Difficulty: Introductory | Prerequisites: Basic understanding of atoms, ions, and chemical bonding
Acid-base chemistry is one of the foundational pillars of general chemistry and comes up repeatedly in later topics (equilibrium, buffers, electrochemistry, biochemistry). This material sits early in most Gen Chem sequences because it introduces proton transfer, a concept you will use constantly. If you are comfortable with the idea that atoms can gain or lose electrons to form ions, you have what you need to start here. If the words "ion," "proton," or "chemical equation" feel unfamiliar, review your intro-to-matter notes first.
Acid-base reactions are proton-transfer reactions: an acid donates a hydrogen ion (H⁺) to a base. The strength of the acid or base depends on how readily it gives up or accepts that proton, and we quantify it with Ka (acids) or Kb (bases). When an acid and base react completely, you get a salt and water, which is called neutralisation.
Acid (Brønsted-Lowry definition)
A substance that donates a proton (H⁺) to another substance.
In simple terms, an acid is anything that hands off a hydrogen ion during a reaction.
Base (Brønsted-Lowry definition)
A substance that accepts a proton (H⁺) from another substance.
Think of a base as a proton sponge: it picks up the H⁺ that the acid releases.
Strong acid / strong base
An acid or base that dissociates completely (or very nearly so) in water.
In simple terms, if you dissolve a strong acid in water, essentially every molecule breaks apart into ions. HCl, HNO₃, and H₂SO₄ are common examples of strong acids; NaOH and KOH are common strong bases.
Weak acid / weak base
An acid or base that only partially dissociates in water, leaving an equilibrium between the intact molecule and its ions.
Think of it as a reluctant donor (or acceptor): most of the molecules stay whole. Acetic acid (CH₃COOH, the acid in vinegar) is a classic weak acid.
pH scale
A logarithmic scale from 0 to 14 used to express the acidity or basicity of a solution. pH 0 is the most acidic, pH 7 is neutral, and pH 14 is the most basic.
In simple terms, pH tells you how acidic or basic a solution is. Lower numbers mean more acidic, higher numbers mean more basic.
Neutralisation reaction
A reaction between an acid and a base that produces a salt and water.
Think of it as the acid and base cancelling each other out.
Titration
A laboratory method for determining the concentration of an unknown acid (or base) by gradually adding a solution of known concentration until the reaction is complete.
In simple terms, you drip a known base into an unknown acid, measuring how much you need to fully react, then back-calculate the acid's concentration.
Acid dissociation constant (Ka)
An equilibrium constant that measures how completely an acid donates its proton in water. A large Ka means a strong acid; a small Ka means a weak acid.
Think of Ka as a "willingness score" for proton donation.
Base dissociation constant (Kb)
The equivalent equilibrium constant for bases, measuring how readily a base accepts a proton in water. Large Kb means strong base; small Kb means weak base.
Dissociation
The process by which a compound separates into ions when dissolved in water.
In simple terms, the molecule breaks apart into charged pieces in solution.
Conjugate acid-base pair
Two species that differ by exactly one proton. When an acid donates H⁺, what remains is its conjugate base; when a base accepts H⁺, the result is its conjugate acid.
Think of them as a before-and-after pair in a proton transfer.
Every acid-base reaction is, at its core, a proton (H⁺) moving from one species to another.
The species that loses the proton is the acid; the species that gains it is the base.
After the transfer, each reactant has a conjugate partner:
Acid → loses H⁺ → becomes its conjugate base
Base → gains H⁺ → becomes its conjugate acid
Strong acids dissociate completely in water. In a solution of HCl, virtually no intact HCl molecules remain; it is all H⁺ and Cl⁻.
Common strong acids: HCl, HBr, HI, HNO₃, H₂SO₄, HClO₄
Weak acids dissociate only partially. Most molecules stay intact, and an equilibrium forms between the undissociated acid and its ions.
Example: CH₃COOH ⇌ CH₃COO⁻ + H⁺
The same logic applies to bases. NaOH (strong) dissociates fully; NH₃ (weak) only partially accepts protons from water.
pH = −log[H⁺]
Ranges from 0 (highly acidic) to 14 (highly basic), with 7 as neutral (pure water at 25 °C).
Each whole-number step represents a tenfold change in H⁺ concentration. A solution at pH 3 is ten times more acidic than one at pH 4.
pOH works the same way for hydroxide ions: pOH = −log[OH⁻]. At 25 °C, pH + pOH = 14.
When an acid reacts with a base, the products are a salt and water.
Classic example:
HCl + NaOH → NaCl + H₂O
The H⁺ from HCl combines with the OH⁻ from NaOH to form water; the remaining ions (Na⁺ and Cl⁻) form the salt.
The net ionic equation for any strong-acid/strong-base neutralisation is simply: H⁺(aq) + OH⁻(aq) → H₂O(l)
A titration lets you find the concentration of an unknown acid or base.
You add a solution of known concentration (the titrant) to the unknown until the equivalence point: the point where moles of acid equal moles of base.
An indicator (such as phenolphthalein) or a pH meter signals when you have reached the equivalence point.
From the volume and concentration of the titrant used, you calculate the unknown concentration.
Ka (acid dissociation constant): the equilibrium constant for the dissociation of an acid in water.
Larger Ka → stronger acid (dissociates more)
Smaller Ka → weaker acid
Kb (base dissociation constant): the equivalent for bases.
Larger Kb → stronger base
Smaller Kb → weaker base
For a conjugate acid-base pair: Ka × Kb = Kw = 1.0 × 10⁻¹⁴ at 25 °C. This means a strong acid has a weak conjugate base, and vice versa.
Dissociation: An acid or base dissolves in water and separates into ions.
Example: HCl(aq) → H⁺(aq) + Cl⁻(aq)
Neutralisation: An acid and a base react to produce a salt and water.
Example: HCl + NaOH → NaCl + H₂O
Reaction of a salt with an acid (reverse direction): A salt of a weak acid reacts with a strong acid to regenerate the weak acid.
Example: NaOAc + HCl → NaCl + CH₃COOH
\text{pH} = -\log[\text{H}^+]\text{pOH} = -\log[\text{OH}^-]\text{pH} + \text{pOH} = 14 \quad (\text{at } 25\,^\circ\text{C})K_a \times K_b = K_w = 1.0 \times 10^{-14} \quad (\text{at } 25\,^\circ\text{C})K_a = \frac{[\text{H}^+][\text{A}^-]}{[\text{HA}]}K_b = \frac{[\text{BH}^+][\text{OH}^-]}{[\text{B}]}Note on the source material: the PDF presents the equation [H⁺] = [OH⁻] / K. This is a rearrangement of the Ka expression for a specific scenario (where [OH⁻] and Ka are known), not a general-purpose formula. The standard forms above are what you should memorise.
Neutralisation reactions are the basis of antacid tablets: the base (often calcium carbonate or magnesium hydroxide) neutralises excess stomach acid (HCl), raising the pH and relieving discomfort.
pH control matters in agriculture, water treatment, and food science. Farmers test soil pH to decide whether to add lime (a base) or sulphur (which generates acid) so that crops grow in their preferred range. Municipal water plants adjust pH to prevent pipe corrosion and ensure safe drinking water.
Students often assume that "strong" and "concentrated" mean the same thing. They do not. Strength refers to how completely an acid dissociates; concentration is how much acid is dissolved per litre. You can have a dilute solution of a strong acid or a concentrated solution of a weak acid.
A pH of 7 is neutral only at 25 °C. At other temperatures Kw changes, and so does the neutral pH. Exams sometimes test whether you know this caveat.
Students sometimes think that neutralisation always produces a solution at pH 7. That is true only when a strong acid reacts with a strong base. A strong acid plus a weak base gives a slightly acidic salt solution; a weak acid plus a strong base gives a slightly basic one.
The logarithmic nature of pH trips people up. A pH change of 1 is a tenfold change in [H⁺], not a simple addition. pH 2 is not "twice as acidic" as pH 4; it is 100 times more acidic.
⚠️ Be able to identify the acid, base, conjugate acid, and conjugate base in any given reaction. This is one of the most common exam questions.
⚠️ Know the common strong acids and strong bases by heart. If it is not on the strong list, treat it as weak.
⚠️ Be comfortable converting between pH, pOH, [H⁺], and [OH⁻]. These conversions appear on nearly every Gen Chem exam.
⚠️ Understand what Ka and Kb tell you. A question that gives you two Ka values and asks which acid is stronger is testing whether you know that larger Ka = stronger acid.
⚠️ Titration calculations (finding concentration from volume and molarity of the titrant) are a staple of problem sets and exams.
True or false: HCl is a weak acid. (False. HCl is a strong acid; it dissociates completely.)
Fill in the blank: pH + pOH = ______ at 25 °C. (14)
True or false: A large Ka value indicates a weak acid. (False. A large Ka indicates a strong acid.)
Fill in the blank: In a neutralisation reaction, an acid and a base produce a ______ and ______. (salt and water)
True or false: pH 3 is ten times more acidic than pH 4. (True.)
Q: Write the net ionic equation for the reaction of any strong acid with any strong base.
A: H⁺(aq) + OH⁻(aq) → H₂O(l)
Q: If Ka for acetic acid is 1.8 × 10⁻⁵, is acetic acid a strong or weak acid? Explain briefly.
A: Weak acid. The Ka is very small (much less than 1), which means only a tiny fraction of acetic acid molecules donate their proton in water.
Q: You titrate 25.0 mL of an unknown HCl solution with 0.100 M NaOH. It takes 30.0 mL of NaOH to reach the equivalence point. What is the concentration of the HCl?
A: Moles of NaOH = 0.0300 L × 0.100 mol/L = 0.00300 mol. Since HCl and NaOH react 1:1, moles of HCl = 0.00300 mol. Concentration of HCl = 0.00300 mol / 0.0250 L = 0.120 M.
Q: What is the pH of a solution with [H⁺] = 1.0 × 10⁻³ M?
A: pH = −log(1.0 × 10⁻³) = 3.0
Q: Explain why mixing a strong acid with a weak base does not produce a neutral (pH 7) solution.
A: The conjugate acid of the weak base is itself a weak acid. After neutralisation, this conjugate acid partially dissociates and donates protons to the solution, making the final pH less than 7.
This material connects directly to chemical equilibrium: Ka and Kb are specific cases of the equilibrium constant Keq, and the same ICE-table method you use for equilibrium problems applies to weak acid/base calculations.
Buffers (the next major topic in most courses) are solutions of a weak acid and its conjugate base. Understanding conjugate pairs here is prerequisite knowledge for buffer problems.
In biochemistry, nearly all enzyme activity depends on the pH of the surrounding solution, which links back to everything in these notes. Amino acids themselves are acid-base species with multiple pKa values.
acid-base reactions, proton transfer, Brønsted-Lowry acid, Brønsted-Lowry base, strong acid, weak acid, strong base, weak base, pH, pOH, hydrogen ion concentration, hydroxide ion, neutralisation, neutralization, salt and water, titration, equivalence point, indicator, acid dissociation constant, Ka, base dissociation constant, Kb, Kw, conjugate acid, conjugate base, conjugate pair, dissociation, ionisation, HCl, NaOH, acetic acid, CH₃COOH, general chemistry, Gen Chem, acid strength, base strength, pH scale, logarithmic scale, proton donor, proton acceptor