Alkane reactions: combustion, halogenation, and radical selectivity

Intuition Beginner

Alkanes are the chemically quietest class of organic molecules. Every carbon is fully saturated with C–H and C–C single bonds, so there are no lone pairs, no pi bonds, and no obvious sites for a reagent to attack. In practice, alkanes undergo only two broad classes of reaction: combustion (burning in oxygen) and halogenation (substitution of a hydrogen by a halogen such as Cl or Br).

Combustion is the reaction you already know from everyday life. Methane (${\text{CH}}_4$) burns in oxygen to give carbon dioxide and water. When oxygen is limited, combustion is incomplete and produces carbon monoxide (CO) or even elemental carbon (soot) instead of ${\text{CO}}_2$. Complete combustion releases more energy per mole of fuel than incomplete combustion. The general equations are:

$${\text{C}}_n{\text{H}}_{2n+2}+\frac{3n+1}{2}{\text{O}}_2⟶n{\text{CO}}_2+(n+1){\text{H}}_2\text{O}$$

Halogenation is more interesting chemically. When a mixture of an alkane and chlorine gas (${\text{Cl}}_2$) is heated or exposed to UV light, a chlorine atom replaces one hydrogen to form an alkyl chloride ($\text{R-Cl}$) plus HCl. The same reaction works with bromine (${\text{Br}}_2$), but bromine is slower and more selective about which hydrogen it replaces.

Selectivity means: if the alkane has more than one type of C–H bond (for example, 1°, 2°, and 3° hydrogens in isobutane), the halogen does not replace them at equal rates. Bromine strongly prefers the weakest C–H bond (the 3° hydrogen). Chlorine is less discriminating — it reacts faster overall but gives a product mixture closer to the statistical ratio of hydrogen types.

Visual Beginner

Reaction	General equation	Conditions	Key feature
Complete combustion	$\text{R-H}+{\text{O}}_2\to {\text{CO}}_2+{\text{H}}_2\text{O}$	Spark or flame	Maximum energy release
Incomplete combustion	$\text{R-H}+{\text{O}}_2\to \text{CO}+{\text{H}}_2\text{O}$	Limited ${\text{O}}_2$	Produces toxic CO
Chlorination	$\text{R-H}+{\text{Cl}}_2\to \text{R-Cl}+\text{HCl}$	Heat or $h\nu$	Fast, low selectivity
Bromination	$\text{R-H}+{\text{Br}}_2\to \text{R-Br}+\text{HBr}$	Heat or $h\nu$	Slow, high selectivity

Relative reactivity per hydrogen type (25 °C):

Hydrogen type	With ${\text{Cl}}_2$	With ${\text{Br}}_2$
1° (primary)	1	1
2° (secondary)	3.8	82
3° (tertiary)	5	1600

Bromine's preference for 3° over 1° positions is roughly 1600:1, while chlorine's is only about 5:1. This dramatic difference has a thermodynamic explanation covered in the intermediate section.

Worked example Beginner

Predict the major monochlorination product(s) of isobutane (2-methylpropane).

Isobutane has the structure (${\text{CH}}_3$)${}_3$CH. There are two types of hydrogen:

• 1° hydrogens: 9 equivalent hydrogens on the three methyl groups.
• 3° hydrogen: 1 hydrogen on the central carbon.

Step 1. Calculate the statistical ratio. If all hydrogens were equally reactive, the 1°:3° product ratio would be 9:1 (because there are 9 one-degree hydrogens and 1 three-degree hydrogen).

Step 2. Apply the relative reactivity per hydrogen. For chlorination at 25 °C, a 3° hydrogen is about 5 times as reactive as a 1° hydrogen. So the reactivity-weighted ratio is:

$$\text{1° product}:\text{3° product}=(9\times 1):(1\times 5)=9:5$$

Step 3. Convert to percentages.

• 1° product (1-chloro-2-methylpropane): $\frac{9}{14}\approx 64\%$
• 3° product (2-chloro-2-methylpropane): $\frac{5}{14}\approx 36\%$

The major product is 1-chloro-2-methylpropane, but a substantial amount of the tertiary chloride also forms. Chlorination is not selective enough to give a single product.

If the same reaction were done with bromine instead of chlorine, the 3°:1° selectivity ratio of ~1600:1 would make 2-bromo-2-methylpropane essentially the only product.

Check your understanding Beginner

Formal definition Intermediate+

A free radical is a species with an unpaired electron. In organic chemistry, the most relevant radicals are carbon-centred (${\text{R}}_3\text{C}\cdot$) and halogen-centred ($\text{X}\cdot$). Radical reactions proceed through a chain mechanism consisting of three phases.

Initiation. A halogen molecule undergoes homolytic cleavage to generate two halogen radicals:

$${\text{Cl}}_2\stackrel{h\nu \text{ or }\Delta }{\to }2\text{Cl}\cdot \Delta H=+242{\text{ kJ mol}}^{-1}$$

The Cl–Cl bond dissociation energy (BDE) is 242 kJ mol${}^{-1}$. UV light or thermal energy supplies this energy.

Propagation. Two steps that each consume one radical and produce one new radical, sustaining the chain:

1. $\text{Cl}\cdot +\text{R-H}\to \text{H-Cl}+\text{R}\cdot$ (hydrogen abstraction)
2. $\text{R}\cdot +{\text{Cl}}_2\to \text{R-Cl}+\text{Cl}\cdot$ (halogen abstraction)

The net reaction is $\text{R-H}+{\text{Cl}}_2\to \text{R-Cl}+\text{HCl}$. A single initiation event can drive thousands of propagation cycles.

Termination. Any combination of two radicals removes radicals from the cycle:

$$\text{Cl}\cdot +\text{Cl}\cdot \to {\text{Cl}}_2\text{R}\cdot +\text{Cl}\cdot \to \text{R-Cl}\text{R}\cdot +\text{R}\cdot \to \text{R-R}$$

Termination is inevitable but slow, because radical concentrations remain very low ($\sim 10^{-8}$ M) during the chain.

Bond dissociation energies (BDEs). The C–H BDE determines how easily a hydrogen is abstracted. Typical values:

C–H bond type	BDE (kJ mol${}^{-1}$)
1° C–H (e.g., ethane)	421
2° C–H (e.g., propane)	401
3° C–H (e.g., isobutane)	385
Allylic C–H (e.g., propene)	364
Benzylic C–H (e.g., toluene)	375

Weaker C–H bonds produce more stable carbon radicals, because the radical centre is stabilised by hyperconjugation (3° > 2° > 1°) and, for allylic/benzylic positions, by resonance delocalisation.

Thermodynamic driving force for propagation. The enthalpy of the H-abstraction step ($\text{Cl}\cdot +\text{R-H}\to \text{HCl}+\text{R}\cdot$) depends on the difference between the H–Cl BDE (431 kJ mol${}^{-1}$) and the C–H BDE being broken:

$$\Delta H_{\text{abstraction}}=\text{BDE}(\text{C-H})-\text{BDE}(\text{H-Cl})$$

For abstraction of a 1° hydrogen: $\Delta H\approx 421-431=-10$ kJ mol${}^{-1}$ (exothermic). For abstraction of a 3° hydrogen: $\Delta H\approx 385-431=-46$ kJ mol${}^{-1}$ (more exothermic).

The corresponding calculation for bromine uses the H–Br BDE of 366 kJ mol${}^{-1}$:

For abstraction of a 1° hydrogen: $\Delta H\approx 421-366=+55$ kJ mol${}^{-1}$ (endothermic). For abstraction of a 3° hydrogen: $\Delta H\approx 385-366=+19$ kJ mol${}^{-1}$ (endothermic, but less so).

The endothermicity of H-abstraction by bromine is the key to its high selectivity, as explained below.

Key mechanism Intermediate+

The Hammond postulate and halogen selectivity. The Hammond postulate states that the transition state of a reaction resembles the species (reactant or product) to which it is closer in energy. For an exothermic step, the transition state resembles the reactants (early TS). For an endothermic step, the transition state resembles the products (late TS).

Chlorine H-abstraction is exothermic for all C–H types. The transition state is reactant-like, meaning the C–H bond is only slightly broken and the radical character on carbon is minimal. The TS energy depends mostly on the reactant energy, so the difference between abstracting a 1°, 2°, or 3° hydrogen is small. Chlorination shows low selectivity.

Bromine H-abstraction is endothermic for all C–H types. The transition state is product-like, meaning the C–H bond is largely broken and significant radical character has developed on carbon. The TS energy reflects the stability of the incipient carbon radical. Since 3° radicals are much more stable than 1° radicals, the 3° TS is much lower in energy. Bromination shows high selectivity.

This is the central mechanistic insight: the more endothermic the H-abstraction step, the more selective the halogen. Fluorination (strongly exothermic, very early TS) is virtually unselective and dangerously vigorous. Iodination (endothermic for all H types, very late TS) is so selective it essentially does not proceed thermodynamically.

Statistical vs reactivity control. The product distribution in a halogenation reaction is determined by both the number of each type of hydrogen and the relative rate of abstraction per hydrogen:

$$\text{yield of product }i=\frac{n_i\times r_i}{{\sum }_j{n}_j\times r_j}$$

where $n_i$ is the number of hydrogens of type $i$ and $r_i$ is the relative reactivity per hydrogen of that type. For chlorination, $r_i$ values are small (1, 3.8, 5) and the statistical factor ($n_i$) dominates. For bromination, $r_i$ values are enormous (1, 82, 1600) and the reactivity factor dominates, overwhelming any statistical advantage of more abundant 1° hydrogens.

Connections Master

• Functional groups and nomenclature 15.02.01. Halogenation of alkanes produces alkyl halides, which are classified as a functional group in the naming hierarchy. Identifying whether the product is a primary, secondary, or tertiary halide requires the nomenclature tools from the preceding unit.

• Stereochemistry 15.01.01. Radical halogenation at a stereogenic centre proceeds with racemisation, because the planar sp${}^2$ carbon radical intermediate is achiral and can be attacked from either face in the second propagation step. This contrasts with ${\text{S}}_{\text{N}}2$ substitution, which inverts configuration.

• Thermodynamics and kinetics 14.04.01. The selectivity difference between Cl and Br is a direct application of the Hammond postulate, which connects thermodynamic driving force ($\Delta H$) to transition-state structure and hence to selectivity. This is one of the most cited examples of kinetic vs thermodynamic control in organic chemistry.

• Radical and pericyclic chemistry 15.08.01. The radical chain mechanism introduced here (initiation, propagation, termination) is the template for all radical reactions. Allylic and benzylic halogenation, radical additions to alkenes, and radical cyclisations all use the same mechanistic framework. The broader treatment of radical chemistry in 15.08.01 extends the selectivity analysis to multi-step radical cascades.

• Carbonyl chemistry 15.07.01. Alkyl halides produced by radical halogenation are precursors to Grignard reagents and organolithium compounds, which are the nucleophilic partners in carbonyl addition reactions. The functional-group interconnection chain is: alkane → radical halogenation → alkyl halide → Grignard → alcohol (via carbonyl addition).

• Spectroscopy 15.11.01. The C–Cl stretch appears in IR near 600–800 cm${}^{-1}$, and the chlorine-bearing carbon is shifted downfield to $\delta$ 40–50 in ${}^{13}$C NMR. These diagnostic signals allow verification of the halogenation product and assessment of regiochemistry.

• Free-radical polymerisation 15.08.02 pending. The chain mechanism (initiation, propagation, termination) in halogenation is mechanically identical to free-radical addition polymerisation, where a radical adds across a C=C bond and the chain grows by repeated propagation steps.

Historical notes Master

The radical theory of halogenation emerged from the work of several chemists across the early twentieth century. Moses Gomberg at the University of Michigan discovered the triphenylmethyl radical in 1900, providing the first evidence that carbon-centred free radicals could exist as persistent species. Gomberg's discovery established that the trivalent carbon radical is a real chemical entity, not merely a formal construct.

The chain mechanism for halogenation was proposed independently by Max Bodenstein (1913) and developed rigorously by Karl Herzfeld and Hugh Taylor in the 1920s. Bodenstein studied the photochemical chlorination of hydrogen (${\text{H}}_2+{\text{Cl}}_2\to 2\text{HCl}$) and recognised that a single photon could initiate a chain reaction converting many molecules — the quantum yield far exceeded 1. The application of Bodenstein's chain concept to organic halogenation followed in the 1930s.

George Hammond formulated the postulate bearing his name in 1955 (Journal of the American Chemical Society 77, 334–338): "If two states, as for example, a transition state and an unstable intermediate, occur consecutively during a reaction process and have nearly the same energy content, their interconversion will involve only a small reorganisation of the molecular structures." The Hammond postulate provided the theoretical framework for understanding why bromine is more selective than chlorine — a fact that had been known empirically since the nineteenth century but lacked a mechanistic explanation.

Derek Barton (Nobel Prize, 1969) exploited the selectivity of radical reactions in developing the Barton reaction (nitrite photolysis, 1960), which uses an intramolecular radical hydrogen transfer to functionalise unactivated C–H bonds at specific positions. The Barton reaction demonstrated that radical selectivity could be harnessed for synthetic purposes, countering the prevailing view that radical reactions were too uncontrolled for use in synthesis.

Quantitative selectivity ratios were measured systematically by G. A. Russell and coworkers in the 1950s and 1960s. The values for chlorine (1 : 3.8 : 5 at 25 °C) and bromine (1 : 82 : 1600 at 25 °C) have been confirmed across multiple substrates and remain the standard reference data.

The development of C–H functionalisation as a field in the twenty-first century (Du Bois, White, Baran, Hartwig, and others) has its roots in the observation that radical halogenation is the simplest and oldest example of selective C–H bond activation. Modern C–H functionalisation methods aim to achieve for C–O, C–N, and C–C bonds what radical halogenation achieves for C–X bonds: selective transformation of a specific C–H bond in the presence of many others.

Bibliography Master

Foundational radical chemistry.

• Gomberg, M., "An Instance of Trivalent Carbon: Triphenylmethyl", Journal of the American Chemical Society 22 (1900), 757–771.
• Bodenstein, M., "Eine Theorie der photochemischen Reaktionsgeschwindigkeiten", Zeitschrift für Physikalische Chemie 85 (1913), 329–397.
• Hammond, G. S., "A Correlation of Reaction Rates", Journal of the American Chemical Society 77 (1955), 334–338.

Textbook references.

• McMurry, J., Organic Chemistry, 10th ed. (Cengage, 2019), Ch. 4.
• Clayden, J., Greeves, N. & Warren, S., Organic Chemistry, 2nd ed. (Oxford UP, 2012), Ch. 19–20.
• Smith, M. B., March's Advanced Organic Chemistry, 7th ed. (Wiley, 2013), Ch. 14.

Advanced treatments.

• Russell, G. A., "Reactivity, Selectivity, and the Hammett Equation for Free-Radical Reactions", in Free Radicals in Organic Chemistry (Wiley, 1973).
• Barton, D. H. R., "The Invention of Radical Reactions. Part XVI. Some Observations on the Decarboxylative Rearrangement of Hypochlorites", Tetrahedron 48 (1992), 2529–2542.
• Walling, C., Free Radicals in Solution (Wiley, 1957).

C–H functionalisation.

• White, M. C., "Alkane C–H Activation by Single-Electron Transfer: A Radical Approach", Science 335 (2012), 807–809.
• Hartwig, J. F., "Catalytic C–H Functionalization: Beyond the Orthodox Paradigm", Nature 455 (2008), 314–322.
• Newhouse, T. & Baran, P. S., "Skeletal Editing: Direct Insertions, Deletions, and Replacements in Organic Synthesis", Angewandte Chemie International Edition 50 (2011), 9866–9878.

Physical organic chemistry.

• Anslyn, E. V. & Dougherty, D. A., Modern Physical Organic Chemistry (University Science Books, 2006), Ch. 11 (radical chemistry).
• Lowry, T. H. & Richardson, K. S., Mechanism and Theory in Organic Chemistry, 3rd ed. (Harper & Row, 1987), Ch. 12.