The Functional Group: A Carbon Wearing an -OH
You have already met dozens of functional groups in passing; now we settle in with one of the most important in all of chemistry and biology. An alcohol is simply any molecule with a hydroxyl group, an -OH, bonded to a saturated (sp3) carbon. That one detail — saturated carbon — is the whole definition, and it matters: the carbon must be a plain tetrahedral carbon, not part of a C=O (that would be a carboxylic acid) and not part of a benzene ring (that, as we will see, is a phenol, a different creature). So ethanol is CH3CH2OH, the alcohol in every glass of wine; methanol is CH3OH; and the simplest alcohol-bearing ring is cyclohexanol.
Picture the oxygen up close. It is sp3-hybridized, wearing two lone pairs and two single bonds in a bent shape much like water — in fact, an alcohol is essentially a water molecule with one hydrogen swapped for a carbon chain. Because oxygen is far more electronegative than either carbon or hydrogen, it pulls electron density toward itself, leaving both the O-H bond and the C-O bond strongly polar. That little wedge of negative charge on oxygen and positive charge on the attached hydrogen is the seed of nearly everything an alcohol does: its solubility, its boiling point, and the two reactive handles it offers — a slightly acidic O-H proton and a pair of nucleophilic lone pairs.
Hydrogen Bonding: Why Alcohols Boil High and Mix with Water
Recall from earlier rungs that the strongest of the ordinary intermolecular forces is the hydrogen bond: a special, extra-strong attraction that appears when an H attached to N, O, or F reaches over to a lone pair on a neighboring N, O, or F. An alcohol's O-H is the textbook hydrogen-bond donor, and its oxygen lone pairs are equally good acceptors. So in a flask of pure ethanol the molecules are not loose strangers — they are stitched together into a shifting, three-dimensional web, each O-H clasping a neighbor's oxygen. This is what we call alcohol hydrogen bonding, and it explains two of the most striking facts about alcohols.
First, boiling point. To boil a liquid you must tear its molecules apart into a gas, and that web of hydrogen bonds is extra tape holding ethanol's molecules down. Ethanol boils at 78 C, while propane — almost the same size and weight but with no O-H to hydrogen-bond — boils at a frigid -42 C. Even dimethyl ether, which has the very same formula as ethanol (C2H6O) but tucks its oxygen between two carbons so it has no O-H donor, boils at -24 C. Same atoms, but the molecule that can donate a hydrogen bond boils a hundred degrees higher. That is not a small effect; it is the difference between a pourable liquid and a gas at room temperature.
Second, solubility in water. An alcohol can hydrogen-bond to water just as happily as to itself, so water welcomes it. But the carbon chain is a freeloader — it cannot hydrogen-bond and only disrupts water's own network, so the longer the chain, the more it drags the molecule toward oil-like insolubility. The result is a clean tug-of-war: methanol, ethanol, and propanol mix with water in all proportions, but by the time you reach 1-hexanol the long greasy tail wins and solubility collapses. A handy rule of thumb is that an alcohol stays appreciably water-soluble up to about four or five carbons per -OH group; beyond that the hydrocarbon part dominates.
Naming and Classifying: 1, 2, 3
Naming an alcohol uses the same IUPAC machinery you already practiced on alkanes, with one new wrinkle: the suffix changes from -ane to -ol, and the -OH group claims top priority for the lowest locant. So you find the longest parent chain that includes the carbon bearing the -OH, number it from whichever end gives that carbon the smaller number, and append -ol with that number. CH3CH2OH is ethanol; the OH on the second carbon of a four-carbon chain is butan-2-ol; a three-carbon chain branching with a methyl is 2-methylpropan-1-ol. The OH outranks double bonds and halogens for the low number — it is the senior group here.
More important than the name is the classification, because it predicts how the alcohol will react in later guides. The label comes straight from the carbon classification you already know: look only at the carbon that carries the -OH, and count how many other carbons are attached to it. Zero or one other carbon makes it a primary (1) alcohol; two makes it secondary (2); three makes it tertiary (3). This is alcohol classification, and it is not bookkeeping for its own sake — whether an alcohol can be oxidized, how it dehydrates, and which substitution pathway it favors all hinge on this single 1/2/3 label.
Classify by counting carbons on the C that holds the -OH: CH3-CH2-OH ethanol 1 (one C on the C-OH) (CH3)2CH-OH propan-2-ol 2 (two C's) (CH3)3C-OH 2-methylpropan-2-ol 3 (three C's) methanol CH3-OH is the special case: zero C's, also counted 1
The Acidity of Alcohols and the Alkoxide Ion
An alcohol's O-H proton is gently acidic — the polar O-H bond can let go of its hydrogen, just barely. When it does, it leaves behind an alkoxide ion, R-O- : an oxygen now carrying a lone pair and a full negative charge. The acidity of an alcohol sits near pKa about 16 to 18, almost identical to water (15.7). That is much weaker than a carboxylic acid (pKa ~4 to 5) and much stronger than an alkane (pKa ~50), which makes alcohols a kind of acidity benchmark you can anchor on. Using the master rule from the acid-base rung, this tells you exactly which bases can make an alkoxide: hydroxide barely manages it (near a tie), so chemists reach for a stronger base like sodium hydride or sodium metal, which drives the deprotonation to completion and fizzes off H2 gas.
Why is the proton even slightly willing to leave? Because oxygen is electronegative enough to hold the resulting negative charge fairly comfortably — far better than carbon could. But notice what is missing: the alkoxide's charge is stuck on one lone oxygen, with no resonance and no second electronegative atom to share the load. That is precisely the difference, traced back to the acidity reasoning of the last rung, between an alcohol and a carboxylic acid: a carboxylate spreads its charge over two oxygens, so it is much more stable and its parent acid is millions of times stronger. Alkyl groups make things slightly worse still, gently pushing electron density toward the oxygen and destabilizing the anion, which is why in solution a bulky tert-butoxide is a touch less stabilized than a small methoxide.
Phenols: When -OH Sits on a Ring
Move that same -OH from a saturated carbon onto a benzene ring and you no longer have an alcohol at all — you have a phenol, and it behaves like a different class of compound. The change in name reflects a real change in chemistry, driven entirely by the ring. The hallmark is acidity: a typical phenol sits near pKa 10, which makes it roughly a million times more acidic than an alcohol (pKa ~16). A phenol is acidic enough that ordinary hydroxide — which struggled to deprotonate an alcohol — now cleanly converts it to its conjugate base, the phenoxide ion. That single fact lets chemists separate phenols from neutral compounds with a simple base wash.
Where does the extra acidity come from? Resonance — the same theme that explained the carboxylate. When a phenol gives up its proton, the negative charge on the resulting phenoxide does not stay trapped on the oxygen; it spreads into the aromatic ring through pi delocalization, reaching onto the carbons at the ortho and para positions. Spreading charge over many atoms is stabilizing, and a more stable conjugate base always means a stronger acid. An alkoxide has no such escape route, so its charge stays cooped up on one oxygen — and that lone difference is the whole million-fold gap.