Molecular Structure of Acids and Bases: Study Guide (Chemistry 8.6)

The one question to ask: how stable is the conjugate base?

When an acid HA gives away a proton, it leaves behind its conjugate base A⁻, carrying the bonding pair of electrons and a negative charge:

HA(aq) + H₂O(l) ⇌ A⁻(aq) + H₃O⁺(aq)

Whether this equilibrium lies far to the right (strong acid) or to the left (weak acid) depends mostly on how comfortable A⁻ is with that extra charge. If A⁻ can hold or spread the charge well, it has little tendency to take the proton back. Then HA is a strong acid and A⁻ is a very weak base. If A⁻ holds the charge poorly, it grabs protons back, HA is weak, and A⁻ is a noticeably basic ion.

So almost every question in this topic comes down to one comparison: which conjugate base is more stable? Three structural features stabilise a conjugate base:

1. Electronegativity of the atom carrying the charge, or of atoms near it.
2. Inductive effects: electronegative atoms elsewhere in the molecule pulling electron density through the bonds.
3. Resonance: the charge spread over two or more atoms.

Which hydrogens are acidic?

A molecule can contain many hydrogen atoms but give away only some of them. The acidic hydrogens are usually bonded to a highly electronegative atom, most often oxygen or a halogen. The O–H or X–H bond is polar, and after the proton leaves, the electronegative atom keeps the pair of electrons and copes with the negative charge.

A hydrogen bonded to carbon is not acidic in water. Carbon and hydrogen have similar electronegativities, the C–H bond is barely polar, and a negative charge on carbon is very unstable.

Look at acetic acid, CH₃COOH. It has four hydrogen atoms, but only one is acidic: the hydrogen on the O–H group. The three hydrogens on the CH₃ group stay put. So acetic acid is monoprotic, and its conjugate base is the acetate ion, CH₃COO⁻.

This is the skill the exam asks for: look at the structure, find the H atoms on O (or on a halogen), and that tells you how many protons can take part in acid-base reactions.

The common strong acids and strong bases

You should know these by name; they are ionised completely in water.

Strong acids	Why the conjugate base is so weak
HCl, HBr, HI	Cl⁻, Br⁻ and I⁻ are large, stable ions with no tendency to take a proton back in water
HNO₃	NO₃⁻ spreads its charge over three O atoms by resonance
H₂SO₄ (first proton)	HSO₄⁻ spreads its charge over several O atoms, helped by the electronegative S and O atoms
HClO₄	ClO₄⁻ spreads its charge over four O atoms, with a strong inductive pull from the Cl

Strong bases are the hydroxides of group 1 (LiOH, NaOH, KOH, …) and the heavier group 2 metals (Ca(OH)₂, Sr(OH)₂, Ba(OH)₂). The base is the OH⁻ ion; its conjugate acid is water, an extremely weak acid. The metal ion, such as Na⁺, has no useful acidic or basic behaviour.

Everything else you meet is usually weak. Two families are worth recognising on sight:

• Weak acids: carboxylic acids, R–COOH (acetic acid, CH₃COOH, is the standard example).
• Weak bases: nitrogen bases with a lone pair, such as ammonia, NH₃, and amines like methylamine, CH₃NH₂; and carboxylate ions, R–COO⁻, the conjugate bases of carboxylic acids.

Electronegativity

Compare hydrogen compounds of elements in the same period. As the atom bonded to H becomes more electronegative, it holds the negative charge in the conjugate base better, so acid strength rises:

CH₄ < NH₃ < H₂O < HF (acid strength increases →)

CH₄ shows no acidic behaviour in water; HF is a weak acid.

Electronegativity also works one step away. In oxoacids (acids with the acidic H on an O bonded to a central atom), a more electronegative central atom pulls electron density from the O–H bond and stabilises the anion:

HOI < HOBr < HOCl (acid strength increases →)

Chlorine is the most electronegative of the three central atoms, so HOCl is the strongest of these three weak acids.

A warning about group trends. Down group 17 the order is HF < HCl < HBr < HI, the opposite of what electronegativity alone predicts. When the charged atom itself changes down a group, the size of the ion dominates: I⁻ is large, spreads its charge over a big volume, and the H–I bond is long and weak. Use electronegativity across a period or for atoms near the acidic site; use size when comparing the halide ions themselves.

Inductive effects

Electronegative atoms do not have to be bonded to the acidic H to matter. A chlorine atom two or three bonds away still pulls electron density through the chain of bonds towards itself. This inductive effect spreads out the negative charge of the conjugate base and stabilises it.

Chlorine-substituted acetic acids show this clearly:

CH₃COOH < ClCH₂COOH < Cl₂CHCOOH < Cl₃CCOOH (acid strength increases →)

Each extra Cl pulls more electron density away from the carboxylate group, so the conjugate base is more stable and the acid gives up its proton more readily. The effect weakens quickly with distance: a Cl atom far from the COOH group has much less effect than one on the neighbouring carbon.

The number of oxygen atoms on a central atom works the same way. In the chlorine oxoacids, each extra O (bonded only to Cl) pulls electron density away from the O–H group and gives the anion more atoms to share its charge:

HClO < HClO₂ < HClO₃ < HClO₄ (acid strength increases →)

Resonance

Compare acetic acid with ethanol, CH₃CH₂OH. Both have an O–H group, but acetic acid is a weak acid and ethanol is so much weaker that it does not noticeably change the pH of water. The difference is in the conjugate bases (Figure 1).

In acetate, the negative charge is delocalised: each oxygen carries about half of it, and the two C–O bonds are identical in length. A spread-out charge is more stable than a concentrated one. Ethoxide cannot spread its charge, so it is a high-energy, strongly basic ion, and ethanol is a very weak acid.

Resonance also explains the strong oxoacids in the table above: NO₃⁻ shares its charge over three O atoms and ClO₄⁻ over four.

Worked example 1: ranking oxoacids

Question. Rank HOBr, HOCl and HClO₃ from weakest to strongest acid. Justify each step using the structure of the conjugate base.

1. Find the acidic H. In all three, the acidic H is on an oxygen bonded to the halogen. The conjugate bases are OBr⁻, OCl⁻ and ClO₃⁻.
2. HOBr vs HOCl. Same structure, different central atom. Cl is more electronegative than Br, so it pulls electron density away from the O and stabilises OCl⁻ more than Br stabilises OBr⁻. HOCl is stronger.
3. HOCl vs HClO₃. Same central atom, but HClO₃ has two extra O atoms on the Cl. They pull electron density from the O–H group (inductive effect), and in ClO₃⁻ the negative charge is spread over three O atoms by resonance. OCl⁻ has its charge on one O. ClO₃⁻ is more stable, so HClO₃ is stronger.

Answer. HOBr < HOCl < HClO₃.

How to write it on an exam. Name the feature (electronegativity of the central atom; number of O atoms), say what it does to the conjugate base (stabilises the negative charge), then link that to acid strength (a more stable conjugate base means a stronger acid).

Worked example 2: acids and bases from structure and data

Question. (a) How many acidic hydrogens does chloroacetic acid, ClCH₂COOH, have? (b) Is ClCH₂COOH a stronger or a weaker acid than acetic acid? (c) Acetic acid has Ka = 1.8 × 10⁻⁵ and the ammonium ion, NH₄⁺, has Ka = 5.6 × 10⁻¹⁰. Which is the stronger base, the acetate ion or ammonia?

(a) Three hydrogens in the formula, but two are on carbon. Only the H on the O–H group is acidic: one acidic hydrogen (monoprotic).

(b) Stronger. The Cl atom on the carbon next to the carboxyl group pulls electron density towards itself (inductive effect). This spreads out the negative charge in ClCH₂COO⁻, which is therefore more stable than CH₃COO⁻, so ClCH₂COOH gives up its proton more readily.

(c) For each conjugate pair, Kb = Kw ÷ Ka.

• Acetate: Kb = 1.0 × 10⁻¹⁴ ÷ 1.8 × 10⁻⁵ = 5.6 × 10⁻¹⁰
• Ammonia: Kb = 1.0 × 10⁻¹⁴ ÷ 5.6 × 10⁻¹⁰ = 1.8 × 10⁻⁵

Ammonia is the stronger base, by a factor of about 3 × 10⁴. This fits the structural picture: acetate’s charge is spread by resonance, so it has little drive to accept a proton. Ammonia’s lone pair sits on a single nitrogen atom, ready to accept one. Both are still weak bases; neither comes close to OH⁻.

Check. The weaker acid of each pair (NH₄⁺, Ka = 5.6 × 10⁻¹⁰) has the stronger conjugate base (NH₃). Ka × Kb = Kw makes this pattern automatic.

Common misconceptions

• “Every hydrogen in an acid’s formula is acidic.” Only hydrogens on electronegative atoms (usually O or a halogen) are. Acetic acid has four H atoms and one acidic proton.
• “Strong acid means concentrated acid.” Strength is about the extent of ionisation, which is fixed by structure. A dilute HCl solution is still a strong acid.
• “The more electronegative the halogen, the stronger the hydrohalic acid.” Down group 17 the order is reversed: HF is weak, HI is the strongest. Ion size and bond strength win there.
• “A strong acid has a strong conjugate base.” The opposite: the stronger the acid, the weaker its conjugate base. Cl⁻ has essentially no basic behaviour in water.
• “Resonance makes the acid molecule more stable, so the acid is weaker.” What matters is the resonance in the conjugate base. Charge spread over several atoms stabilises A⁻ and makes HA stronger.
• “Ammonia is a base because it contains OH⁻.” NH₃ contains no hydroxide. It is a base because its nitrogen lone pair accepts a proton from water, which produces OH⁻.
• Explaining strength only with bond polarity. Polarity helps identify acidic hydrogens, but full explanations should talk about the stability of the conjugate base.

Where this leads

Next, Topic 8.7, pH and pKa, uses pKa values, which you can now predict from structure, to decide whether the acid or base form of a substance dominates at a given pH, and to choose indicators. The strength differences explained here are what produce the different curve shapes in Acid-Base Titrations. Try the practice questions now, then use the revision notes and the checklist. For resonance structures, revisit Resonance and Formal Charge.