Difficulty: Intermediate | Prerequisites: Acid-Base Fundamentals (Part 1 of these notes), periodic table trends
Four main factors determine how acidic or basic a molecule is: charge stability on the conjugate base, electronegativity of the atom bearing the charge, resonance delocalisation, and hybridisation (s-character). When comparing across a row of the periodic table, electronegativity wins; when comparing down a column, size wins.
Charge stability
The degree to which a charged species (usually the conjugate base) can tolerate its charge without being highly reactive. More stable conjugate bases come from stronger acids.
Electronegativity
The tendency of an atom to attract electrons toward itself. More electronegative atoms stabilise negative charges better, making the corresponding acid stronger.
Resonance stabilisation (delocalisation)
The spreading of a charge over multiple atoms through overlapping p-orbitals. More resonance structures that spread the charge out means greater stability.
Inductive effect
The pulling of electron density through sigma bonds by electronegative atoms. Electron-withdrawing groups near an acidic proton increase acidity by stabilising the conjugate base.
Hybridisation
The mixing of atomic orbitals. sp-hybridised carbons have 50% s-character and hold electrons more tightly than sp³ carbons (25% s-character), making sp C-H bonds more acidic.
s-character
The proportion of s-orbital character in a hybrid orbital. Higher s-character means electrons are held closer to the nucleus, stabilising negative charge.
Protonation state
The form of a molecule at a given pH. At pH < pKa, the protonated (acid) form dominates. At pH > pKa, the deprotonated (base) form dominates. At pH = pKa, both forms exist in equal amounts.
The central question for predicting acidity: how stable is the conjugate base?
If the conjugate base (A⁻) is more stable, the acid (AH) is stronger.
If the conjugate base is less stable (more reactive), the acid is weaker.
There is always an energy barrier that must be crossed for the proton to leave. The height of that barrier relates to how well the resulting anion can handle the negative charge.
Anything that withdraws electron density from the conjugate base increases acidity.
Electronegativity is the most direct factor: more electronegative atoms stabilise negative charges better.
Across a row of the periodic table, electronegativity increases left to right, and so does acid strength:
C-H (pKa ~50) < N-H (pKa ~35) < O-H (pKa ~16) < F-H (pKa ~5)
Having more electronegative atoms elsewhere in the molecule also matters (inductive effect). Electronegative substituents pull electron density away from the conjugate base, stabilising it.
Example: F₃C-COOH is more acidic than CH₃-COOH because the three fluorines withdraw electron density.
Resonance structures spread the negative charge across multiple atoms.
More delocalisation means more stability, which means a stronger acid.
Example of ranking conjugate base stability:
A carboxylate ion (RCO₂⁻) has the charge spread over two oxygen atoms (resonance), making carboxylic acids (pKa ~5) much more acidic than alcohols (pKa ~16), where the charge sits on a single oxygen.
Ranking by stability: charge spread over two oxygens (most stable) > charge on one electronegative atom > charge on carbon (least stable).
sp hybridised orbitals have 50% s-character.
sp² hybridised orbitals have 33% s-character.
sp³ hybridised orbitals have 25% s-character.
Electrons in orbitals with more s-character are held closer to the nucleus and are more stable.
Therefore: sp C-H bonds are more acidic than sp² C-H bonds, which are more acidic than sp³ C-H bonds.
When comparing atoms in the same column, size wins over electronegativity.
Larger atoms spread the charge over a bigger volume, stabilising the conjugate base.
Example: HI > HBr > HCl > HF in acid strength, even though F is the most electronegative. The iodide ion is so large it easily accommodates the negative charge.
Rule of thumb: across a row, electronegativity wins. Down a column, size wins.
To rank acids: identify each conjugate base, assess its stability using the four factors above, and rank accordingly. The acid whose conjugate base is most stable is the strongest acid.
To rank bases: the strongest base has the least stable conjugate acid (highest pKa for its conjugate acid).
Example ranking from the lecture (carboxylate species):
An anion with charge spread over two oxygens by resonance (pKa ~5, most stable) > a phenoxide-type anion (pKa ~10) > an alkoxide (pKa ~16) > a carbanion (pKa ~40-50, least stable)
At pH = pKa, both the protonated and deprotonated forms exist in equal amounts ("perfect balance").
At pH < pKa (more acidic environment), the protonated (acid) form dominates.
At pH > pKa (more basic environment), the deprotonated (base) form dominates.
Reference points:
pH = 1: strongly acidic, both acid forms present
pH = 7: neutral
pH = 14: strongly basic, both base forms present
The acidity trends covered here are exactly why aspirin (acetylsalicylic acid, pKa ~3.5) is absorbed in the acidic environment of the stomach (pH ~2), where it remains mostly in its protonated, uncharged form and can cross cell membranes. Once it reaches the more basic environment of the bloodstream (pH ~7.4), it deprotonates and becomes trapped inside cells.
Buffer systems in biology rely on pH-pKa relationships. The bicarbonate buffer in your blood maintains pH near 7.4 because the pKa of carbonic acid (~6.4) sits close enough to buffer effectively. Understanding these equilibria is essential in pharmacology, biochemistry, and environmental chemistry.
"Electronegativity always determines acid strength." This is only true across a row. Down a column, size (polarisability) dominates. HF is the weakest hydrohalic acid despite fluorine being the most electronegative halogen.
"More resonance structures always means more stable." Quality matters. A resonance structure that places a negative charge on a more electronegative atom contributes more to stability than one that places it on carbon.
"A molecule with a lower pKa is always a stronger base." Lower pKa means stronger acid, not stronger base. To compare base strength, look at the pKa of the conjugate acid: the base whose conjugate acid has the higher pKa is the stronger base.
"At pH = 7, everything is deprotonated." Whether a molecule is protonated at pH 7 depends on its pKa. A carboxylic acid (pKa ~5) is mostly deprotonated at pH 7, but an amine (pKa ~10 for conjugate acid) is mostly protonated at pH 7.
⚠️ Ranking questions are extremely common. Expect to be given 3-4 molecules and asked to rank them by acidity or basicity, using the four factors.
⚠️ Know the hierarchy: across a row, electronegativity wins. Down a column, size wins. Within similar structures, resonance and induction break ties.
⚠️ Be comfortable moving between "which is more acidic?" and "which conjugate base is more stable?" These are the same question asked from different angles.
⚠️ Protonation-state questions at a given pH will appear. Memorise: pH < pKa means the acid form dominates; pH > pKa means the base form dominates.
⚠️ Hybridisation effects on acidity (sp > sp² > sp³) are commonly tested and often overlooked by students.
True or False: HI is a stronger acid than HF. (True, size wins down a column.)
Fill in the blank: Across a row of the periodic table, ______ determines relative acid strength. (electronegativity)
True or False: A carboxylate ion is less stable than an alkoxide ion. (False, the carboxylate has resonance stabilisation across two oxygens.)
Fill in the blank: At pH = pKa, ______% of the molecule is in the protonated form. (50)
True or False: An sp³ C-H bond is more acidic than an sp C-H bond. (False, sp has more s-character and is more acidic.)
Q: Rank the following in order of increasing acidity: CH₄, NH₃, H₂O, HF.
A: CH₄ < NH₃ < H₂O < HF. Across the second row, electronegativity increases from C to F, stabilising the conjugate base and making each successive hydride a stronger acid.
Q: Why is acetic acid (CH₃COOH, pKa ~5) so much more acidic than ethanol (CH₃CH₂OH, pKa ~16)?
A: The conjugate base of acetic acid (acetate, CH₃COO⁻) is stabilised by resonance: the negative charge is delocalised over two equivalent oxygen atoms. The ethoxide ion (CH₃CH₂O⁻) has no such resonance, so the charge is localised on one oxygen.
Q: Between an SH group and an NH group, which is more acidic, and why?
A: SH is more acidic. Sulphur is larger than nitrogen (further down the periodic table), so the conjugate base (RS⁻) spreads its negative charge over a larger volume. This is a case where size wins, as S and N are in different rows.
Q: A compound has pKa = 4.5. At pH 7.4 (physiological pH), what is the dominant form?
A: The deprotonated (base/conjugate base) form dominates. Since pH (7.4) > pKa (4.5), the equilibrium favours loss of the proton.
Q: Rank these carbanions by stability: CH₃⁻ (from sp³ C-H), vinyl carbanion (from sp² C-H), acetylide anion (from sp C-H).
A: Acetylide (sp, most stable) > vinyl (sp²) > methyl (sp³, least stable). More s-character means the electrons are held closer to the carbon nucleus, stabilising the anion.
Conjugate-base stability directly predicts leaving-group ability in substitution and elimination reactions (Chapters 7-8). Good leaving groups are the conjugate bases of strong acids.
The factors covered here (electronegativity, resonance, induction) reappear when predicting the reactivity of carbonyl compounds and the stability of enolates in later chapters.
Understanding protonation states at physiological pH is essential for amino acid chemistry (Chapter 24) and drug design in medicinal chemistry courses.
Acid strength factors, conjugate base stability, electronegativity and acidity, resonance stabilisation, inductive effect, electron-withdrawing group, hybridisation and acidity, s-character, atom size and acidity, periodic trends in acidity, protonation state, pH vs pKa, Henderson-Hasselbalch, ranking acidity, ranking basicity, carbanion stability, leaving-group ability, charge delocalisation