Discovery of the structure of benzene

The structure of benzene, a ring of six carbon atoms, was worked out over more than a century. Michael Faraday first isolated the substance in 1825, August Kekulé proposed the ring in 1865, and X-ray crystallography in the late 1920s showed that the ring is flat and symmetrical.

Background

In the early nineteenth century London was lit by oil gas, made by dropping whale or fish oil into a hot furnace and compressing the result into copper cylinders. The gas lost some of its illuminating power over time. In 1825 the Portable Gas Works asked Faraday, working at the Royal Institution, to find out why.

Faraday found that part of the gas condensed into an oily liquid residue. He recorded in his diary on 26 April 1825 that oil gas had been delivered to the Institution, and he began to examine its distillates. On 9 May he visited the gas works himself, and from 16 May the main work began on 211 volumes of oil gas delivered to the Institution. It was a gruelling process of fractional distillation to remove water and other unwanted compounds.

The Royal Institution later described the episode as research without an immediate application that led to an extraordinary legacy.

The discovery

After repeated rounds of freezing, boiling, condensing and compressing, Faraday decided by 18 May that he had isolated something new. In his diary he often called it a “peculiar substance.” He announced it to the Royal Society on 16 June 1825.

His paper does not use the word benzene. Faraday etched the name “bicarburet of hydrogen” on the glass vial that survives at the Royal Institution, because he described its composition as two parts carbon to one part hydrogen by weight. The name benzene took hold later as other chemists found new ways to make the compound, and Faraday approved of it.

The formula is C6H6, and the molar mass is about 78.11 grams per mole. PubChem lists a boiling point of about 80 degrees Celsius. The liquid belongs to a family of compounds that earlier chemists called aromatic because many of them have a sweet, sharp smell.

How the structure was worked out

Chemists could count the atoms but could not explain how six carbon atoms and six hydrogen atoms fit together. Benzene was unusually stable and behaved differently from other compounds with a similar formula.

In 1865 Kekulé suggested that benzene was a regular hexagon with a hydrogen at each corner. He later described the idea as coming from a daydream of a snake seizing its own tail, a story that has become one of the best-known anecdotes in chemistry. He later modified the model to treat benzene as a rapid equilibrium between two cyclohexatriene forms.

Even so, a ring with three fixed double bonds should be highly reactive, and benzene is not. The structure remained unproven for decades. The Royal Institution notes that at benzene’s centenary in 1925 the structure had still not been settled.

The answer came from crystallography. In 1928 a laboratory at the Royal Institution under William Henry Bragg was using X-rays to probe molecular structures. Benzene is a liquid at room temperature, so the team studied hexamethylbenzene, a crystalline relative. Kathleen Lonsdale’s 1929 paper showed that the ring is flat and hexagonal, with all the carbon-carbon bonds the same length. Electrons are spread around the ring in a delocalized cloud, and that delocalization accounts for the stability.

Competing structures

Kekulé’s was not the only proposal. The University of Bristol’s chemistry pages note that Henry Armstrong suggested a “centroid” structure and Albert Ladenburg proposed a prism-shaped one. At the Royal Institution, James Dewar offered further alternatives. Each tried to explain why benzene did not behave like an ordinary molecule with three double bonds. The Lonsdale result supported a symmetric ring with shared electrons rather than any of these alternatives.

Why the ring is stable

The University of Bristol explains that the best description of benzene is indeed hexagonal, with every carbon-carbon bond the same length and intermediate between a single and a double bond. The p-orbitals on neighboring carbon atoms overlap to form a delocalized molecular orbital that extends around the ring. This gives the molecule extra stability, called resonance energy, and makes it much less reactive than a ring with three isolated double bonds would be.

The same source compares benzene with several other arrangements of the same atoms, among them benzvalene, Dewar benzene and prismane. Benzene is the most stable of them, and the energies from highest to lowest differ by almost 546 kilojoules per mole. That large gap helps explain why benzene is the form chemists actually find.

Faraday’s simple compositional analysis and Kekulé’s geometric insight were therefore both partly right. The ring was correct, but the picture of fixed single and double bonds was not.

Impact and legacy

Aromatic chemistry changed industry. In 1856 the eighteen-year-old chemistry student William Henry Perkin was trying to make artificial quinine, a valuable antimalarial drug, from aniline, a coal tar derivative related to benzene. Instead he made a purple dye, mauveine, the first synthetic dye.

Benzene rings now occur in plastics, pharmaceuticals, perfumes and pigments. Toluene, which has one extra methyl group on the ring, and the xylenes, which have two, are both widely used solvents and are found in gasoline. The ring also became a central example in teaching how electrons are shared in a molecule.

Benzene’s carcinogenic effects are now well known. In earlier times it was used to decaffeinate coffee, degrease engines and as an ingredient in shaving cream. Today the pure liquid is limited to controlled industrial processes. This entry is general background rather than advice on handling any chemical.

Editorially reviewed by Specialty Digest Editorial TeamLast reviewed November 5, 2026Researched and drafted with AI assistanceReport an issue