Organic Chemistry · Polar Covalent Bonds; Acids and Bases

Acid and Base Strength

9 min read

Science note: pK_a values cited (HCl −7, H3O+ −1.74, HF 3.2, acetic acid 4.74, chloroacetic acids 2.86/1.29/0.64, NH4+ 9.3, phenol 10, water 15.7, ethanol 16, alkyne 25, NH3 38, alkane 50) are standard textbook values; no experimental data were fabricated.

On this page 9 sections

In 30 seconds
Why this matters
The college version
Eli explains
Worked example
Key takeaway
Check yourself
Study tools
Sources & references

In 30 seconds

Acid strength is measured by how far the dissociation equilibrium lies toward products — quantitatively, by K_a or pK_a. This topic explains the structural factors that control pK_a: the of the atom bearing the proton, the strength of the H–X bond, the hybridization () of that atom, inductive effects from nearby substituents, resonance in the , and the charge of the species. Because every base has a conjugate acid, base strength is simply the mirror image of acid strength: strong bases have weak conjugate acids, and the pK_a of the conjugate acid is the quantitative handle. The same factors that stabilize a conjugate base make the parent acid stronger, so one framework — "whatever stabilizes the conjugate base" — predicts nearly every acidity trend in the chapter and the book.

Why this matters

Predicting which proton comes off a molecule, and which base can remove it, is the core skill of organic synthesis and mechanism. Choosing a base for a deprotonation (lithium diisopropylamide vs. sodium hydroxide vs. bicarbonate) is a pK_a decision; estimating whether a reaction's equilibrium lies left or right is a pK_a comparison. In biology and medicine, acid–base strength governs drug ionization and absorption, enzyme active-site protonation states, and buffer systems in blood. The same reasoning also explains why carboxylic acids are acidic but alcohols are not, and why phenols are more acidic than alcohols — two facts that appear constantly in later chapters on carbonyl chemistry and aromatic compounds.

The college version

Core Concepts

The pKa scale

The acid-dissociation constant for HA in water is K_a = [H₃O⁺][A^-]/[HA], and pK_a = -log₁₀ K_a. The pK_a scale runs from about −10 (superacids) to 50 (alkanes). Rule of thumb: a change of one pK_a unit is a tenfold change in K_a; a change of five units is a 10⁵-fold change. Representative values: HCl ≈ −7, H₃O⁺ ≈ −1.74, HF ≈ 3.2, CH₃COOH ≈ 4.74, NH₄⁺ ≈ 9.3, phenol ≈ 10, water ≈ 15.7, ethanol ≈ 16, terminal alkyne ≈ 25, ammonia ≈ 38, alkane ≈ 50.

Factor 1: Electronegativity of the atom holding the proton

Across a row, acidity increases with electronegativity because a more electronegative atom holds the negative charge of the conjugate base better: CH₄ (pK_a ≈ 50) < NH₃ (38) < H₂O (15.7) < HF (3.2).

Factor 2: Bond strength down a group

Down a column, electronegativity decreases but bond strength drops even faster; the weaker the H–X bond, the easier the proton leaves: HF (3.2) < HCl (−7) < HBr (−9) < HI (−10). Bond strength, not electronegativity, dominates this trend.

Factor 3: Hybridization and s-character

A carbon with more s-character in its hybrid orbitals is more electronegative: sp > sp² > sp³. Thus terminal alkynes (sp, pK_a ≈ 25) are much more acidic than alkenes (sp², ≈ 44) and alkanes (sp³, ≈ 50). This is why acetylide ions exist and are used to build carbon–carbon bonds, while alkyl anions are essentially unattainable under normal conditions.

Factor 4: Inductive effects

An electronegative substituent near the acidic proton withdraws electron density and stabilizes the conjugate base, raising acidity. Substituted acetic acids illustrate the effect: CH₃COOH pK_a ≈ 4.74, ClCH₂COOH ≈ 2.86, Cl₂CHCOOH ≈ 1.29, Cl₃CCOOH ≈ 0.64. Each added chlorine lowers the pK_a by roughly 1.5–2 units, and the effect fades with distance (β-chloro acids are weaker than α-chloro acids).

Factor 5: Resonance in the conjugate base

If the conjugate base can delocalize its negative charge, the acid is dramatically stronger. Carboxylic acids (pK_a ≈ 4.74) beat alcohols (pK_a ≈ 16) by about eleven orders of magnitude because the carboxylate anion spreads the charge over two equivalent oxygens, while an alkoxide holds it on one. Phenol (pK_a ≈ 10) beats ethanol (16) because phenoxide delocalizes the charge into the ring. Resonance is the strongest of the stabilizing effects for common organic acids.

Factor 6: Charge

Charged acids are much stronger than their neutral relatives: H₃O⁺ (−1.74) vs. H₂O (15.7); NH₄⁺ (9.3) vs. NH₃ (38). A positive charge destabilizes the acid form relative to the neutral conjugate base, pushing dissociation forward.

Base strength

Base strength is the inverse of the pK_a of the conjugate acid: a base whose conjugate acid has a high pK_a is a strong base. Ranking common bases: NH₂^- (conjugate acid pK_a ≈ 38) and CH₃CH₂O^- (16) are strong; OH^- (15.7) moderate-strong; CH₃COO^- (4.74) weak; Cl^- (−7) very weak. The same structural factors apply with the sign flipped: stabilize the conjugate acid, and the base gets stronger.

Common Confusions

Do Not Confuse	With	Difference
Strong acid	Concentrated acid	Strength is inherent (K_a); concentration is how much acid is in solution.
Large K_a	Large pK_a	Large K_a = strong; large pK_a = weak. They move in opposite directions.
Base strength	Nucleophilicity	Basicity is proton affinity (an equilibrium property); nucleophilicity is attack on carbon (a kinetic property) — discussed in later chapters.
Electronegativity trend	Bond-strength trend	Across a row electronegativity wins; down a group bond strength wins.
Resonance effect	Inductive effect	Resonance acts through π systems and does not fade with distance; induction acts through σ bonds and fades with distance.
Hydroxide as strongest base	Alkoxide/amide bases	Hydroxide's conjugate acid is water (15.7); alkoxides (16) and amide ion (38) are stronger bases.

Eli explains

The same idea, in plain words

Explain it like I’m 10

Acid strength is like how easily a toy lets go of a marble. An acid that holds its proton loosely is strong, and one that grips it tightly is weak. What matters most is how comfortable the empty hand (the conjugate base) is after the marble leaves: the more ways it can spread the leftover charge — over hungrier atoms, through nearby helpers, or across a whole chain of atoms — the happier it is, and the easier the acid lets go.

Worked examples

Compare acetic acid, CH₃COOH, pK_a ≈ 4.74, with ethanol, CH₃CH₂OH, pK_a ≈ 16. Both donate a proton from an O–H bond, so electronegativity and bond strength are equal. The difference is in the conjugate bases. Ethoxide, CH₃CH₂O^-, holds the entire negative charge on one oxygen. Acetate, CH₃COO^-, spreads the same charge over two equivalent oxygens through resonance. Because charge spread over two atoms is far more stable than charge confined to one, acetate is dramatically more stable, and acetic acid is about 10¹¹ times more acidic:

$K a (CH 3 COOH) K a (CH 3 CH 2 OH) = 10 (pK a (EtOH) - pK a (AcOH)) = 10 (16 - 4.74) = 10 11.26 \approx 1.8 \times 1011$

Rule of thumb: a pK_a gap of 11 units is an 11-order-of-magnitude gap in the equilibrium constant.

Rank the hydrohalic acids by strength using bond energy. The H–X bond dissociation energies fall down the group (H–F is strongest, H–I weakest), so acidity rises: HI (−10) > HBr (−9) > HCl (−7) > HF (3.2). Now predict whether Cl^- can deprotonate acetic acid. Cl^- is the conjugate base of HCl, whose pK_a is −7, so chloride is a very weak base. Acetic acid has pK_a 4.74. The equilibrium

$CH 3 COOH + Cl - ⇌ CH 3 COO - + HCl$

compares two acids: acetic acid (4.74) is a far weaker acid than HCl (−7), so the reaction does not proceed — chloride cannot deprotonate acetic acid. General rule: the base whose conjugate acid has the higher pK_a wins, and equilibrium favors the weaker acid and weaker base.

The pK_a of acetic acid is 4.74 at 25 °C. Convert to K_a:

$K a = 10 -pK a = 10 -4.74 = 1.8 \times 10 -5$

The value 1.8 × 10^-5 carries units of molarity (mol/L), matching the definition K_a = [H₃O⁺][A^-]/[HA], where the numerator units M × M divide by the denominator M to give M. Now estimate the pH of a 0.10 M acetic acid solution. With x = [H₃O⁺], the equilibrium expression is:

$K a = x 2 0.10 - x \approx x 2 0.10$

Solving:

$x = \sqrt K a \times 0.10 M = \sqrt 1.8 \times 10 -5 \times 0.10 M 2 = \sqrt 1.8 \times 10 -6 M 2 = 1.3 \times 10 -3 M$

Dimensional check: K_a (M) × concentration (M) = M², and the square root returns M. Then:

$pH = - log 10 (1.3 \times 10 -3) = 2.87$

A 0.10 M solution of a weak acid has pH near 3, far above the pH ≈ 1.0 that a 0.10 M strong acid would give — a direct demonstration of what "weak acid" means.

Key takeaways

Smaller pK_a = stronger acid; one pK_a unit = tenfold change in K_a.
Across a row: acidity rises with electronegativity (CH4 < NH3 < H2O < HF).
Down a group: acidity rises as the H–X bond strength falls (HF < HCl < HBr < HI).
More s-character = more acidic C–H: alkyne (25) > alkene (44) > alkane (50).
Inductive withdrawal stabilizes conjugate bases: Cl3CCOOH (0.64) ≫ CH3COOH (4.74).
Resonance beats induction: carboxylic acids (4.74) are ~10¹¹ times more acidic than alcohols (16).
Charged acids are far stronger: H3O+ (−1.74) vs. H2O (15.7).
Base strength is the mirror image: a strong base ⇄ a conjugate acid with high pK_a.
One framework unifies everything: factors that stabilize the conjugate base strengthen the acid.

Check yourself

6 review questions from the chapter. Try each one, then open the answer.

Rank H₂O, HF, NH₃, and CH₄ by increasing acid strength. Which factor explains the order?
Show answer
CH₄ (50) < NH₃ (38) < H₂O (15.7) < HF (3.2); the electronegativity of the atom bearing the proton.
Why is HI a stronger acid than HF even though iodine is less electronegative?
Show answer
The H–I bond is much weaker than H–F, so the proton leaves more easily; bond strength dominates down a group.
Why are terminal alkynes (pK_a ≈ 25) far more acidic than alkanes (pK_a ≈ 50)?
Show answer
The sp hybrid carbon has more s-character and is more electronegative, better stabilizing the acetylide conjugate base.
Why is acetic acid about 10¹¹ times more acidic than ethanol?
Show answer
The acetate conjugate base is resonance-stabilized over two oxygens, while ethoxide confines the charge to one oxygen.
Which is the stronger base, Cl^- or CH₃COO^-? Explain using the pK_a of the conjugate acids.
Show answer
CH₃COO^- is the stronger base: its conjugate acid has pK_a 4.74, far above HCl's −7; a higher conjugate-acid pK_a means a stronger base.
What is K_a for an acid with pK_a = 4.74?
Show answer
K_a = 10^-4.74 = 1.8 × 10^-5 M.

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Study tools & related lessonsKey vocabulary · Related

Key vocabulary

pKₐ: Negative base-10 log of K_a; smaller value = stronger acid.
conjugate base: The acid minus its proton; its stability determines acid strength.
electronegativity: An atom's pull on shared electrons.
bond dissociation energy: Energy needed to break a bond; weaker H–X bond = stronger acid.
s-character: Fraction of s-orbital character in a hybrid orbital.
inductive effect: Electron withdrawal through σ bonds by electronegative atoms.
resonance stabilization: Delocalization of charge in the conjugate base.
strong vs. weak base: Determined by the pK_a of the base's conjugate acid.

Sources & references

openstax.org — Organic Chemistry

This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.

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In 30 seconds

Why this matters

The college version

Core Concepts

The pKa scale

Factor 1: Electronegativity of the atom holding the proton

Factor 2: Bond strength down a group

Factor 3: Hybridization and s-character

Factor 4: Inductive effects

Factor 5: Resonance in the conjugate base

Factor 6: Charge

Base strength

Common Confusions

The same idea, in plain words

Worked examples

Key takeaways

Check yourself

Keep learning

Key vocabulary

Related lessons

Sources & references