Carboxylic acids — Organic Chemistry Notes

Acetic acid, CH₃COOH, and ethanol, CH₃CH₂OH, both carry an O–H on a two-carbon frame. Acetic acid gives up that proton about a hundred billion times more readily. Its C=O is also a poorer target for a nucleophile than the C=O of acetone.

Both differences come from one interaction between the C=O and the OH on the same carbon. This page shows that interaction, then follows it through acidity, reactivity, and the one common way these acids fall apart on heating.

Structure

In acetic acid, one carbon is double-bonded to an oxygen and single-bonded to an OH and a CH₃. That C=O plus OH unit is the carboxyl group, COOH, and its carbon is the carboxyl carbon. A carboxylic acid is any compound with a carboxyl group, written RCOOH.

The carboxyl carbon is bonded to three atoms, one of them by a double bond, so it is sp². Its three groups lie flat, about 120° apart: the trigonal planar shape from Molecular geometry.

The OH oxygen has a lone pair right next to the C=O pi bond. That pair can move into the C–O bond while the C=O pi electrons move onto the other oxygen. The result is a second resonance contributor, with a positive charge on the OH oxygen and a negative charge on the carbonyl oxygen.

Follow the two arrows on the left: a lone pair on the OH oxygen moves into the C–O bond, and the C=O pi electrons move onto the top oxygen. That gives the structure on the right.

Because it separates a + and a − charge, this contributor counts for less than the uncharged one, but it still counts. The C–OH bond gains some double-bond character, and it is measurably shorter than the C–O bond of an alcohol (143 pm). This donation happens in the neutral acid, before any proton leaves.

Acetic acid, at 60 g/mol, boils at 118 °C. Acetone, at 58 g/mol, boils at 56 °C. The acid boils higher because its molecules hydrogen-bond to one another, and acetone, with no O–H, cannot. Two acid molecules can pair up into a dimer, held by two hydrogen bonds at once: each molecule’s O–H hydrogen reaches the other molecule’s C=O oxygen.

Each dashed line runs from one molecule’s O–H hydrogen to the other molecule’s C=O oxygen. Count the ring they close: eight atoms.

The dimer is strong enough to survive boiling. Much of acetic acid vapor is made of dimers of mass 120, so a measurement of the vapor’s molecular weight just above the boiling point comes out well above 60. That was early evidence that the dimer is a real species.

Why the O–H is so acidic

Ethanol has a pKa of 16. Acetic acid has a pKa of 4.76. Both lose a proton from an O–H, so compare what each leaves behind. Ethanol leaves ethoxide, CH₃CH₂O⁻, which keeps the whole negative charge on one oxygen. Acetic acid leaves acetate, CH₃COO⁻. The general name for an anion RCOO⁻ is a carboxylate.

On the left the charge has one place to be. On the right, the arrows move it from the lower oxygen to the upper one, and the two drawings differ only in which oxygen is which.

Acetate has two resonance contributors, and they are equivalent: each puts the charge on a different oxygen. Equivalent contributors give the most stabilization, because the charge is then shared equally. A charge shared by two oxygens is far more stable than one held by a single atom. The gap between ethanol and acetic acid, about eleven pKa units, is the hundred-billion factor from the start of the page. Carboxylic acids in general sit at pKa 4 to 5.

Measurements confirm the sharing. In a carboxylate, both C–O bonds are 126 pm long, between a C–O single bond (143 pm) and a C=O double bond (123 pm). Each oxygen carries half the negative charge.

The two contributors are not two molecules. Acetate drawn with the double bond to one oxygen, then to the other, is one ion drawn two ways. Nothing flips back and forth. The real ion has two identical C–O bonds and half a charge on each oxygen, so it matches neither drawing on its own.

Factors affecting acidity compares two acids with its ARIO checklist: the Atom that holds the charge, Resonance, Induction (a pull on electrons through the sigma bonds) and Orbital. Ethanol and acetic acid tie on atom, since oxygen holds the charge in both. Most courses credit the whole eleven-unit gap to resonance, and exams expect that answer. Some analyses of the numbers give a large share to the C=O group itself. Its carbon, already partly positive, sits right next to the oxygen that takes the charge, and it steadies that charge by induction and electrostatic attraction. On that view, a good part of the gap would remain even without delocalization. How to split the gap between the two effects is still argued.

Substituent effects on top of resonance

An electronegative atom near the carboxyl group pulls electron density away through the sigma bonds. That pull is the induction of ARIO, and it steadies the carboxylate further. Acetic acid has a pKa of 4.76. Chloroacetic acid, ClCH₂COOH, has 2.86; dichloroacetic acid, 1.29; trichloroacetic acid, 0.65. Each chlorine adds to the effect.

Induction fades quickly with distance. Put the chlorine three carbons away, in 4-chlorobutanoic acid, and the pKa is 4.52, barely below butanoic acid’s 4.82. Factors affecting acidity drew this fall-off one bond at a time.

Oxalic acid, HOOC–COOH, has two carboxyl groups. A compound like this is a diacid, and it loses its two protons one after the other. So it has two pKa values, and they always differ: 1.27 for the first proton and 4.27 for the second.

Read the bottom line of each panel. The gap between the two pKa values closes as the marked carbons push the two carboxyl groups apart.

The first proton leaves more easily than acetic acid’s does. The other COOH group is electron-withdrawing, so it steadies the new carboxylate by induction.

The second proton leaves less easily than the first. The molecule already carries one negative charge, and removing a second proton puts a second negative charge close to it. Like charges repel. For malonic acid, HOOC–CH₂–COOH, and succinic acid, HOOC–CH₂CH₂–COOH, the second proton is even harder to remove than acetic acid’s. Oxalic acid is the exception: its second pKa, 4.27, is still below acetic acid’s 4.76.

Both effects weaken as the two groups move apart, so the gap between the two pKa values shrinks. It is 3.0 units for oxalic acid, whose carboxyl carbons are bonded to each other, and 1.4 units for succinic acid, with two carbons between them.

A benzene ring passes these effects along too. Benzoic acid, a COOH on a benzene ring, has a pKa of 4.20. Put a nitro group, NO₂, directly across the ring from the COOH, at the para position. The nitro group pulls electrons out of the ring, and the pKa drops to 3.44. A methoxy group, OCH₃, in the same place pushes electrons into the ring, and the pKa rises to 4.47. Aromatic Chemistry explains why each group pushes or pulls, and uses the same effects to decide how benzene rings react; here they show up as a change in pKa.

The group sits across the ring from the COOH, the para position. Compare each pKa with benzoic acid’s 4.20.

Practical acidity: the separation you will actually run

Carboxylic acids (pKa 4 to 5) and phenols (pKa 10) can be told apart with sodium bicarbonate, NaHCO₃. pKa showed how to read an equilibrium from two pKa values; the example below runs it.

Worked example — separating benzoic acid from phenol

A mixture of benzoic acid (pKa 4.20) and phenol (pKa 10.0) is dissolved in ether. Separate them.

Step 1 — pick a base whose conjugate acid’s pKa falls between the two. Bicarbonate works, since carbonic acid’s pKa is 6.4. Benzoic acid is 2.2 units below that, so its equilibrium lies about 160 to 1 toward the carboxylate. Phenol is 3.6 units above it, so almost none of the phenol reacts.

Step 2 — shake the ether with aqueous NaHCO₃. Sodium benzoate is an ionic salt, so it moves into the water layer. The carbonic acid formed alongside it breaks down to water and CO₂, and you see bubbles. Phenol stays neutral and stays in the ether.

Step 3 — separate the layers. The ether layer holds the phenol. The water layer holds the sodium benzoate.

Step 4 — get the acid back. Add HCl to the water layer. Benzoic acid becomes neutral again, does not dissolve well in water, and comes out as a solid. Filter it off.

Sodium hydroxide would have failed. Its conjugate acid is water, pKa 15.7, so hydroxide deprotonates both compounds, and both end up in the water layer.

The same reasoning applies in the body. Blood has a pH of 7.4, about three units above a typical carboxylic acid’s pKa. So a carboxylic acid drug in the blood is almost entirely in its carboxylate form. That changes how well it dissolves, how easily it crosses cell membranes, and where it ends up in the body.

Worked example — predicting relative acidity

Rank acetic acid, trifluoroacetic acid (CF₃COOH), phenol and ethanol from most to least acidic.

Ethanol (pKa 16) is the weakest. Its conjugate base has no resonance. Phenol (10) comes next: its anion spreads the charge into the ring, but onto carbons, which hold negative charge poorly. Acetic acid (4.76) spreads its charge over two oxygens. Trifluoroacetic acid (0.23) does the same, and its three fluorines add strong induction.

Order: CF₃COOH > CH₃COOH > phenol > ethanol. Resonance sorts the four into tiers. Induction then adjusts within a tier.

The carbonyl carbon: less electrophilic than a ketone

Acetone’s carbonyl carbon is bonded to two CH₃ groups. They donate only a little, through the sigma bonds, as Aldehydes & ketones showed. Acetic acid’s carbonyl carbon is bonded to a CH₃ and an OH, and the OH lone pair feeds electron density onto the carbonyl carbon, as the first figure on this page showed. That cancels part of the carbon’s partial positive charge. A carboxylic acid is therefore noticeably less reactive toward nucleophilic attack than a ketone.

Compare the two δ+ labels. Only the acid has a lone pair next door to feed its carbonyl carbon.

There is a second obstacle: the acidic proton. A strongly basic nucleophile removes that proton before it can attack the carbon. An amine, for example, simply forms a salt with the acid. The carboxylate left behind carries a negative charge, so it attracts a nucleophile even less than the acid did.

LiAlH₄ also removes the proton first, and H₂ bubbles off. It is one of the few reagents reactive enough to go on and attack the carboxylate. Most nucleophiles stop at the salt. So a chemist usually turns a carboxylic acid into a more reactive compound, such as an acid chloride, before working at its carbonyl.

A Grignard reagent cannot add to a carboxylic acid’s C=O. Nucleophilic addition showed that acidic protons destroy Grignard reagents. The acid’s O–H does this in a fast acid–base reaction, so the Grignard reagent never reaches the carbonyl. Making a ketone from an acid needs a detour through a more reactive compound such as an acid chloride, covered in Acid chlorides & anhydrides.

Getting to and from carboxylic acids

Butanoic acid can be made from at least four different starting materials. Oxidize butan-1-ol or butanal. Turn 1-bromopropane into a Grignard reagent and let it attack CO₂. CO₂ has a C=O, so the Grignard carbon attacks it just as it attacks a ketone in Nucleophilic addition, and the chain gains one carbon. Or cut oct-4-ene in two at its C=C.

The marked carbon in each starting material ends up as the COOH carbon. From 1-bromopropane that carbon comes from CO₂; from oct-4-ene each alkene carbon becomes one.

The table gives the general version of each route. Each one has its own conditions, and the last column says what to watch for.

To make RCOOH from	Reagent	Watch for
1° alcohol RCH₂OH	KMnO₄, or CrO₃/H₂SO₄ (Jones)	PCC stops at the aldehyde instead
aldehyde RCHO	KMnO₄, Ag₂O, or Jones	aldehydes oxidize very easily
nitrile RCN	H₃O⁺ and heat; or HO⁻ and heat, then H₃O⁺	preview: goes through the amide, no oxidant needed; covered in Nitriles
alkyl halide RX	1. Mg 2. CO₂ 3. H₃O⁺	adds one carbon; see Grignard reagents
alkene	hot KMnO₄, or O₃ then H₂O₂	cleaves the C=C, so count carbons

From the acid, SOCl₂ gives the acid chloride, RCOCl. This is the standard first step, because the acid chloride is the most reactive of the carboxylic acid derivatives, the compounds made by replacing the acid’s OH with another group. It can be turned into any of the others, as Nucleophilic acyl substitution will show.

Two routes skip the acid chloride. Heating the acid with an alcohol and a little strong acid gives an ester, a reaction called Fischer esterification (Esters & amides). And LiAlH₄ reduces the acid all the way to a primary alcohol; NaBH₄ does not.

How β-keto acids fall apart

Butanoic acid boils at 164 °C and distills unchanged. 3-Oxobutanoic acid, CH₃COCH₂COOH, falls apart on gentle warming, around 50 to 100 °C. It loses CO₂ and leaves acetone. Losing CO₂ this way is called decarboxylation.

The difference is a ketone in the right place. Label the carbons outward from the carboxyl carbon: the next one is the α (alpha) carbon, and the one after that is the β (beta) carbon. In 3-oxobutanoic acid, the β carbon is a ketone carbon. An acid built this way, RCOCH₂COOH, is a β-keto acid.

The molecule can curl so that its carboxyl O–H reaches the ketone oxygen. Six atoms then form a ring: the ketone oxygen, the ketone carbon, the α carbon, the carboxyl carbon, the OH oxygen and the moving hydrogen. Three pairs of electrons move around that ring at once. The ketone oxygen takes the hydrogen, using the electrons of its C=O pi bond. The electrons of the O–H bond become the second C=O of CO₂. The electrons of the bond from the α carbon to the carboxyl carbon become the pi bond of a new C=C, between the α carbon and the ketone carbon. All three shifts happen in one step, which is why gentle warming is enough.

That last arrow is why the ketone must sit at the β carbon. The electrons of the breaking bond need a pi bond to form, and the α carbon can form one only to a neighbor that is giving up its own pi bond: the ketone carbon. If the ketone carbon sits one carbon further out, the α carbon is not bonded to it, so the breaking bond’s electrons have no pi bond to form and no C=C appears. If the ketone carbon is the α carbon itself, it is the carbon losing the bond, and nothing next to it can take the electrons into a new C=C.

Follow the three arrows round the ring on the left. Then follow the highlighted CH₂ carbon: it ends up in the enol’s C=C, and then as a CH₃ of acetone.

The first product is not acetone. It is an enol, a C=C with an OH on one of its carbons, because the hydrogen went onto the ketone oxygen. The enol then tautomerizes to the ketone, as it did in Alkynes. The equilibrium lies far on the ketone side, so acetone is what you isolate. A mechanism drawn straight from the ring to acetone skips the enol.

Malonic acid (drawn in the diacid figure above) decarboxylates the same way, but needs stronger heating, near its melting point of about 135 °C. Its second carboxyl group plays the part of the ketone, and the product is acetic acid. A substituted malonic acid, RCH(COOH)₂, with two COOH groups on one carbon, does the same.

This reaction returns as the last step of the malonic and acetoacetic ester syntheses in Enolate Chemistry. For now, hold the narrow point: among ordinary acids, a carbonyl at the β position is what makes one lose CO₂ on gentle warming. A few other acids lose CO₂ on stronger heating, trichloroacetic acid and oxalic acid among them, but not through this ring.

What carries forward

The carboxylic acid is the parent of the derivatives in the next topics: esters, amides, acid chlorides and anhydrides. Each has an atom with a lone pair bonded to the carbonyl carbon, and how strongly that lone pair donates sets how reactive the derivative is. Turning an acid into its acid chloride before further reactions is a standard move in synthesis. And the carboxylate, with its two identical C–O bonds, is the clearest measured example of resonance in the course.