Acid-Base Chemistry – OC 101, Unit 1
Source: Organic Chemistry I lecture notes, University of Minnesota Twin Cities
Difficulty: Introductory
Prerequisites: General chemistry fundamentals (atomic structure, bonding basics, electronegativity). If you are not comfortable with electron pairs and electronegativity, review those topics first.
Acid-base chemistry is one of the first major conceptual frameworks you will meet in organic chemistry. It underpins reaction mechanisms, nucleophilicity, leaving-group ability, and almost every transformation you will study from here on. The three theories covered here (Arrhenius, Brønsted-Lowry, Lewis) each widen the lens: Arrhenius is the narrow, aqueous-only view from general chemistry; Brønsted-Lowry introduces proton transfer in any solvent; Lewis extends the idea to electron-pair donation and acceptance, which is exactly the language organic chemistry uses for curved-arrow mechanisms. If you are coming in cold, know that this material bridges gen chem and orgo, and you will lean on it constantly.
Acids donate protons (or accept electron pairs); bases accept protons (or donate electron pairs). Three theories describe this with increasing generality: Arrhenius (aqueous only), Brønsted-Lowry (proton transfer in any medium), and Lewis (electron-pair transfer). Strength is measured by pH, pOH, and the acid dissociation constant Ka.
Acid (general)
A substance that produces hydrogen ions (H+) when dissolved in water. In simple terms, an acid is a proton donor.
Base (general)
A substance that produces hydroxide ions (OH−) when dissolved in water. In simple terms, a base is a proton acceptor.
Arrhenius acid
A substance that increases the concentration of H+ ions in aqueous solution. Think of it as the narrowest definition of an acid, limited to water-based systems.
Arrhenius base
A substance that increases the concentration of OH− ions in aqueous solution.
Brønsted-Lowry acid
A proton (H+) donor. This definition works in any solvent, not just water.
Brønsted-Lowry base
A proton (H+) acceptor.
Conjugate acid-base pair
Two species related by the transfer of a single proton. When an acid donates a proton, it becomes its conjugate base; when a base accepts a proton, it becomes its conjugate acid.
Lewis acid
An electron-pair acceptor. Think of it as a species with an empty or electron-poor orbital looking to receive electrons.
Lewis base
An electron-pair donor. In organic chemistry, nucleophiles are Lewis bases.
pH
A logarithmic scale measuring the concentration of H+ ions in solution. pH 7 is neutral; below 7 is acidic; above 7 is basic.
pOH
A logarithmic scale measuring the concentration of OH− ions. pOH 7 is neutral; below 7 is basic; above 7 is acidic. pH + pOH = 14 in aqueous solution at 25 °C.
Ka (acid dissociation constant)
A quantitative measure of acid strength in solution. A larger Ka means a stronger acid (it dissociates more completely).
pKa
The negative log of Ka. A smaller pKa means a stronger acid. This is the scale organic chemists use most often.
Neutralisation reaction
A reaction between an acid and a base that produces a salt and water.
Displacement reaction (acid-base context)
A reaction involving the transfer of a proton from one species to another, producing a new substance and often releasing a gas.
Arrhenius theory (narrowest scope)
Acid: increases H+ concentration in aqueous solution
Base: increases OH− concentration in aqueous solution
Only applies to reactions in water
Does not account for electron transfer or non-aqueous solvents
Useful as a starting point, but too limited for organic chemistry
Brønsted-Lowry theory (the workhorse for proton-transfer reactions)
Acid: a proton (H+) donor
Base: a proton (H+) acceptor
Works in any solvent, not only water
Introduces conjugate acid-base pairs: every acid-base reaction produces a conjugate acid and a conjugate base
Example: HCl + H₂O → H₃O+ + Cl−
HCl donates a proton → it is the Brønsted-Lowry acid
H₂O accepts the proton → it is the Brønsted-Lowry base
H₃O+ is the conjugate acid of water
Cl− is the conjugate base of HCl
Lewis theory (broadest scope, essential for organic mechanisms)
Acid: an electron-pair acceptor
Base: an electron-pair donor
Applies to covalent and ionic compounds alike
Does not require a proton to be involved at all
This is the definition you will use most in organic chemistry, because curved-arrow mechanisms track electron-pair movement
A Lewis acid-base pair refers to two species related by the transfer (or sharing) of an electron pair
Neutralisation reactions
An acid reacts with a base to produce a salt and water
Classic example: HCl + NaOH → NaCl + H₂O
The H+ from the acid combines with the OH− from the base to form water; the remaining ions form the salt
Neutralisation is the basis of titration experiments in the lab
Displacement reactions
A proton transfers from an acid to another substance (often a metal), producing a new compound and often releasing hydrogen gas
Example: H₂SO₄ + Cu → CuSO₄ + H₂
Note: this specific example (sulfuric acid + copper) is actually a redox reaction, not a simple proton transfer. It requires hot, concentrated H₂SO₄. The source material simplifies this, but be aware that the mechanism is more involved than a straightforward acid-base displacement
pH scale
Measures H+ concentration: pH = −log[H+]
pH 7 = neutral (pure water at 25 °C)
pH < 7 = acidic
pH > 7 = basic
pOH scale
Measures OH− concentration: pOH = −log[OH−]
pOH 7 = neutral
pOH < 7 = basic
pOH > 7 = acidic
Key relationship: pH + pOH = 14 (at 25 °C in aqueous solution)
Ka (acid dissociation constant)
Quantifies how completely an acid dissociates in solution
Larger Ka = stronger acid (dissociates more)
Strong acids (HCl, HBr, HI, H₂SO₄, HNO₃, HClO₄) have very large Ka values
Weak acids (acetic acid, carbonic acid) have small Ka values
pKa
pKa = −log(Ka)
Smaller pKa = stronger acid
Organic chemists almost always use pKa rather than Ka, because the numbers are easier to compare
A difference of one pKa unit represents a tenfold difference in acid strength
pH = −log[H+]
pOH = −log[OH−]
pH + pOH = 14 (aqueous solution, 25 °C)
Ka = [H+][A−] / [HA] (for a generic weak acid HA)
pKa = −log(Ka)
Neutralisation: HCl + NaOH → NaCl + H₂O
Acid-base chemistry is everywhere outside the lecture hall. Antacid tablets neutralise excess stomach acid (HCl) by delivering a weak base such as calcium carbonate or magnesium hydroxide. Swimming pool pH is monitored and adjusted with acids and bases to keep the water safe and the chlorine effective. In biochemistry, enzyme activity depends on the pH of the surrounding environment, and blood pH is tightly buffered near 7.4; even small deviations can be life-threatening.
Students often think "strong acid" means "concentrated acid." It does not. Strength refers to the degree of dissociation (Ka), while concentration refers to the amount of acid dissolved per unit volume. You can have a dilute solution of a strong acid.
Students often confuse Brønsted-Lowry and Lewis definitions. Every Brønsted-Lowry acid is a Lewis acid, but not every Lewis acid is a Brønsted-Lowry acid (Lewis acids need not involve protons at all). BF₃, for example, is a Lewis acid but has no proton to donate.
Students sometimes assume that a low pH number means "less acid." The opposite is true. Lower pH = higher H+ concentration = more acidic.
Students forget that water is amphoteric: it can act as either an acid or a base depending on what it reacts with.
⚠️ Be able to classify a species as an acid or base under all three theories (Arrhenius, Brønsted-Lowry, Lewis). Exam questions often ask you to identify which definition applies.
⚠️ Know how to identify conjugate acid-base pairs in a given reaction. This is a bread-and-butter exam question.
⚠️ Understand the relationship between Ka, pKa, and acid strength. You will need to predict which side of an equilibrium is favoured by comparing pKa values.
⚠️ Be comfortable converting between pH, pOH, [H+], and [OH−]. Calculation questions are common.
⚠️ The distinction between strong/weak and concentrated/dilute is a perennial exam trap.
True or False: A Lewis acid must contain a hydrogen atom.
Fill in the blank: In the reaction HCl + H₂O → H₃O+ + Cl−, the conjugate base of HCl is ______.
True or False: A solution with pH 3 is more acidic than a solution with pH 5.
Fill in the blank: pH + pOH = ______ in aqueous solution at 25 °C.
True or False: The Arrhenius definition of acids and bases can be applied to non-aqueous solvents.
Answers: 1. False (Lewis acids accept electron pairs, no proton required). 2. Cl−. 3. True. 4. 14. 5. False (Arrhenius applies only to aqueous solutions).
Q: Define an acid and a base according to the Brønsted-Lowry theory.
A: A Brønsted-Lowry acid is a proton (H+) donor. A Brønsted-Lowry base is a proton (H+) acceptor.
Q: In the reaction HCl + H₂O → H₃O+ + Cl−, identify the conjugate acid-base pairs.
A: Pair 1: HCl (acid) and Cl− (its conjugate base). Pair 2: H₂O (base) and H₃O+ (its conjugate acid).
Q: What is the main limitation of the Arrhenius theory compared with the Brønsted-Lowry theory?
A: The Arrhenius theory only applies to aqueous solutions and does not account for proton transfer in non-aqueous solvents or for reactions that do not involve H+ or OH− directly.
Q: A solution has a pH of 4. What is its pOH? Is the solution acidic or basic?
A: pOH = 14 − 4 = 10. The solution is acidic (pH < 7).
Q: Explain why BF₃ is classified as a Lewis acid even though it contains no hydrogen atoms.
A: BF₃ has an empty p-orbital on boron, making it an electron-pair acceptor. Under the Lewis definition, any species that accepts an electron pair is an acid, regardless of whether a proton is involved.
Q: If acid A has a Ka of 1.8 × 10⁻⁵ and acid B has a Ka of 6.3 × 10⁻³, which is the stronger acid?
A: Acid B is stronger. A larger Ka means greater dissociation in solution.
This material connects directly to nucleophiles and electrophiles in organic chemistry: a nucleophile is a Lewis base (electron-pair donor) and an electrophile is a Lewis acid (electron-pair acceptor). Every mechanism you draw with curved arrows is tracking Lewis acid-base interactions.
It also connects to equilibrium and thermodynamics: predicting which side of an acid-base reaction is favoured uses the same equilibrium reasoning from general chemistry, applied through pKa comparisons.
Finally, understanding conjugate acid-base pairs is essential for buffer chemistry and for biological systems where pH regulation matters (enzyme kinetics, drug design, blood chemistry).
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