Organic chemistry began as the chemistry of life and, for a long time, many scientists believed that organic substances could only arise within living organisms by a mysterious “vital force.” That belief collapsed in 1828 when Friedrich Wöhler heated ammonium cyanate and obtained urea, a bona fide organic molecule, without any biological help. In one elegant experiment, Wöhler showed that organic compounds obey the same physical laws as all other substances, opening the floodgates to modern organic synthesis, pharmaceuticals, materials, and biochemistry.
At the heart of this branch of chemistry sits carbon, an element that behaves like a molecular Lego. Its two headline traits—tetravalency and catenation—explain the size, diversity, and architectural richness of organic molecules. Tetravalency means a neutral carbon atom forms four covalent bonds, typically arranged toward the corners of a tetrahedron. Catenation means carbon atoms bond to each other in long chains, branched networks, and rings. When these two features combine, the result is an almost unlimited catalog of structures that can host non-carbon atoms or functional groups such as hydroxyl, carbonyl, carboxyl, amino, nitro, and halogen substituents. From the simplest alkane to DNA and graphene, the theme is the same: four bonds and a taste for self-linking give rise to the chemistry of life and of modern materials.
What counts as an organic compound and what does not
In today’s broad usage, an organic compound is any compound whose framework is built from carbon, usually bonded to hydrogen and often to oxygen, nitrogen, sulfur, phosphorus, or halogens. Classic counter-examples that are generally treated in inorganic or physical chemistry include carbides, the elemental oxides of carbon (CO and CO₂), carbonate and hydrogen carbonate salts, and simple cyanides of alkali metals. Everything else—from hydrocarbons and alcohols to amino acids, sugars, pharmaceuticals, dyes, polymers, and fuels—sits squarely within organic chemistry.
The big families at a glance
Hydrocarbons contain only carbon and hydrogen. Alkanes are saturated, alkenes contain C=C double bonds, and alkynes host C≡C triple bonds. Aromatic hydrocarbons such as benzene have cyclic, conjugated π systems with characteristic stability and substitution patterns.
Alcohols and phenols include a hydroxyl group; phenols differ by having the hydroxyl directly on an aromatic ring, which makes them more acidic and more reactive in electrophilic substitution. Carboxylic acids possess the –COOH group and can donate a proton, forming carboxylate salts; their derivatives include esters, acid chlorides, and amides. Aldehydes and ketones share the carbonyl group; aldehydes have at least one hydrogen attached to the carbonyl carbon, while ketones have two carbon substituents. Sugars are polyhydroxy aldehydes or ketones with rich stereochemistry and the ability to cyclize. Nitrogen compounds span amines, amides, nitriles, nitro compounds, and urea; they are central to biomolecules, materials, and catalysis.
Why thermodynamic data matters for organic molecules
A table of thermodynamic quantities is a treasure map for chemists. Four columns are particularly useful for organic compounds: the standard enthalpy of combustion, the standard enthalpy of formation, the standard Gibbs energy of formation, and the standard molar entropy. Each number answers a different practical question, and together they let you compare fuels, forecast reaction spontaneity, and estimate equilibrium positions.
The enthalpy of combustion measures the heat released when one mole of a compound burns completely in oxygen to form CO₂ and H₂O under standard conditions. It is negative because the process is exothermic. For fuels, the magnitude of this number per unit mass or per unit volume is the energy density: larger magnitude means more heat available. Methane’s value of roughly −890 kJ mol⁻¹ makes it an efficient, clean fuel on a per-carbon basis; octane’s much larger molar value reflects its greater number of C–C and C–H bonds, which liberate more heat upon oxidation.
The enthalpy of formation tells you how much heat is absorbed or released when one mole of a substance forms from its elements in their standard states. Many common organic molecules have negative formation enthalpies, signalling that, relative to their elements, they are thermodynamically stable. Positive numbers indicate a compound sits “uphill” from its elements and may prefer to decompose unless kinetic barriers intervene.
The Gibbs energy of formation folds enthalpy and entropy together, answering the question of overall thermodynamic favorability at a given temperature. Negative values indicate that making the compound from its elements is spontaneous under standard conditions. In reaction calculations, you sum the ΔG°f values of products minus reactants to estimate ΔG° for the process and then link that to the equilibrium constant with ΔG° = −RT ln K.
Entropy measures dispersal of energy and matter. Gas-phase molecules tend to have higher molar entropies than liquids or solids because their translational and rotational motions are less constrained. Comparing entropies teaches you why many combustions are entropy-favorable: a condensed organic fuel and oxygen become gaseous CO₂ and water vapor at flame temperatures, creating more accessible microstates.
Reading the numbers like a pro
When two isomeric hydrocarbons are compared, the one with more branching generally has a less negative enthalpy of combustion per mole because branching slightly stabilizes the molecule relative to its fully oxidized products. This detail explains why gasoline formulators like branched alkanes: they resist knocking and deliver controlled energy release. Adding oxygen to a molecule, as in alcohols, tends to make the enthalpy of combustion per carbon less negative because the oxygenated carbon is already partially oxidized. Thus ethanol yields less heat per kilogram than the hydrocarbon it may replace, even if it burns cleaner.
A negative enthalpy of formation for ethanol or benzene means that, at 298 K, the hypothetical process of assembling the compound from elemental carbon, hydrogen, and oxygen would release heat. A compound with a strongly negative Gibbs energy of formation is thermodynamically stable with respect to its elements, yet that tells you nothing about reaction speed. Diamond converts to graphite in the thermodynamic ledger, but kinetic barriers keep diamonds intact on human timescales; many energetic organic molecules exhibit a similar enthalpy–kinetics tension.
From data to design: fuels, pharmaceuticals, and processes
Combustion enthalpies help you rank and select fuels for engines, rockets, and domestic use while anticipating CO₂ footprints. Gibbs energies guide synthetic chemists toward reactions that are thermodynamically downhill or toward conditions that shift equilibrium in their favor, such as removing a product water to drive esterification. Formation entropies and enthalpies underpin vapor-liquid equilibria in distillation design, while entropy insights explain why gas-evolving reactions are often pulled forward at high temperature. Even in drug discovery, relative stabilities and hydrogen-bond strengths echo through binding thermodynamics, where favorable enthalpies and entropies must cooperate to deliver a potent, selective ligand.
Worked example: estimating heat from burning methane and ethanol
Suppose you burn one mole of methane and one mole of ethanol completely in oxygen at standard conditions. Using representative values, methane releases about 890 kJ mol⁻¹ and ethanol around 1368 kJ mol⁻¹. Although ethanol gives off more heat per mole, methane provides more energy per gram because its molar mass is only 16 g, whereas ethanol’s is 46 g. On a per-mass basis, methane’s specific energy is roughly 55.6 kJ g⁻¹, while ethanol’s is about 29.7 kJ g⁻¹. This simple calculation shows why natural gas is such an efficient fuel and why volumetric energy density and storage logistics also matter in real systems.
Practical interpretation of the data shown
A representative organic data sheet lists dozens of compounds. Hydrocarbons such as methane, ethyne, ethene, ethane, and benzene show large negative combustion enthalpies, a clear fingerprint of their use as fuels. Alcohols and phenols display negative formation enthalpies and moderate entropies in the liquid state, reflecting hydrogen bonding and partial ordering. Carboxylic acids like formic and acetic acid are even more strongly stabilized by resonance and hydrogen bonding, which is consistent with their relatively high boiling points. Aldehydes and ketones such as methanal and ethanal present negative formation enthalpies and higher gaseous entropies than their liquid counterparts, illustrating general phase trends. Sugars, including glucose, fructose, and sucrose, have massively exothermic combustion enthalpies because they are packed with oxidizable C–H and C–C bonds, even though their formation from elements is also exergonic under standard biochemical conventions. Nitrogen compounds such as urea and amines exhibit distinctive thermodynamic signatures due to the presence of polar N–H bonds and possibilities for hydrogen bonding and ionic association.
Using Hess’s law and formation data to compute reaction enthalpies
Hess’s law states that enthalpy is a state function; the heat effect of a reaction depends only on initial and final states, not on the path. To compute the enthalpy change for an organic reaction, you sum the formation enthalpies of the products and subtract the sum for reactants. Consider the hydrogenation of ethene to ethane. With typical formation enthalpies of about +52.26 kJ mol⁻¹ for ethene and −84.68 kJ mol⁻¹ for ethane, and zero for H₂ in its standard state, the reaction enthalpy comes out near −136.94 kJ mol⁻¹, revealing why hydrogenation is exothermic and why catalysts are required not to make it possible, but to make it fast and selective.
Safety and environmental context
Combustion is exothermic and oxygen-hungry; adequate ventilation and flame control are essential in any lab or industrial setting. Thermodynamic favorability does not guarantee selectivity; incomplete combustion yields toxic CO, soot, and volatile organics. The same tables that help you design a fuel blend also help you estimate heat loads for safety valves and cooling jackets. Sustainable chemistry aims to pair favorable thermodynamics with minimal hazard and waste, replacing stoichiometric oxidants with air or green oxidants, and using catalysts that lower energy demand while suppressing by-products.
Reference mini-tables drawn from the data
The following condensed tables reproduce a few representative entries so you can see how to read and apply them. Values are standard-state, 298 K approximations in the units shown.
Representative hydrocarbons
| Substance | ΔcH° (kJ mol⁻¹) | ΔfH° (kJ mol⁻¹) | ΔfG° (kJ mol⁻¹) | S° (J K⁻¹ mol⁻¹) |
|---|---|---|---|---|
| Methane, CH₄(g) | −890 | −74.81 | −50.72 | 186.26 |
| Ethene, C₂H₄(g) | −1411 | 52.26 | 68.15 | 219.56 |
| Ethane, C₂H₆(g) | −1560 | −84.68 | −32.82 | 229.60 |
| Benzene, C₆H₆(l) | −3268 | 49.0 | 124.3 | 173.3 |
Selected oxygenates
| Substance | ΔcH° (kJ mol⁻¹) | ΔfH° (kJ mol⁻¹) | ΔfG° (kJ mol⁻¹) | S° (J K⁻¹ mol⁻¹) |
|---|---|---|---|---|
| Ethanol, C₂H₅OH(l) | −1368 | −277.69 | −174.78 | 160.7 |
| Phenol, C₆H₅OH(s) | −3054 | −164.6 | −50.42 | 144.0 |
| Formic acid, HCOOH(l) | −255 | −424.72 | −361.35 | 128.95 |
| Acetic acid, CH₃COOH(l) | −875 | −484.5 | −389.9 | 159.8 |
Sugars and nitrogen compounds
| Substance | ΔcH° (kJ mol⁻¹) | ΔfH° (kJ mol⁻¹) | ΔfG° (kJ mol⁻¹) | S° (J K⁻¹ mol⁻¹) |
|---|---|---|---|---|
| Glucose, C₆H₁₂O₆(s) | −2808 | −1268 | −910 | 212 |
| Sucrose, C₁₂H₂₂O₁₁(s) | −5645 | −2222 | −1545 | 360 |
| Urea, CO(NH₂)₂(s) | −632 | −333.51 | −197.33 | 104.60 |
| Methylamine, CH₃NH₂(g) | −1085 | −22.97 | 32.16 | 243.41 |
These numbers immediately support practical conclusions. Ethene’s positive formation enthalpy and Gibbs energy relative to its elements reflect the stored chemical potential in a double bond, which is released upon hydrogenation or combustion. Glucose and sucrose exhibit very large negative combustion enthalpies because extensive oxidation to CO₂ and H₂O yields strong C=O and O–H bonds, liberating substantial heat.
Carbon’s bonding toolkit and why it matters to thermodynamics
The same hybridization patterns that determine structure also influence thermodynamic data. sp³ carbons in alkanes are relatively low in strain and yield predictable combustion heats. Introducing sp² hybridization in alkenes raises enthalpy of formation because π bonds are weaker than σ bonds, even though conjugation can delocalize electrons and recover stability, as seen in benzene’s special resonance stabilization. Strained rings, such as cyclopropane, store energy that appears as a less negative combustion enthalpy per carbon and a higher formation enthalpy; relieving angle strain on reaction provides a thermodynamic push.
From Wöhler’s urea to modern organic frontiers
Wöhler’s preparation of urea from an inorganic salt shattered the vitalism myth and created a unified view of chemistry. Today, the same principles extend to organometallic catalysis that stitches carbon chains together with almost surgical precision, to polymer chemistry that transforms simple monomers into durable materials, and to green chemistry that optimizes enthalpy and entropy to minimize energy use and waste. The periodic table has not changed since 1828, but our command of thermodynamics and kinetics lets us assemble molecules with properties tuned for medicine, electronics, agriculture, and sustainability.
Frequently asked questions
Why are combustion enthalpies of hydrocarbons so negative?
Burning a hydrocarbon replaces relatively weaker C–C and C–H bonds with very strong C=O bonds in CO₂ and O–H bonds in H₂O. The formation of these stronger bonds releases large amounts of heat, so the enthalpy of combustion is strongly exothermic.
How do ΔfH° and ΔfG° differ in meaning?
The enthalpy of formation measures heat flow at constant pressure when forming a compound from its elements, while the Gibbs energy of formation includes entropy, telling you the overall spontaneity. A compound can have a negative ΔfH° but, if entropy is unfavorable, a less negative or even positive ΔfG°.
Does a negative ΔfG° mean a compound is reactive or unstable?
No. Negative ΔfG° means the compound is thermodynamically favored versus its elements. Reactivity depends on kinetics—activation barriers and mechanisms. Stable compounds can react vigorously if initiated, and thermodynamically unstable ones can persist if barriers are high.
Why does branching lower the magnitude of the enthalpy of combustion per mole?
Branching stabilizes alkanes slightly through hyperconjugation and reduced strain. A more stable starting molecule sits closer in energy to its combustion products, so the heat released is slightly less negative than for a straight-chain isomer of the same formula.
What did Wöhler actually prove with urea?
By converting ammonium cyanate to urea, Wöhler demonstrated that an organic compound characteristic of living systems could be synthesized from inorganic precursors, disproving the vital force theory and launching modern organic synthesis.
Are sugars good fuels if their combustion enthalpy is large?
Per mole, sugars release substantial heat, but they contain oxygen already and have high polarity, reducing energy density per kilogram compared with hydrocarbons. They are excellent biological fuels because enzymes can extract energy stepwise; as industrial fuels, their moisture and processing costs often limit use.
How can I use these tables in exams and labs?
Use ΔfH° to compute reaction enthalpies via Hess’s law, ΔfG° to estimate equilibrium constants and spontaneity, ΔcH° to compare fuels, and S° to justify temperature effects and gas-evolving reaction tendencies. Always specify standard state and temperature.
Why do phenols behave differently from alcohols if both have –OH groups?
Phenols have the hydroxyl attached to an aromatic ring, enabling resonance stabilization of the phenoxide ion. This makes phenols more acidic, raises certain substitution reactivities, and influences thermodynamic properties such as boiling point and formation enthalpy.
Which compounds are usually excluded from “organic” in textbooks?
Carbon oxides (CO, CO₂), carbonates and hydrogen carbonates, and many carbides are typically treated outside organic chemistry, even though they contain carbon. Organic chemistry focuses on carbon frameworks bonded mainly to hydrogen and heteroatoms in covalent networks.
How do entropy values help in mechanism thinking?
Reactions that increase the number of gas molecules usually gain entropy and become more favorable at higher temperatures. Conversely, cyclization or association processes that reduce particle count often suffer entropy penalties and may need enthalpic help or low-temperature conditions.
