An alkane is a hydrocarbon, a molecule made entirely of carbon and hydrogen, where every bond between atoms is a single bond. This makes alkanes the simplest family of organic compounds. They follow a predictable formula: for any straight or branched alkane with n carbon atoms, there are exactly 2n + 2 hydrogen atoms, written as CnH2n+2. Methane (CH4), the main component of natural gas, is the smallest alkane. From there, the family scales up one carbon at a time: ethane, propane, butane, and so on.
Why Alkanes Are Called “Saturated”
You’ll often see alkanes described as saturated hydrocarbons. “Saturated” means every carbon atom holds as many hydrogen atoms as it possibly can. There are no double or triple bonds between carbons, just single bonds throughout. Each carbon forms four bonds total, connecting to some combination of other carbons and hydrogens. This arrangement gives each carbon atom a roughly tetrahedral shape, with bond angles close to 109.5 degrees.
Compare this to alkenes, which contain at least one carbon-carbon double bond, or alkynes, which have a triple bond. Those unsaturated compounds are more chemically reactive because their double and triple bonds can open up and add new atoms. Alkanes, fully loaded with hydrogen, don’t have that option, which is why they tend to be far less reactive.
How Alkanes Are Named
The first four alkanes have traditional names: methane (1 carbon), ethane (2), propane (3), and butane (4). From five carbons onward, the names come from Greek number prefixes: pentane, hexane, heptane, octane, nonane, decane.
For branched alkanes, the IUPAC naming system follows a few key steps. First, find the longest continuous chain of carbon atoms. That chain determines the base name. Then identify any branches (called alkyl groups) hanging off the main chain and number the chain from whichever end gives those branches the lowest possible position numbers. A six-carbon chain with a one-carbon branch at the second position, for example, becomes 2-methylhexane.
When multiple identical branches exist, prefixes like di-, tri-, and tetra- indicate how many. If different types of branches are present, they’re listed alphabetically. One common mistake: the longest chain doesn’t always run in a straight horizontal line. It can twist and turn through the structure, so you need to trace carefully.
Structural Isomers
Starting at four carbon atoms, alkanes can be arranged in more than one way while keeping the same molecular formula. These different arrangements are called structural isomers. Butane (C4H10) has 2 isomers: the straight chain and a branched version. Pentane (C5H12) has 3, and hexane (C6H14) has 5. The number of possible isomers grows rapidly with each additional carbon. These isomers aren’t just theoretical curiosities. They have slightly different physical properties, like different boiling points, which matters in applications like fuel refining.
Cycloalkanes
When a chain of carbon atoms loops back on itself to form a ring, the result is a cycloalkane. Closing the ring means losing two hydrogen atoms compared to the straight-chain version, so cycloalkanes follow the formula CnH2n. Cyclopentane and cyclohexane are the most common examples.
Small rings come with a problem called ring strain. Carbon atoms prefer bond angles of about 109.5 degrees, but in cyclopropane (a three-carbon ring) the angles are forced down to 60 degrees, and in cyclobutane they’re squeezed to about 90 degrees. That mismatch stores extra energy in the molecule, making small-ring cycloalkanes less stable and more reactive than their larger-ring relatives. Cyclohexane, with six carbons, can flex into a shape that achieves nearly ideal bond angles, so it’s quite stable.
Physical Properties
Alkanes are nonpolar molecules. The difference in how strongly carbon and hydrogen attract electrons is tiny, so carbon-hydrogen bonds carry almost no electrical charge. The only forces pulling alkane molecules toward each other are weak London dispersion forces, which arise from temporary fluctuations in electron distribution. These weak attractions explain two key trends.
First, boiling and melting points rise steadily with chain length. Methane boils at -164°C, making it a gas at room temperature. Ethane boils at -89°C, propane at -42°C, and butane at -1°C. All gases. Pentane, boiling at 36°C, is the first alkane that’s liquid under normal conditions. By octane (8 carbons, boiling point 125°C) and decane (10 carbons, 174°C), you’re firmly in liquid territory. Longer chains mean more surface area for those London dispersion forces to act on, so it takes more heat to pull the molecules apart.
Second, alkanes are virtually insoluble in water. Water molecules are held together by strong hydrogen bonds. For an alkane to dissolve, it would need to break into those hydrogen bonds and replace them with much weaker London dispersion forces. The energy doesn’t balance out, so water and alkanes simply don’t mix. This is why oil (mostly long-chain hydrocarbons) floats on water. Alkanes dissolve readily in organic solvents, though, because those solvents are also held together by London dispersion forces. Swapping one set of weak attractions for another costs almost no energy.
Chemical Reactions
Alkanes are among the least reactive organic compounds. Their single bonds are strong and their nonpolar nature gives other reagents little to grab onto. But two reactions stand out.
Combustion is by far the most important. When an alkane burns in oxygen, every bond in the molecule breaks and reforms into carbon dioxide and water, releasing a large amount of heat. Propane burning, for instance, produces three molecules of CO2 and four of water per molecule of fuel. This is the reaction behind everything from gas stoves to car engines. If oxygen supply is limited, the combustion is incomplete and produces carbon monoxide instead of CO2, which is toxic.
Halogenation is the other notable reaction. In the presence of energy (usually UV light or heat), a halogen like chlorine can replace one of the hydrogen atoms on an alkane. The reaction proceeds through a chain mechanism involving highly reactive fragments called free radicals. One chlorine molecule splits into two reactive chlorine atoms, each of which can pull a hydrogen off the alkane and set off a chain of further reactions. This is one of the primary ways chemists convert unreactive alkanes into more useful starting materials.
Everyday Uses
Alkanes are the backbone of the fossil fuel economy. Natural gas is 70 to 90 percent methane, with smaller amounts of ethane, propane, and butane. It heats homes, generates electricity, and fuels industrial processes. Propane and butane can be compressed into liquids at relatively low pressures, which is why they work well in portable tanks for grills and in butane cigarette lighters.
Gasoline is a blend of liquid alkanes, primarily in the five-to-twelve carbon range, with octane being the most well-known component. Longer chains, from about 17 to 35 carbons, form the basis of lubricating oils. Solid alkanes with even longer chains show up as paraffin wax in candles. Beyond fuel and lubrication, alkanes serve as raw materials for the chemical industry. Through reactions like the halogenation described above, they’re converted into the building blocks for plastics, solvents, and countless other products.
Methane and Climate
Methane, the simplest alkane, is also a potent greenhouse gas. Over a 100-year period, the EPA estimates methane traps 27 to 30 times more heat in the atmosphere than the same mass of carbon dioxide. Major sources include natural gas leaks, livestock digestion, landfills, and wetlands. Because methane breaks down in the atmosphere faster than CO2, its short-term warming effect is even more dramatic, making it a key target in efforts to slow climate change.