Difficulty: Intermediate | Prerequisites: Acids and Bases Defined, pH/pOH Scales, Strong vs. Weak Acids and Bases (all Ch. 12)
Titrations use a known-concentration solution to find an unknown concentration by reacting them until both acid and base are fully consumed (the endpoint). Buffers are mixtures of a weak acid and its conjugate base that resist pH changes when small amounts of acid or base are added. Dissolving ionic compounds in water can shift the pH depending on whether the component ions come from strong or weak acid/base parents.
Titration
A laboratory technique for determining an unknown concentration by reacting a solution of known concentration (the titrant) with the unknown solution until the reaction is complete. It is a form of stoichiometry with solutions.
Think of it as: you drip one solution into another, measuring exactly how much it takes to use up all the acid or base.
Endpoint (equivalence point)
The point in a titration at which all of the acid and all of the base reactants have been consumed. No excess of either remains.
In simple terms, the endpoint is when the reaction is perfectly finished, with nothing left over on either side.
Indicator
A substance that changes colour at or near a specific pH, signalling that the endpoint of a titration has been reached.
Think of it as a colour alarm that tells you to stop adding titrant.
Buffer
An aqueous solution that resists changes in pH when small amounts of acid or base are added. It consists of a weak acid and a similar concentration of its conjugate base.
In simple terms, a buffer is a chemical shock absorber for pH.
Conjugate base (in a buffer context)
The species formed when a weak acid loses a proton. In a buffer, the conjugate base is present in comparable concentration to the weak acid, ready to neutralise any added H⁺.
Think of it as the "backup" that mops up extra acid before the pH can shift much.
Salt hydrolysis
The reaction of ions from a dissolved ionic compound with water, which can produce H⁺ or OH⁻ and shift the pH away from 7.
In simple terms, some salts make water acidic or basic because their ions react with water.
A titration is stoichiometry with solutions: you use molarity and volume to find moles, then use the balanced equation to convert between acid and base moles.
Setup: the solution of unknown concentration sits in a flask. The solution of known concentration (titrant) is added from a burette until the endpoint is reached.
An indicator in the flask changes colour at the endpoint to signal when to stop.
Reaction: 2 HCl(aq) + Ca(OH)₂(aq) → 2 H₂O(l) + CaCl₂(aq)
Given: 25.0 mL of HCl (unknown M), titrated with 0.0495 M Ca(OH)₂ until endpoint at 47.8 mL added.
Step 1: Find moles of Ca(OH)₂.
0.0495 mol/L × 0.0478 L = 2.366 × 10⁻³ mol Ca(OH)₂
Step 2: Use the mole ratio from the balanced equation (2 mol HCl : 1 mol Ca(OH)₂).
2.366 × 10⁻³ mol × (2 mol HCl / 1 mol Ca(OH)₂) = 4.732 × 10⁻³ mol HCl
Step 3: Find molarity of HCl.
M = 4.732 × 10⁻³ mol / 0.0250 L = 0.189 M HCl
Resist changes in pH when small amounts of acid or base are added.
Composition: a weak acid and a similar amount of its conjugate base.
How they work:
If H⁺ is added, the conjugate base (e.g. C₂H₃O₂⁻) neutralises it, forming the weak acid (HC₂H₃O₂). The pH barely changes.
If OH⁻ is added, the weak acid donates a proton to neutralise it, forming the conjugate base and water. The pH barely changes.
Examples of buffer pairs:
Acetic acid / sodium acetate (HC₂H₃O₂ / NaC₂H₃O₂)
Carbonic acid / sodium bicarbonate (H₂CO₃ / NaHCO₃)
A strong acid alone (e.g. HCl) cannot form a buffer. Buffers require a weak acid/base equilibrium.
When an ionic compound dissolves, its ions may or may not react with water. Whether they do depends on whether the parent acid or base was strong or weak.
NaCl (from strong acid HCl + strong base NaOH):
Na⁺ + H₂O → no reaction
Cl⁻ + H₂O → no reaction
Result: neutral solution, pH = 7
NaCN (from weak acid HCN + strong base NaOH):
Na⁺ + H₂O → no reaction
CN⁻ + H₂O ⇌ HCN + OH⁻ (CN⁻ is the conjugate base of a weak acid, so it reacts)
Result: basic solution, pH > 7
NH₄Cl (from strong acid HCl + weak base NH₃):
NH₄⁺ + H₂O ⇌ NH₃ + H₃O⁺ (NH₄⁺ is the conjugate acid of a weak base, so it reacts)
Cl⁻ + H₂O → no reaction
Result: acidic solution, pH < 7
Rule of thumb:
Strong acid + strong base parent → neutral salt
Strong base + weak acid parent → basic salt
Strong acid + weak base parent → acidic salt
Titration: finding moles from molarity and volume
moles = Molarity × Volume (in litres)
Titration: finding unknown molarity
M_unknown = moles_unknown / Volume_unknown (in litres)
Always use the balanced equation's mole ratio to convert between acid moles and base moles before dividing by volume.
Buffer composition rule
A buffer requires a weak acid + its conjugate base in similar concentrations. No formula to memorise here beyond the equilibrium expression for the weak acid, but the Henderson-Hasselbalch equation (pH = pK_a + log([A⁻]/[HA])) may appear in later chapters or courses.
Titrations are used in water treatment to measure acidity and in the food industry to check the acid content of juices and wines. Blood is a buffer system: the carbonic acid/bicarbonate pair (H₂CO₃ / HCO₃⁻) holds blood pH near 7.4, and even small departures cause serious medical problems.
Students often think any acid-base mixture is a buffer. It is not. A buffer specifically requires a weak acid and its conjugate base in similar amounts. A strong acid mixed with a strong base just neutralises.
Students sometimes confuse the endpoint of a titration with pH 7. The endpoint is where moles of acid equal moles of base (stoichiometrically); the pH at the endpoint is only 7 if both the acid and the base are strong. A weak acid titrated with a strong base has an endpoint above pH 7.
Students often assume all dissolved salts are neutral. They are not. The pH depends on whether the ions come from strong or weak parents.
Students forget to use the mole ratio from the balanced equation in titration calculations. If the ratio is 2:1 rather than 1:1, skipping it gives the wrong answer.
⚠️ Titration problems nearly always require you to convert mL to L before multiplying by molarity. Forgetting this is the most common arithmetic mistake.
⚠️ Know how to predict whether a salt solution is acidic, basic, or neutral. The method: identify the parent acid and parent base, then check whether each is strong or weak.
⚠️ Exam questions on buffers often ask why a strong acid cannot form a buffer. The answer: strong acids dissociate completely, so there is no equilibrium between the acid and its conjugate base to absorb added H⁺ or OH⁻.
⚠️ Be ready to explain the mechanism of a buffer in words: what happens when H⁺ is added (conjugate base neutralises it) and what happens when OH⁻ is added (weak acid neutralises it).
True or false: The endpoint of a titration is the point where all acid and base reactants have been consumed. (True)
Fill in the blank: A buffer consists of a ______ acid and a similar amount of its ______. (weak; conjugate base)
True or false: A solution of NaCl dissolved in water has a pH of 7. (True)
Fill in the blank: NH₄Cl dissolved in water produces a(n) ______ solution because NH₄⁺ is the conjugate acid of a weak base. (acidic)
True or false: A strong acid like HCl can form a buffer with its conjugate base Cl⁻. (False)
Q: You titrate 25.0 mL of an unknown HCl solution with 0.0495 M Ca(OH)₂. The endpoint is reached after 47.8 mL of Ca(OH)₂ has been added. The balanced equation is 2 HCl + Ca(OH)₂ → 2 H₂O + CaCl₂. What is the molarity of the HCl?
A: Moles of Ca(OH)₂ = 0.0495 × 0.0478 = 2.366 × 10⁻³ mol. Moles of HCl = 2.366 × 10⁻³ × 2 = 4.732 × 10⁻³ mol. M = 4.732 × 10⁻³ / 0.0250 = 0.189 M HCl.
Q: Explain why a mixture of HC₂H₃O₂ and NaC₂H₃O₂ acts as a buffer, but a mixture of HCl and NaCl does not.
A: HC₂H₃O₂ is a weak acid with an equilibrium between the acid and its conjugate base (C₂H₃O₂⁻). When H⁺ is added, the conjugate base absorbs it; when OH⁻ is added, the weak acid neutralises it. HCl is a strong acid that dissociates completely, so there is no intact acid left to neutralise added OH⁻, and Cl⁻ is far too weak a base to absorb added H⁺.
Q: Predict whether a solution of NaCN in water is acidic, basic, or neutral. Explain.
A: Basic. Na⁺ comes from a strong base (NaOH) and does not react with water. CN⁻ comes from a weak acid (HCN), so it reacts with water: CN⁻ + H₂O ⇌ HCN + OH⁻. The production of OH⁻ makes the solution basic.
Q: An ionic compound dissolves in water and produces a solution with pH < 7. What can you conclude about its parent acid and parent base?
A: The compound is derived from a strong acid and a weak base. The conjugate acid of the weak base donates protons to water, producing H₃O⁺ and lowering the pH.
Titrations build directly on the stoichiometry and molarity skills from earlier chapters. Buffer chemistry connects to the Henderson-Hasselbalch equation and equilibrium constants (K_a), which are explored in greater depth in later coursework. Understanding salt hydrolysis ties back to the strong/weak distinction and conjugate acid-base pairs covered earlier in this chapter.
titration, endpoint, equivalence point, indicator, burette, stoichiometry with solutions, buffer, buffer solution, weak acid conjugate base, acetic acid sodium acetate buffer, carbonic acid bicarbonate buffer, Henderson-Hasselbalch, salt hydrolysis, pH of salts, NaCl neutral, NaCN basic, NH₄Cl acidic, strong acid weak base salt, strong base weak acid salt, CHEM 2, general chemistry, Chapter 12