Carbon and its Compounds

Sodium solves the noble-gas problem by simply giving its one spare electron away; chlorine solves it by taking one. Carbon cannot play either game. With six electrons total, arranged 2, 4, carbon has four electrons sitting in its outermost shell, and both obvious ionic moves turn out to be dead ends: losing all four electrons would leave a tiny six-proton nucleus straining to hold onto just two remaining electrons, an enormous amount of energy for very little payoff, while gaining four electrons would leave that same small nucleus straining in the opposite direction, trying to hold ten electrons in place, four more than it has protons to attract. Carbon solves the problem a completely different way: instead of transferring electrons outright, it shares them. Two atoms, each contributing one electron to a shared pair, both count that shared pair toward their own outer shell, and both reach a stable, noble-gas-like arrangement without either atom losing or gaining anything outright. This shared pair is called a covalent bond, and it is the single decision this entire chapter unpacks.

The simplest covalent molecule of all belongs to hydrogen. A hydrogen atom has just one electron in its single K shell, and needs one more to match helium's stable pair, so two hydrogen atoms share their one electron each, forming H2 and a single covalent bond, drawn as a simple line between the two atoms. Chlorine, needing one more electron to complete its octet, does the same to form Cl2. Oxygen needs two more electrons each, so two oxygen atoms share two pairs between them, a double bond. Nitrogen needs three more each, sharing three pairs, a triple bond, the strongest and shortest of the three. Carbon, needing four, reaches methane: one carbon atom shares one electron each with four separate hydrogen atoms, forming four single covalent bonds, CH4, and both carbon and every hydrogen atom finish with a full, stable outer shell. Covalently bonded molecules share something else too: the bond holding the two atoms of one molecule together is strong, but the forces attracting one whole molecule to its neighbour are comparatively weak, which is exactly why carbon compounds, methane among them, tend to have low melting and boiling points and, since no ions are ever formed by sharing, are generally poor conductors of electricity.

Chemists have catalogued millions of known carbon compounds, a number that dwarfs the compounds formed by every other element on the periodic table put together, and the reason comes down to two specific properties of carbon working together. The first is catenation: carbon's unusual ability to bond with other carbon atoms, again and again, building long chains, branched chains, or closed rings, held together by carbon-carbon bonds strong and stable enough to support truly long structures. Silicon can form similar chains, but only up to seven or eight atoms before the compound becomes too reactive to survive; carbon has no such ceiling. The second is tetravalency: with four electrons available to share, each carbon atom can bond to up to four other atoms, whether more carbon, or hydrogen, oxygen, nitrogen, sulphur, chlorine, almost anything, and carbon's comparatively small atomic size lets its nucleus hold on to every one of those shared electron pairs unusually tightly, making the resulting bonds strong and the compounds stable. For a long time, chemists believed carbon compounds extracted from living things carried something beyond ordinary chemistry, a mysterious "vital force" that only a living organism could supply, and that these organic compounds could never be built from scratch in a laboratory. Friedrich Wöhler shattered that belief in 1828, synthesising urea, a genuine biological waste product, directly from ammonium cyanate, a compound with no biological origin at all. The vital force theory did not survive the discovery; the name "organic chemistry" for the study of carbon compounds did.

Take two carbon atoms and link them with a single bond, then use hydrogen atoms to fill every remaining valency, and the result is ethane, C2H6, every single valency of every atom satisfied by a single bond. Compounds built entirely this way, with carbon atoms connected only by single bonds, are called saturated compounds, and they tend to be comparatively unreactive. But carbon and hydrogen can combine differently: C2H4, ethene, needs a double bond between its two carbons before every valency balances out, and C2H2, ethyne, needs a triple bond. Compounds containing a carbon-carbon double or triple bond are called unsaturated compounds, and they are noticeably more reactive than their saturated cousins, precisely because that extra shared pair (or two) of electrons in the double or triple bond is available to react further. Compounds built only from carbon and hydrogen, saturated or not, are called hydrocarbons, and this single distinction sorts them into three families: saturated hydrocarbons are alkanes, unsaturated ones with at least one double bond are alkenes, and unsaturated ones with at least one triple bond are alkynes.

Line up carbon atoms one after another and the chain can run to almost any length: methane with one carbon, ethane with two, propane with three, butane with four, all the way up through pentane and hexane, a genuine family with a clear naming pattern. But length is not the only choice available. Take butane's four carbon atoms and there are two different ways to arrange the same skeleton: one straight chain, C-C-C-C, and one branched, with a shorter chain hanging off the middle. Both skeletons, once filled out with exactly enough hydrogen atoms to satisfy every valency, share the exact same molecular formula, C4H10, yet they are genuinely different molecules with different structures and different properties. Compounds sharing one formula but differing in structure like this are called structural isomers, and the possibilities only grow as chains get longer. Carbon chains do not have to stay open either: cyclohexane, C6H12, closes its six carbons into a ring instead of a line, and benzene, C6H6, forms a similarly closed ring, a foundational structure across an enormous share of organic chemistry. Straight, branched, or ringed, any of these skeletons can turn out saturated or unsaturated, the shape of the skeleton and the nature of its bonds are two entirely separate questions.

Replace one hydrogen atom in a hydrocarbon chain with a different atom or group, and something disproportionate happens: the whole compound's behaviour can change completely, no matter how long or short the carbon chain underneath happens to be. The replacing atom is called a heteroatom, and specific heteroatoms or small groups of atoms attached this way are called functional groups, since they define a compound's characteristic function almost entirely on their own. Chlorine or bromine replacing a hydrogen gives a haloalkane. An -OH group gives an alcohol. Other common groups give an aldehyde, a ketone, or a carboxylic acid, each with its own distinct chemistry. This single idea, that a small attached group can dictate a compound's properties almost independently of the chain it sits on, is what makes organic chemistry navigable at all despite the millions of individual compounds involved: learn how one functional group behaves, and that knowledge transfers to every compound carrying that same group, regardless of chain length.

Methanol, ethanol, propanol and butanol behave almost identically in every chemical reaction, despite having entirely different carbon chain lengths, because every one of them carries the exact same functional group, -OH. Line them up by chain length and something else becomes visible: each successive member differs from the one before it by exactly one -CH2- unit, CH3OH to C2H5OH, C2H5OH to C3H7OH, and so on. A family of compounds built this way, one functional group, an unbroken sequence of chain lengths, each step apart by a single -CH2- unit, is called a homologous series, and the same pattern holds across every functional group, alkanes, alkenes, alcohols, carboxylic acids, each has its own homologous series. Within any one series, physical properties shift gradually and predictably as the chain lengthens, melting points and boiling points climb, solubility patterns shift, simply because the molecules are getting bigger and heavier. Chemical properties, though, barely move at all within a series, since chemical behaviour is decided almost entirely by the functional group, and every member of the series carries the identical one.

A name like propanol carries real, decodable information, once the pattern behind it is known. Count the carbon atoms in the chain first: three carbons gives the base name propane. Then note whether a functional group is present, and attach it as a prefix or suffix: a halogen becomes a prefix, chloro- or bromo-; most other functional groups become a suffix instead. Where the suffix begins with a vowel, the base name's final "e" is dropped before attaching it, propane minus "e" plus "one" gives propanone, the systematic name for a three-carbon ketone. An unsaturated chain swaps its ending too, "ane" becomes "ene" for a double bond or "yne" for a triple bond, so propene is a three-carbon compound with one double bond, and propyne the equivalent with a triple bond. Alcohols take the suffix "-ol", aldehydes "-al", carboxylic acids "-oic acid". Once this system is internalised, a name like butanoic acid or chloropropane stops being a label to memorise and becomes something that can be decoded, or built, directly from first principles.

Carbon, in every one of its forms, burns in oxygen to release carbon dioxide along with heat and light, C + O2 -> CO2 + heat and light, and most compounds of carbon burn the same generous way, methane and ethanol both releasing plenty of heat and light as they oxidise completely. But not every carbon flame looks the same. Saturated hydrocarbons tend to burn with a clean flame; unsaturated ones burn with a yellow, visibly sooty flame, leaving a telltale black deposit on any metal plate held above it, since there is not quite enough oxygen reaching every carbon atom to convert it fully to carbon dioxide. Even a saturated fuel produces the same sooty, incomplete result if its air supply is restricted, which is exactly why a gas or kerosene stove is built with adjustable air holes: enough air mixed in gives a clean blue flame and full, efficient combustion, while blocked or restricted air holes waste fuel and blacken the bottom of every pot on the stove. Coal and petroleum, being ancient, compressed remains of long-dead plants and sea life, fossil fuels, also carry small amounts of nitrogen and sulphur, and burning them releases oxides of both, a major source of the sulphur and nitrogen oxide pollution found in city air today.

Warm ethanol gently and add alkaline potassium permanganate drop by drop, and the purple colour disappears at first, only persisting once excess has been added. What is happening is oxidation: potassium permanganate (or, just as commonly, acidified potassium dichromate) is adding oxygen to the ethanol, converting it step by step into ethanoic acid, which is exactly why substances capable of donating oxygen this way are called oxidising agents. A second, quite different kind of change is addition: unsaturated hydrocarbons, their double or triple bonds still hungry for more, will add hydrogen directly across that extra bond in the presence of a catalyst such as palladium or nickel, converting themselves into saturated compounds. This exact reaction, run at industrial scale with a nickel catalyst, is how liquid vegetable oils, built from long unsaturated carbon chains, are hydrogenated into solid fats; animal fats, largely saturated already, do not need this step, and the unsaturated fats health advice recommends are simply oils that have skipped it. A third kind of change belongs to saturated hydrocarbons, which are otherwise fairly unreactive: exposed to sunlight, chlorine gas reacts with them anyway, replacing hydrogen atoms one at a time, CH4 + Cl2 -> CH3Cl + HCl. Because one atom is simply swapped for another rather than anything being added or removed overall, this is called a substitution reaction.

Ethanol, commonly called alcohol, is a liquid at room temperature, fully soluble in water in any proportion, and it is the active ingredient of every alcoholic drink, as well as a genuinely useful solvent in medicines like tincture iodine and cough syrup. Drop a small piece of sodium into ethanol and hydrogen gas bubbles off immediately, 2Na + 2CH3CH2OH -> 2CH3CH2ONa + H2, sodium ethoxide left behind in solution, the same kind of hydrogen-releasing reaction sodium has with water. Heat ethanol instead, to 443 K with an excess of concentrated sulphuric acid acting as a dehydrating agent, and it loses a water molecule entirely, converting to ethene, unsaturated and reactive where ethanol was neither. Ethanol's effect on the body is worth being precise about, since "alcohol" covers more than one compound with very different risks. Ethanol itself, even in small amounts, slows the central nervous system and impairs coordination and judgement, and pure, undiluted ethanol (absolute alcohol) can be lethal. Methanol is far more dangerous still: the liver oxidises it into methanal, which coagulates proteins in living cells much as heat coagulates an egg, and even a small quantity can cause blindness or death. This is precisely why industrial ethanol, exempt from the taxes and controls on drinking alcohol, is deliberately laced with methanol and dyed blue, denatured alcohol, poisonous and clearly marked, specifically to prevent it being consumed.

Ethanoic acid, commonly called acetic acid, belongs to the carboxylic acid family, identifiable by its acidic behaviour, though unlike a mineral acid such as hydrochloric acid, which ionises completely, carboxylic acids are weak acids, only partly ionising in water. A 5-8% solution of it in water is ordinary vinegar, widely used to preserve pickles. Its pure form freezes at 290 K, cold enough to solidify in a chilly winter, which is exactly why it earned the older name glacial acetic acid. Warm ethanoic acid with ethanol and a few drops of concentrated sulphuric acid as a catalyst, and the two combine into a sweet-smelling ester, CH3COOH + C2H5OH -> CH3COOC2H5 + H2O, esterification, the very reaction behind many perfumes and flavourings. Run that same ester back against sodium hydroxide and it splits apart again into alcohol and the sodium salt of the acid, saponification, named for exactly what it produces: soap. Ethanoic acid reacts with an ordinary base too, forming a salt and water, NaOH + CH3COOH -> CH3COONa + H2O, and with carbonates and hydrogencarbonates it releases carbon dioxide directly, 2CH3COOH + Na2CO3 -> 2CH3COONa + H2O + CO2, a gas easily confirmed by turning limewater milky, exactly the same confirming test used for carbon dioxide throughout this entire subject.

Shake oil into a test tube of plain water and it barely mixes at all, separating back into its own layer within moments; add a few drops of soap solution first and the oil stays mixed for noticeably longer. Soap molecules are sodium or potassium salts of long-chain carboxylic acids, and each molecule has a genuinely split personality: a long hydrocarbon tail that dissolves happily in oil but not in water, and a charged, ionic head that dissolves happily in water but not in oil. Dropped into a mixture of oil and water, soap molecules arrange themselves around each oil droplet with their hydrocarbon tails pointing inward, toward the oil, and their ionic heads facing outward, toward the surrounding water, forming a cluster called a micelle. Wrapped this way, the oil droplet, and whatever dirt was trapped inside it, becomes soluble enough in water to rinse cleanly away, which is the entire cleaning mechanism of soap in one sentence. Hard water complicates this: calcium and magnesium ions react with ordinary soap to form an insoluble, curdy scum, wasting soap and leaving foam hard to raise, which is exactly the problem detergents were engineered to solve. Detergents, typically sodium salts of sulphonic acids or ammonium salts with a halide, keep the same hydrocarbon-tail-plus-ionic-head structure that makes micelles work, but their particular ionic ends simply do not form an insoluble precipitate with calcium or magnesium, letting them lather and clean effectively even in the hardest water.