Free Radicals Organic Chemistry

How Radical Reactions Actually Work in the Lab (When They Go Right)

The most common mistake students make when studying free radicals is thinking of them as textbook intermediates. In practice, they are the reason half your reactions in the fume hood turn brown, smell terrible, and give you three products instead of one. Understanding the mechanism is one thing. Managing the reaction so it actually does what you want is another.

Free radical chemistry involves species with unpaired electrons that propagate chain reactions through initiation, propagation, and termination steps. The simplest example is halogenation of an alkane with bromine under UV light. You set up a quartz tube, shine a mercury lamp, and watch the red-brown color slowly disappear as HBr fumes evolve. You collect the product by distillation and hope the NMR doesn't show too much dibrominated material. That is the whole game in a nutshell. What beginners miss about this topic is the selectivity difference between chlorine and bromine radicals. Chlorine is fast and indifferent. It abstracts hydrogens with almost no discrimination, giving a statistical mixture of products. Bromine is slow and picky. It will happily skip a primary carbon and attack a tertiary one even if there are six primary hydrogens waiting. This is the Bell-Evans-Polanyi principle in action. The transition state for bromination has more radical character, so the stability of the resulting carbon radical matters much more. I learned this the hard way during a graduate synthesis when I tried to brominate a substrate with both secondary and tertiary positions and got 60% of the wrong isomer because I used chlorination conditions by habit instead of switching to NBS in CCl4. The workaround was straightforward but costly. I ran a silica column, recycled the major impurity back through the reaction, and accepted a 40% overall yield loss. Never again. Now I check the bond dissociation energies before choosing a halogenating agent every single time.

Here is the practical part that most textbooks gloss over. You need to think about kinetics and thermodynamics simultaneously. The propagation step where the radical abstracts a hydrogen from the substrate is usually rate-determining. The activation energy for this step depends directly on the strength of the C-H bond being broken. A tertiary C-H bond is about 96 kcal/mol while a primary one is closer to 101 kcal/mol. That 5 kcal/mol difference translates into roughly a 1000-fold rate difference at room temperature for bromine abstraction. For chlorine the difference shrinks to maybe 50-fold because the chlorine radical is so reactive that the transition state arrives early and bond strengths matter less. Peroxide initiators like AIBN or benzoyl peroxide decompose thermally to generate the initial radicals. AIBN decomposes at around 65°C with a half-life of roughly 10 hours. That is convenient because it gives you a steady trickle of radicals over a typical reaction period without blowing up the temperature. Benzoyl peroxide works similarly but is a contact explosive when dry. Handle it wet. Store it refrigerated. Replace it every few months because it degrades even in solution and old batches give inconsistent initiation rates. Chain length is another practical consideration you should calculate. In ideal halogenation the chain length can exceed 10,000. That means one photon or one initiating radical event leads to thousands of product molecules before termination stops the chain. In practice, chain lengths drop to 100-500 when you have impurities, oxygen, or polar solvents that intercept radicals. Oxygen is the enemy here. Dissolved O2 reacts with carbon radicals to form peroxy radicals that terminate chains or divert the reaction into hydroperoxide products. If your reaction is supposed to be radical-mediated and it shows no induction period and proceeds slowly from the start, check your solvent for dissolved oxygen. Sparging with argon for 15 minutes usually fixes it. Running the reaction under positive argon pressure with a balloon setup is the lazy but effective version.

Termination steps are bimolecular and second-order in radical concentration. That means at low radical flux the reaction runs efficiently. At high flux, like when you crank up the UV lamp intensity or add too much peroxide, termination dominates and your yield drops. This is why controlled, gentle initiation is better than aggressive initiation. I once used 20 mol% benzoyl peroxide for a radical addition because a paper said "excess initiator ensures complete conversion" and got a tar that wouldn't crystallize. The correct amount was 2 mol%. The paper was wrong or the authors were sloppy or both. One counter-intuitive point about radical reactions: they are often faster in polar solvents than nonpolar ones, which seems to contradict the textbook emphasis on CCl4. The reason is that polar solvents stabilize the polar transition states in certain radical abstractions and can also dissolve reagents better. CCl4 is traditionally used because it is inert and has a high boiling point, not because it is optimal. Modern protocols frequently use acetonitrile or even water for some radical transformations. The classic Hunsdiecker reaction, for example, runs in CCl4 but alternative versions use MeCN with similar results. Bu3SnH is the go-to reagent for radical reductions and cyclizations. It is effective but toxic. Tributyltin compounds accumulate in the environment and are difficult to remove from products. If you are doing a synthesis on scale, consider using silane-based alternatives like PMHS (polymethylhydrosiloxane) with a radical initiator. PMHS is cheap, nontoxic, and easier to work up. The reaction is slower and sometimes requires higher temperatures, but for a multigram synthesis the trade-off is worth it. I switched my lab protocol from Bu3SnH to PMHS/AIBN in toluene at 110°C and cut the workup time from 45 minutes to about 10 minutes per reaction. Product purity improved because tin contamination is gone.

Another practical nuance: radical reactions are sensitive to concentration. Dilute conditions favor intramolecular cyclization over intermolecular polymerization or dimerization. If you are doing a 5-exo or 6-exo radical cyclization, run it at 0.01 to 0.05 M. At 0.5 M you will get significant dimer and trimer byproducts. This is not always obvious from the mechanism drawing. The mechanism shows a clean cycle. The flask shows a mess. Monitoring radical reactions by TLC can be unreliable because radical intermediates are transient and the starting material may disappear via multiple pathways. Running a small aliquot and quenching it with a bit of TEMPO or hydroquinone before spotting on TLC prevents further reaction on the plate. Otherwise you might see a disappearing spot that actually represents continued radical consumption after the sample lands on silica. The bottom line for Free Radicals Organic Chemistry in practice is that mechanism understanding gets you started but controlling the reaction environment gets you product. Temperature, oxygen exclusion, initiator loading, solvent choice, and concentration are all knobs you need to turn deliberately. Pick one, optimize it, move to the next. Do not change two variables at once unless you enjoy analyzing messy chromatograms for three days straight.