

5. Punctuate. Numbers are separated from numbers by commas (2,3-) and from letters by hyphens (3-methyl). The complete name is written as one word ending in the parent.

Common alkyl substituents and their alphabetizing letter: isopropyl (under i), sec-butyl and tert-butyl (the italicized sec- and tert- prefixes are ignored, so both alphabetize under b), but isobutyl alphabetizes under i because "iso" is part of the name.

Worked Example — Naming a Branched Alkane

Name the compound whose skeleton is a 6-carbon chain bearing a methyl group on the third carbon and ethyl groups on the third and fourth carbons (drawn from the left):

$CH_3-CH_2-C(CH_3)(C_2H_5)-CH(C_2H_5)-CH_2-CH_3$

Step 1 — Parent chain. The longest continuous chain is 6 carbons long $\rightarrow$ hexane. (Check: tracing through the branch carbons does not yield anything longer than 6.)

Step 2 — Number for lowest locants. Numbering from the left places substituents at C3, C3, C4. Numbering from the right places them at C3, C4, C4. Compare first point of difference: $\{3,3,4\}$ vs $\{3,4,4\}$ — the left numbering wins because 3 < 4 at the second position.

Step 3 — Substituents. Two ethyl groups (C3 and C4) and one methyl group (C3).

Step 4 — Alphabetize. "Ethyl" (e) comes before "methyl" (m). The two ethyls are collected as diethyl, but we alphabetize using "ethyl," not "diethyl."

Step 5 — Assemble: $\;$ 3,4-diethyl-3-methylhexane.

Trap: A common mistake is to call this "3-methyl-3,4-diethylhexane" by alphabetizing on the prefix "di." The multiplying prefix di- is invisible to alphabetization — only the substituent root ("ethyl") counts, so ethyl precedes methyl.

Checkpoint — Parent Chain & Locants

Cycloalkanes

A cycloalkane is named by adding the prefix cyclo- to the parent name of the ring: a six-membered ring is cyclohexane, a three-membered ring is cyclopropane, and so on. The ring itself is usually taken as the parent.

Numbering a substituted ring: assign C1 to a substituted carbon, then number around the ring in the direction that gives the lowest set of locants to the remaining substituents. When substituents differ, the one first in alphabetical order receives the lower locant at the first point of difference.

• One substituent: no locant is needed (e.g. methylcyclohexane — the ring carbon bearing the methyl is "C1" by default).
• Two or more substituents: number to give the lowest locants, breaking ties alphabetically.

Ring vs. chain as parent: if the ring has more carbons than the largest attached chain, the ring is the parent (e.g. 1-ethyl-2-methylcyclohexane). If a chain attached to the ring is longer than the ring, the ring becomes a cycloalkyl substituent on that chain (e.g. a propyl chain bearing a cyclohexane ring is named cyclohexyl-substituted propane $\rightarrow$ "1-cyclohexylpropane" style naming, i.e. the ring is the substituent).

Worked check: A cyclopentane ring with a methyl on one carbon and a chlorine two carbons away. Start numbering at a substituted carbon and go the direction giving lowest locants. With substituents alphabetized (chloro before methyl), we get 1-chloro-3-methylcyclopentane rather than "3-chloro-1-methyl…", because the lower locant set is assigned to the first-cited (alphabetically lower) substituent at the first point of difference.

Checkpoint — Constitutional Isomers & Cyclo Naming

Exit Ticket — Putting a Name Together

Conformational Analysis

Part 2 of 7 — Conformations of Acyclic Alkanes

A single $C-C$ $\sigma$ bond allows nearly free rotation. The different three-dimensional arrangements a molecule adopts by rotating about such bonds are called conformations (or conformers / rotamers). Crucially, interconverting conformers does not break any bonds — they are not isomers in the structural sense and cannot be separated under normal conditions.

The key variable is the dihedral (torsion) angle: the angle between two bonds on adjacent carbons when viewed down the $C-C$ axis. As this angle changes, the molecule's potential energy rises and falls. The energy differences are small (a few $\text{kcal/mol}$), so at room temperature molecules spin rapidly through all conformations — but they spend most of their time in the lowest-energy ones.

Two kinds of strain raise the energy of a conformation:

• Torsional strain — resistance to eclipsing of bonds on adjacent atoms (electron–electron / bond–bond repulsion when bonds align). Maximal in eclipsed conformers.
• Steric strain — repulsion between atoms or groups forced too close in space (van der Waals overlap). Important when bulky groups approach each other.

The tool we use to visualize all this is the Newman projection.

Newman Projections: Staggered vs. Eclipsed

A Newman projection views the molecule straight down one $C-C$ bond. The front carbon is drawn as a dot with three bonds radiating from the center; the back carbon is a circle with three bonds emerging from its edge.

Ethane ($CH_3-CH_3$) is the simplest case, with two limiting conformations:

In the staggered conformation the front C–H bonds bisect the gaps between the back C–H bonds, minimizing bond–bond repulsion. In the eclipsed conformation the front and back bonds line up directly, maximizing torsional strain.

The barrier to rotation in ethane is about $3\;\text{kcal/mol}$, which corresponds to roughly $1\;\text{kcal/mol}$ per eclipsed $H{-}H$ interaction (three pairs eclipse simultaneously). This barrier is low enough that ethane rotates millions of times per second at room temperature, yet high enough that the staggered form is meaningfully favored.

Key point: Eclipsed conformers are transition states / energy maxima on the rotational pathway, not stable wells. Staggered conformers are the minima. The molecule continuously passes through eclipsed arrangements but dwells in staggered ones.

Checkpoint — Ethane Conformers

Butane: Gauche and Anti

When the rotating bond carries larger groups, steric strain joins torsional strain. The classic case is rotation about the C2–C3 bond of butane ($CH_3-CH_2-CH_2-CH_3$), where each of the two central carbons bears a methyl group.

As we rotate the back carbon through $360^\circ$, four named conformers recur:

Anti ($180^\circ$) is the most stable: the two methyl groups are as far apart as possible (staggered and anti), so neither steric nor torsional strain is significant.

Gauche ($60^\circ$) is still staggered (no torsional strain) but the two methyls are only $60^\circ$ apart, close enough to feel mutual van der Waals repulsion. This gauche interaction costs about $0.9\;\text{kcal/mol}$ — a number worth memorizing, because it reappears in cyclohexane analysis.

Totally eclipsed ($0^\circ$) is the global maximum: the two methyls eclipse each other directly, combining the worst torsional and steric penalties.

Worked ranking: From most to least stable, butane's conformers about C2–C3 are anti ($180^\circ$) > gauche ($60^\circ$) > methyl/H eclipsed ($120^\circ$) > methyl/methyl eclipsed ($0^\circ$). Notice both staggered forms (anti, gauche) sit below both eclipsed forms — torsional strain dominates the energy landscape, with steric strain breaking ties among similar conformers.

Checkpoint — Butane Conformers

Exit Ticket — Strain Concepts

Cyclohexane Chair Conformations

Part 3 of 7 — The Chair, Ring Flips, and Diaxial Strain

Although a flat hexagon suggests $120^\circ$ internal angles, a real cyclohexane ring is not planar. By puckering, the ring lets every carbon adopt nearly perfect tetrahedral ($109.5^\circ$) angles and keeps every adjacent $C-H$ pair staggered. The result is the famous chair conformation — essentially strain-free (no angle strain, no torsional strain).

Cyclohexane also visits higher-energy shapes — the half-chair (the rotational transition state, $\approx 10\;\text{kcal/mol}$), the twist-boat ($\approx 5\text{–}6\;\text{kcal/mol}$), and the boat ($\approx 7\;\text{kcal/mol}$, suffering "flagpole" steric strain and eclipsing) — but at any instant the overwhelming majority of molecules are chairs.

In the chair, the twelve hydrogens divide into two geometrically distinct sets:

• Axial hydrogens point straight up or down, parallel to the ring's vertical axis (alternating up/down around the ring).
• Equatorial hydrogens point outward, roughly along the "equator" of the ring.

Each carbon bears exactly one axial and one equatorial hydrogen.

The Ring Flip

A chair can interconvert with its mirror-image chair by a process called a ring flip, passing through the half-chair, twist-boat, and boat forms. The barrier ($\approx 10\;\text{kcal/mol}$) is low enough that ring-flipping occurs thousands of times per second at room temperature.

The defining consequence of a ring flip:

Every axial position becomes equatorial, and every equatorial position becomes axial. A group's "up vs. down" orientation in space is preserved, but its "axial vs. equatorial" label is swapped.

For unsubstituted cyclohexane the two chairs are identical in energy, so this is invisible. But the moment we add a substituent, the two chairs differ — and the molecule spends more time in whichever chair places the substituent equatorial.

Why equatorial is preferred — 1,3-diaxial strain: An axial substituent on, say, C1 points straight up, directly toward the other two axial groups on the same face (the axial hydrogens on C3 and C5). These 1,3-diaxial interactions are sterically unfavorable — they are, in fact, exactly the gauche relationships from butane, transplanted onto the ring. Placing the substituent equatorial relieves this crowding.

Checkpoint — Chair Geometry & Ring Flip

A-Values: Quantifying the Equatorial Preference

The energetic penalty for placing a given group axial instead of equatorial is tabulated as its A-value (in $\text{kcal/mol}$). A larger A-value means a stronger equatorial preference.

The A-value equals the free-energy difference between the axial and equatorial chairs, so it determines the equilibrium ratio. For methylcyclohexane (A $\approx 1.7$), the equilibrium is roughly 95% equatorial : 5% axial at room temperature.

The standout is tert-butyl. Its A-value ($\approx 4.7\;\text{kcal/mol}$) is so large that the axial chair is essentially never populated — tert-butylcyclohexane is "locked" with the tert-butyl group equatorial. This is why tert-butyl is used as a conformational anchor in problems: it fixes which chair you are looking at.

A-value bookkeeping: A methyl axial costs $\approx 1.7\;\text{kcal/mol}$, which is about two gauche-butane interactions ($2 \times 0.9 \approx 1.8$). This is not a coincidence — each 1,3-diaxial interaction between the axial methyl and an axial ring hydrogen is a gauche relationship.

Cis/Trans Isomerism in Rings & Disubstituted Chairs

Because a ring has two distinct faces, two substituents on different ring carbons can be on the same face (cis) or opposite faces (trans). Unlike conformers, cis and trans ring isomers are genuine stereoisomers — they cannot interconvert by ring-flipping or bond rotation; converting one to the other would require breaking bonds. (A ring flip changes axial/equatorial labels but never changes cis/trans relationships.)

To find the most stable chair of a disubstituted cyclohexane, place the larger group equatorial and check whether the geometry forces the other group axial.

Worked example — 1,4-dimethylcyclohexane:

• trans-1,4: the two methyls can be both equatorial (e,e) in one chair — this is the lowest-energy arrangement, with no axial methyls. (The other chair is diaxial and much higher.) Therefore trans-1,4-dimethylcyclohexane is more stable than cis.
• cis-1,4: one methyl is forced axial and the other equatorial (a,e) in both chairs — ring-flipping just swaps which methyl is axial, so there is always exactly one axial methyl ($\approx 1.7\;\text{kcal/mol}$ penalty). It can never reach an all-equatorial state.

The 1,4 pattern in general: trans-1,4 gives e,e (most stable); cis-1,4 is forced a,e. For 1,2 and 1,3 the pattern inverts — e.g. for 1,3-dimethyl, the cis isomer can be e,e while trans is forced a,e.

Trap: Students often assume "trans is always more stable." That is true for 1,4 (and 1,2), but for 1,3-disubstituted cyclohexanes the cis isomer is the one that can be diequatorial — so cis-1,3 beats trans-1,3. Always work out the actual axial/equatorial outcome rather than memorizing "trans wins."

Checkpoint — A-Values, Cis/Trans, Stability

Exit Ticket — Picking the Better Chair

Ring Strain in Cycloalkanes

Part 4 of 7 — Angle Strain, Torsional Strain, and Ring Size

The internal angles of a flat regular polygon are fixed by geometry: $60^\circ$ for a triangle (cyclopropane), $90^\circ$ for a square (cyclobutane), $108^\circ$ for a pentagon (cyclopentane), $120^\circ$ for a hexagon. But $sp^3$ carbon "wants" the tetrahedral angle of $109.5^\circ$. The deviation from this ideal produces angle (Baeyer) strain.

Total ring strain has up to three contributors:

• Angle strain — compression or expansion of $C-C-C$ bond angles away from $109.5^\circ$.
• Torsional strain — eclipsing of $C-H$ bonds on adjacent ring carbons (worst when the ring is forced flat).
• Steric (transannular) strain — atoms across the ring crowding each other (matters in medium rings).

Real rings are not flat (except cyclopropane, which has no choice): cyclobutane, cyclopentane, and cyclohexane pucker to trade a little angle strain for a large relief of torsional strain. The chair of cyclohexane is the perfect solution — zero angle strain and zero torsional strain.

Measuring Strain: Heat of Combustion per $CH_{2}$

We cannot weigh "strain" directly, but we can measure it through heats of combustion ($\Delta H_{comb}$). Every cycloalkane $C_nH_{2n}$ burns to $CO_2$ and $H_2O$; a strained ring stores extra energy that is released on combustion, so it burns "hotter" per $CH_2$ unit than a strain-free reference.

Dividing $\Delta H_{comb}$ by the number of $CH_2$ groups gives a per-$CH_2$ value. A completely strain-free $CH_2$ releases about $157.4\;\text{kcal/mol}$ (the value for a long unstrained chain). Excess above this reflects ring strain:

Worked interpretation: Cyclopropane's $\Delta H_{comb}$ per $CH_2$ is the largest of any cycloalkane, signaling the most strain per carbon. The total strain happens to be similar for cyclopropane and cyclobutane (~26–27 kcal/mol), but spread over fewer carbons in cyclopropane, so per $CH_2$ cyclopropane is clearly the most strained. Cyclohexane sits at the baseline — essentially zero strain.

Checkpoint — Sources of Strain

Ring-by-Ring Tour

Cyclopropane (most strained). Three carbons necessarily lie in one plane, fixing the angle at $60^\circ$ — a catastrophic deviation from $109.5^\circ$. The bonds cannot point directly at one another, so they form "banana bonds" (bent bonds) with poor orbital overlap. All C–H bonds are eclipsed. Result: high reactivity for a "saturated" compound; cyclopropane rings open under conditions alkanes ignore.

Cyclobutane (still significantly strained). Internal angle near $90^\circ$. The ring puckers to relieve some torsional strain, leaving substantial angle strain. Total strain is comparable to cyclopropane, but distributed over four carbons.

Cyclopentane (nearly strain-free). A flat pentagon would have $108^\circ$ angles — almost perfect! — but a flat ring eclipses all C–H bonds. To dodge that torsional strain it adopts a puckered "envelope" (one carbon out of plane) or twist conformation. Net strain is small (~6.5 kcal/mol total, ~1.3 per $CH_{2}$).

Cyclohexane (strain-free). The chair achieves $\approx 109.5^\circ$ angles and fully staggered bonds simultaneously. Essentially zero strain. This is why six-membered rings are by far the most common ring size in nature and pharmaceuticals.

Medium rings (7–12). These reintroduce modest strain from transannular (across-the-ring) steric crowding and imperfect staggering, before large rings (>14) become essentially strain-free again.

Reactivity link: Strain is stored energy, so the most strained rings (cyclopropane, cyclobutane) are the most reactive toward ring-opening reactions. Cyclohexane, being strain-free, behaves like an ordinary unreactive alkane. "Saturated" does not always mean "unreactive" when strain is present.

Checkpoint — Ring Size and Stability

Exit Ticket — Comparing Strain

Physical Properties of Alkanes

Part 5 of 7 — Boiling Points, Branching, Solubility, and IMFs

Alkanes are nonpolar. The $C-C$ bond is between identical atoms, and the $C-H$ bond has only a tiny electronegativity difference ($\approx 0.4$), so alkanes have essentially no permanent dipole. With no dipole–dipole forces and no hydrogen bonding, the only intermolecular force available to alkanes is the weakest one: London dispersion forces (LDF), also called van der Waals or induced-dipole/induced-dipole forces.

London forces arise from instantaneous, fluctuating electron distributions that induce complementary dipoles in neighbors. Two features make them stronger:

1. More electrons / larger surface area — bigger molecules are more polarizable, so LDF grows with molecular size.
2. More contact area — molecules that can pack closely along their surfaces attract more strongly than compact, ball-like molecules.

Because all of an alkane's physical behavior traces back to weak, size-dependent London forces, the property trends are remarkably systematic and predictable.

Boiling Point: Size and Branching

Trend 1 — Boiling point rises with molecular weight (chain length). Each added $CH_2$ adds electrons and surface area, strengthening London forces, so more thermal energy is needed to separate molecules into the gas phase. The first four straight-chain alkanes ($CH_4$, ethane, propane, butane) are gases at room temperature; $C_5$–$C_{17}$ are liquids; $C_{18}$ and up are waxy solids. Roughly, each additional carbon raises the boiling point by $20\text{–}30\,^{\circ}\text{C}$ in the lower members.

Trend 2 — Branching lowers boiling point. For a fixed formula, a branched isomer boils lower than its straight-chain relative. Branching makes the molecule more compact and spherical, reducing the surface area available for intermolecular contact, which weakens London forces.

Worked comparison — the three $C_5H_{12}$ isomers:

• n-pentane (straight): bp $\approx 36\,^{\circ}\text{C}$ — most extended, most surface contact, highest bp.
• isopentane (2-methylbutane, one branch): bp $\approx 28\,^{\circ}\text{C}$.
• neopentane (2,2-dimethylpropane, most branched, nearly spherical): bp $\approx 9.5\,^{\circ}\text{C}$ — least surface contact, lowest bp.

All three have the same molecular weight (72 g/mol), so the ~26 °C spread is due entirely to shape: the more spherical the molecule, the weaker its London forces and the lower its boiling point.

Checkpoint — Boiling Point Trends

Solubility, Density, and Melting Point

Solubility — "like dissolves like." Alkanes are nonpolar and cannot hydrogen-bond, so they are insoluble in water (a polar, hydrogen-bonding solvent). Dissolving an alkane in water would require breaking favorable water–water hydrogen bonds without compensating attraction — entropically and enthalpically unfavorable (the hydrophobic effect). Alkanes are, however, fully miscible with other nonpolar solvents (other hydrocarbons, ethers, $CCl_4$). They are themselves excellent nonpolar solvents — this is the chemistry behind oils and greases.

Density. All liquid alkanes are less dense than water ($\rho \approx 0.62\text{–}0.80\;\text{g/mL}$). Combined with their insolubility, this is why oil floats on water and forms a distinct upper layer — the basis of oil spills and of separations in a separatory funnel.

Melting point. Melting points also generally rise with molecular weight, but less smoothly than boiling points, because melting depends on how well molecules pack into a crystal lattice, not just on IMF strength. Alkanes with an even number of carbons often pack more efficiently than odd-numbered neighbors, producing a characteristic "sawtooth" alternation in the melting-point trend. Symmetry matters too: highly symmetric neopentane has an anomalously high melting point for its size because its spherical shape packs neatly into a crystal.

Key contrast: Boiling point is governed almost entirely by IMF strength (size + surface area), so its trend is smooth. Melting point adds a packing/symmetry dependence, so its trend is bumpier — a symmetric molecule can boil low (weak LDF) yet melt high (efficient packing), as neopentane does.

Checkpoint — Solubility & Packing

Exit Ticket — Predicting Properties

Reactions Workshop: Combustion & Radical Halogenation

Part 6 of 7 — Putting Reactivity to Work

Alkanes are famously unreactive ("paraffins" — parum affinis, "little affinity"): no $\pi$ bonds, no lone pairs, no polar bonds for nucleophiles or electrophiles to attack. Only two reactions matter at this level, and both involve breaking strong C–H bonds homolytically:

1. Combustion — complete oxidation with $O_2$ to $CO_2$ and $H_2O$, releasing large amounts of heat. The basis of alkanes as fuels.
2. Free-radical halogenation — substitution of an H by a halogen (Cl or Br) under heat or light, proceeding through a radical chain mechanism.

This workshop drills the mechanism and the selectivity arithmetic you will be tested on.

Combustion and Balancing

Complete combustion of any alkane follows:

$C_nH_{2n+2} + \left(\dfrac{3n+1}{2}\right) O_2 \rightarrow n\,CO_2 + (n+1)\,H_2O$

The reaction is highly exothermic (large negative $\Delta H$): for example, octane releases roughly $-1300\;\text{kcal/mol}$. Insufficient oxygen gives incomplete combustion, producing $CO$ (toxic) and soot ($C$) instead of $CO_2$.

Worked balance — propane ($C_3H_8$): Here $n = 3$, so the coefficient of $O_2$ is $(3\cdot 3 + 1)/2 = 5$. Products: $3\,CO_2$ and $4\,H_2O$, giving $C_3H_8 + 5\,O_2 \rightarrow 3\,CO_2 + 4\,H_2O$ with $\Delta H \approx -530\;\text{kcal/mol}$.

Combustion is not synthetically useful (it destroys the carbon skeleton), but its heats of combustion are exactly the data we used in Part 4 to quantify ring strain.

The Radical Chain Mechanism

Free-radical halogenation (e.g. $CH_4 + Cl_2 \xrightarrow{h\nu} CH_3Cl + HCl$) proceeds in three stages. Memorize the stage names and what each does:

1. Initiation. Light or heat homolytically cleaves the weak halogen–halogen bond, generating two halogen radicals. This is the only step that creates radicals from a closed-shell molecule:

$Cl_2 \xrightarrow{h\nu} 2\,\text{Cl}\cdot$

2. Propagation (two steps that repeat). Radicals are consumed and regenerated, so a single initiation event can drive thousands of cycles:

• $\text{Cl}\cdot + CH_4 \rightarrow HCl + \text{CH}_3\cdot$ (abstracts H, forms a carbon radical)
• $\text{CH}_3\cdot + Cl_2 \rightarrow CH_3Cl + \text{Cl}\cdot$ (forms product, regenerates $\text{Cl}\cdot$)

Note the chlorine radical is regenerated, so it carries the chain forward — this is what makes it a chain reaction.

3. Termination. Any two radicals combine, removing radicals from the system and ending a chain:

$\text{Cl}\cdot + \text{Cl}\cdot \rightarrow Cl_2 \qquad \text{CH}_3\cdot + \text{Cl}\cdot \rightarrow CH_3Cl \qquad \text{CH}_3\cdot + \text{CH}_3\cdot \rightarrow C_2H_6$

Diagnostic: The hallmark of initiation is "2 radicals made from 0." The hallmark of propagation is "radical in, radical out" (count stays the same). The hallmark of termination is "2 radicals in, 0 out." The trace of ethane ($C_2H_6$) found in chlorination of methane is direct evidence of the $\text{CH}_3\cdot + \text{CH}_3\cdot$ termination step.

Checkpoint — Mechanism Steps

Selectivity: Which Hydrogen Reacts?

When an alkane has different types of hydrogens (primary 1°, secondary 2°, tertiary 3°), halogenation can occur at each. Two factors set the product ratio:

• Number of hydrogens of each type (statistical factor).
• Relative reactivity per hydrogen, because radical stability is $3° > 2° > 1°$ (more substituted radicals are more stable). Abstracting an H to make a more stable radical is easier.

The two halogens differ dramatically:

This trade-off is Hammond's postulate in action: bromination has a late, product-like transition state, so it "feels" the radical-stability difference strongly and is selective; chlorination's early transition state barely distinguishes the H types.

Worked example — monochlorination of propane ($CH_3CH_2CH_3$): Propane has six 1° H (two $CH_3$ groups) and two 2° H (the central $CH_2$).

• Yield at 1° positions $\propto 6 \times 1 = 6$
• Yield at 2° position $\propto 2 \times 3.8 = 7.6$ Ratio (2° : 1°) $= 7.6 : 6$, i.e. about 56% 2-chloropropane and 44% 1-chloropropane. Even though 2° H are outnumbered 3-to-1, the higher per-H reactivity makes 2-chloropropane the slight major product.

Contrast with bromination of propane: using $1640 : 82 : 1 \rightarrow$ relative 2° yield $= 2 \times 82 = 164$ vs 1° yield $= 6 \times 1 = 6$, giving $\approx 96\%$ 2-bromopropane. The selectivity of bromine overwhelms the statistical factor.

Checkpoint — Selectivity Arithmetic

Exit Ticket — Choosing a Reagent

Synthesis & Cumulative Review

Part 7 of 7 — Tying It All Together

You now hold the full toolkit for alkanes and cycloalkanes. This capstone integrates the four pillars of the unit:

1. Nomenclature — encoding exact connectivity into an unambiguous IUPAC name (longest chain, lowest locants, alphabetized substituents). The grammar that lets chemists communicate structure.
2. Conformational analysis — predicting the populated 3-D shapes from torsional and steric strain (Newman projections; anti/gauche; chair, axial/equatorial, ring flip, A-values).
3. Ring strain — explaining stability and reactivity of cycloalkanes from angle and torsional strain, quantified by heats of combustion.
4. Reactions — combustion (the fuel reaction) and free-radical halogenation (initiation/propagation/termination, with selectivity).

The connecting thread across all four is stability driven by minimizing strain and maximizing favorable orbital/steric arrangements — the same logic whether you are choosing a chair, ranking a conformer, or predicting a halogenation product.

How the Pieces Connect

A single worked thread shows the concepts reinforcing one another.

Consider trans-1,2-dimethylcyclohexane undergoing monobromination at a tertiary-like position:

• Nomenclature tells us the substrate: a cyclohexane ring (parent, 6 C beats any chain) with two methyls, trans, at C1 and C2.
• Conformational analysis tells us its shape: the trans-1,2 isomer can place both methyls equatorial (e,e) — the favored chair — minimizing 1,3-diaxial strain.
• Ring strain reminds us the six-membered ring itself is strain-free, so the ring will not open; only substitution is possible.
• Reactions let us predict that $Br_2/h\nu$ — the selective halogen — will abstract the most stable (most substituted) hydrogen, e.g. a $3°$ ring C–H, giving predominantly one product.

Numbers worth carrying forward (units: $\text{kcal/mol}$):

Unifying idea: "Most stable = least strained." A methyl prefers equatorial for the same reason butane prefers anti — to escape gauche/1,3-diaxial crowding. A bromine radical prefers the 3° site because the resulting radical is most stable. Strain accounting governs both conformational preference and chemical selectivity.

Cumulative Checkpoint — Nomenclature & Isomerism

Cumulative Checkpoint — Conformation & Strain

Final Synthesis — Reactions in Context