Difficulty: Intermediate | Prerequisites: Lewis structures, functional groups, electronegativity basics
Big picture: Acid-base chemistry sits beneath nearly every reaction in organic chemistry. This section introduces the Brønsted-Lowry framework for understanding proton transfer, then builds toward the quantitative tools (Ka and pKa) that let you predict which way an equilibrium will lie. If you can read a pKa table and compare conjugate bases, you can predict the outcome of most acid-base reactions. You should already be comfortable drawing Lewis structures and identifying lone pairs and sigma bonds.
In Brønsted-Lowry chemistry, acids donate protons and bases accept them, producing conjugate pairs. The strength of an acid is measured by Ka (how much it dissociates in water) and its more convenient logarithmic form, pKa. Lower pKa means a stronger acid, and the equilibrium of any acid-base reaction favours the side with the weaker acid (higher pKa).
Brønsted-Lowry acid
A molecule (or atom in a molecule) that can donate a proton (H⁺). Structurally, it has a sigma bond between the acidic atom and hydrogen.
Think of it as: the species that gives up its hydrogen.
Brønsted-Lowry base
A molecule (or atom in a molecule) that can accept a proton, using a lone pair of electrons.
Think of it as: the species with a free lone pair ready to grab a proton.
Conjugate base
The species formed after an acid donates its proton. It is the deprotonated form of the acid.
In simple terms, this means: whatever the acid becomes once it loses H⁺.
Conjugate acid
The species formed after a base accepts a proton. It is the protonated form of the base.
In simple terms, this means: whatever the base becomes once it gains H⁺.
Sigma bond (σ bond)
A single covalent bond between two atoms, represented by a single line. In acid-base chemistry, this is the bond between the acidic atom and the hydrogen that may dissociate.
Lone pair
An unshared (non-bonding) pair of electrons on an atom, represented as two dots. This is what a Brønsted-Lowry base uses to accept a proton.
Ka (acid dissociation constant)
The equilibrium constant for the dissociation of an acid in water, with the concentration of water removed from the expression because water is present in vast excess (~55 M).
Think of it as: a number that tells you how completely an acid falls apart in water. Large Ka means lots of dissociation.
pKa
The negative logarithm of Ka: pKa = -log(Ka). A more convenient scale for comparing acid strengths.
Think of it as: a compact way to rank acids. Lower pKa = stronger acid. Each unit of pKa corresponds to a tenfold change in Ka.
Keq (equilibrium constant)
The ratio of product concentrations to reactant concentrations at equilibrium. Ka is derived from Keq by removing the constant water term.
Formal charge
The charge assigned to an atom in a molecule assuming all bonding electrons are shared equally. For acidic atoms, it is most often +1 or 0 (rarely -1). For basic atoms, it is most often negative or neutral.
Amphoteric
Describes a molecule capable of acting as either an acid or a base, depending on the reaction partner. Water and ammonia are classic examples.
An acid donates a proton; a base accepts a proton. Together they produce a conjugate base (from the acid) and a conjugate acid (from the base).
The general reaction: A-H + B: → A:⁻ + B-H⁺
Acid-base reactions are best understood as equilibria, not one-way reactions. Protons are always in transit between competing lone pairs.
A Brønsted-Lowry acid is any atom with a sigma bond to hydrogen.
The formal charge of the acidic atom is most commonly +1 or 0 (rarely -1), though the overall molecular charge can be negative (e.g. bisulfate HSO₄⁻, bicarbonate HCO₃⁻).
A "stronger" acid has a more polarised, more easily dissociated sigma bond to hydrogen. After reaction, its conjugate base has a more diffuse (delocalised) lone pair that is harder to reprotonate.
A Brønsted-Lowry base is any atom with a lone pair of electrons (shown as two dots).
The formal charge of the basic atom is most often negative or neutral (rarely positive).
A negative formal charge does not guarantee reactivity: ions like nitrate (NO₃⁻), bromide (Br⁻), chloride (Cl⁻), and sulfate (SO₄²⁻) are quite stable and unreactive as bases.
A "stronger" base has a more focused (localised), bond-like lone pair. After reaction, its conjugate acid has a sigma bond to hydrogen that is harder to dissociate.
When a strong acid reacts with a strong base, the equilibrium favours products: the transition state shows a long (mostly broken) A-H bond and a short (mostly formed) B-H bond, and the reaction releases enthalpy (-ΔH).
When a weak acid reacts with a weak base (or the reverse of a favoured reaction), the equilibrium favours reactants: the transition state shows a reluctant bond-breaking, and the reaction absorbs enthalpy (+ΔH).
Unsymmetrical equilibrium arrows (one longer than the other) indicate which side is favoured.
Any acid can be placed in water and the resulting pH measured. The equilibrium constant Keq for proton dissociation is determined from the concentrations of products and reactants.
Because water is the solvent and present in enormous excess (~55 M), its concentration is effectively constant and is removed from the expression. This gives Ka.
Ka = [conjugate base][H₃O⁺] / [acid]
pKa = -log(Ka)
Strong acids dissociate completely: large Ka, negative pKa (e.g. HCl: Ka = 1 × 10⁸, pKa = -8).
Weak acids barely dissociate: small Ka, positive pKa (e.g. acetic acid: Ka = 1.58 × 10⁻⁵, pKa = 4.8).
Extremely weak acids are essentially non-acidic: tiny Ka, very large pKa (e.g. isobutane: Ka = 1 × 10⁻⁵³, pKa = 53).
HCl in water: pH measurement gives a 10,000-to-1 ratio of dissociated to unionised acid. Ka = 1 × 10⁸, pKa = -8. A "strong acid" means you have a very high probability of encountering a free proton in solution.
Acetic acid in water: pH gives a 1-to-63,000 ratio of dissociated to unionised acid. Ka = 1.58 × 10⁻⁵, pKa = 4.8. A "weak acid" means protons dissociated from the acid are rare.
Isobutane (alkane C-H) in water: the odds of finding a dissociated proton are 1 in 10⁵³. This is effectively zero. Alkane C-H bonds are not meaningfully acidic.
Compare the pKa of the acid on the left with the pKa of the conjugate acid on the right.
The equilibrium favours the side with the weaker acid (higher pKa).
To quantify how strongly an equilibrium is favoured: pKreaction = pKa(acid) - pKa(conjugate acid). A negative pKreaction means products are favoured; a positive one means reactants are favoured.
Expression | Formula | Notes |
|---|---|---|
Equilibrium constant | Keq = [A⁻][H₃O⁺] / [AH][H₂O] | Products over reactants |
Acid dissociation constant | Ka = Keq × [H₂O] = [A⁻][H₃O⁺] / [AH] | Water term removed (constant ~55 M) |
pKa | pKa = -log(Ka) | Lower pKa = stronger acid |
Reaction favourability | pKreaction = pKa(acid) - pKa(conjugate acid) | Negative = products favoured; positive = reactants favoured |
The pKreaction gives a rough order-of-magnitude estimate: a pKreaction of -11 means products are favoured by about 10¹¹ to 1.
Students often think "strong" and "weak" are absolute categories. They are not. Acid-base chemistry operates on a continuum, and the labels "stronger" and "weaker" are always relative to the other species in the reaction.
Students often think a negative formal charge on an atom means it must be a reactive base. It does not. Anions like nitrate, bromide, and chloride are quite stable and unreactive as bases because of structural factors (size, resonance, electronegativity).
Students often confuse the direction of the pKa scale. A lower pKa means a stronger acid (more dissociation), not a weaker one. If you remember that HCl has a pKa of -8 and methane has a pKa of 48, the direction becomes intuitive.
Students often think that the equilibrium favours the side with the stronger acid. The opposite is true: equilibrium favours the side with the weaker acid (higher pKa), because the stronger acid dissociates more readily and pushes the reaction toward products.
⚠️ You will almost certainly be asked to predict the favoured side of an acid-base equilibrium using pKa values. The rule: equilibrium favours the side with the weaker acid (higher pKa).
⚠️ Know how to derive Ka from Keq (remove the water term) and convert between Ka and pKa. These steps are frequently tested.
⚠️ Be comfortable reading a pKa table: strong acids at the top (negative pKa), weak acids in the middle, essentially non-acidic molecules at the bottom (pKa > 40).
⚠️ The concept that "strong acid + strong base → weak acid + weak base" is a simplification. The more precise framing is: the equilibrium favours the side with the more stable conjugate base and the weaker conjugate acid.
True or false: A pKa of -3 indicates a stronger acid than a pKa of 5. (True)
Fill in the blank: Ka = Keq × ____. (The concentration of water, [H₂O])
True or false: Bromide (Br⁻) is a strong base because it carries a negative charge. (False, it is a very stable, unreactive conjugate base)
Fill in the blank: The equilibrium of an acid-base reaction favours the side with the ____ acid. (Weaker / higher pKa)
True or false: pKa = -log(Ka). (True)
Q: Define a Brønsted-Lowry acid and a Brønsted-Lowry base in structural terms.
A: A Brønsted-Lowry acid is any atom (or atom in a molecule) with a sigma bond to hydrogen. A Brønsted-Lowry base is any atom with a lone pair of electrons capable of accepting a proton.
Q: Given that acetic acid has a pKa of 5 and water has a pKa of 16, calculate pKreaction for the reaction of acetic acid with hydroxide. Which side is favoured?
A: pKreaction = pKa(acid) - pKa(conjugate acid) = 5 - 16 = -11. Products are favoured by approximately 10¹¹ to 1.
Q: Why is the concentration of water removed from the equilibrium expression when calculating Ka?
A: Water is the solvent and present in such excess (~55 M) that its concentration is effectively constant and unaffected by the equilibrium. Multiplying Keq by this constant gives Ka.
Q: An acid has a Ka of 1 × 10⁻⁷. What is its pKa, and would you classify it as strong, weak, or essentially non-acidic?
A: pKa = -log(1 × 10⁻⁷) = 7. This is a weak acid (comparable to hydrogen sulfide, H₂S).
Q: Why does a negative formal charge on an atom not guarantee that it will act as a base?
A: Structural factors such as atom size, electronegativity, resonance delocalisation, and solvation stability can make an anion so stable that its lone pair is effectively unreactive. Examples include nitrate, chloride, and bromide.
This material connects directly to predicting acid and base strength using the four structural factors (elemental, resonance, hybridisation, inductive effects), covered in Part 2 of these notes. It also underpins nucleophilicity and electrophilicity concepts later in the course: a good base is often a good nucleophile, though the two do not always correlate.
The pKa table (Table 7.1) will reappear in nearly every reaction mechanism chapter. Memorising the approximate pKa ranges for common functional groups (carboxylic acids ~5, alcohols ~16, amines ~38, alkanes ~50) pays off throughout the course.
Brønsted-Lowry, proton transfer, proton donor, proton acceptor, conjugate pair, acid dissociation constant, Ka, pKa, pKreaction, equilibrium constant, Keq, strong acid, weak acid, formal charge, sigma bond, lone pair, amphoteric, hydronium, conjugate base stability, acid-base equilibrium, organic acid-base chemistry, OSU organic chemistry, Chapter 7.4