Measuring Concentration Through Controlled Reaction
What a titration actually measures, and why the equivalence point isn't always pH 7
Titration is a laboratory technique for determining the unknown concentration of an acid or base by reacting it, in a slow and controlled way, with a solution of known concentration (the titrant) until the reaction is exactly complete. A measured volume of the unknown solution is placed in a flask, and titrant is added gradually — typically from a burette, which allows precise volume measurement — until the reaction reaches its equivalence point: the point at which the moles of acid exactly equal the moles of base that have reacted, based on the balanced neutralization equation.
The equivalence point is a stoichiometric concept — it's defined purely by mole ratios, not by pH. This is the single most important and most frequently misunderstood idea in titration: the equivalence point does not automatically correspond to pH 7. Whether the solution at the equivalence point is actually neutral, acidic, or basic depends entirely on the identity of the specific acid and base involved in the titration, because the salt formed at the equivalence point may or may not undergo hydrolysis (a topic covered in depth in the Salt Hydrolysis lesson).
A titration curve is a plot of pH (on the y-axis) versus volume of titrant added (on the x-axis), and its overall shape — including where the equivalence point falls on the pH axis, and how sharply the pH changes near that point — depends directly on whether the acid and base involved are strong or weak.
💡 Three Titration Curve Shapes, Three Different Equivalence Point pH Values
Strong acid + strong base: the equivalence point occurs at exactly pH 7. Neither ion produced by the neutralization reaction (for example, Na⁺ and Cl⁻ from HCl + NaOH) undergoes any meaningful hydrolysis with water, so the resulting salt solution is neutral. This titration curve also shows the sharpest, steepest pH jump near the equivalence point, making the endpoint easiest to detect precisely.
Weak acid + strong base: the equivalence point occurs above pH 7 (basic). At the equivalence point, the weak acid has been fully converted into its conjugate base, which is a comparatively strong enough base (relative to the original weak acid's strength) to hydrolyze water and produce a basic solution — exactly the salt hydrolysis logic that connects directly to the inverse Ka/Kb relationship covered in the Ka & Kb lesson. This curve also shows a noticeably less steep pH jump near the equivalence point compared to the strong/strong case, and includes a distinctive buffering region partway through the titration.
Strong acid + weak base: the equivalence point occurs below pH 7 (acidic), by the mirror-image logic — the weak base has been fully converted into its conjugate acid, which hydrolyzes water to produce a mildly acidic solution.
Curve
Reading a titration curve
A titration curve has several recognizable landmark features. The initial pH (before any titrant is added) reflects the pH of the original unknown solution alone. As titrant is gradually added, the pH changes gradually at first — for a weak acid or weak base being titrated, this gradual region is actually a natural buffering region, since a mixture of the weak acid and its increasingly forming conjugate base is present simultaneously, resisting large pH changes exactly like the buffer systems covered in the Buffers lesson. Near the equivalence point, the pH changes very rapidly with only a tiny additional volume of titrant — this steep, near-vertical jump is the most useful visual feature of the curve, since its midpoint marks the equivalence point itself. Beyond the equivalence point, the curve levels off again, now reflecting the pH contributed by excess titrant.
For a weak acid titrated with a strong base, the pH exactly halfway to the equivalence point equals the acid's pKa — a direct, practical link back to the Henderson-Hasselbalch equation from the Buffers lesson, and a common way titration data is used to experimentally determine an unknown acid's Ka.
Ind
Choosing the right indicator
Because the equivalence point is defined by stoichiometry, not by a visible signal, chemists typically use a chemical indicator — a substance that changes color at a specific pH range — to visually detect when the reaction has reached (or very nearly reached) the equivalence point. The point at which the indicator actually changes color is technically called the endpoint, and choosing an indicator whose color-change range closely matches the equivalence point's actual pH is essential for the endpoint to closely and accurately approximate the true equivalence point.
For a strong acid/strong base titration (equivalence point at pH 7), almost any indicator that changes color near neutral pH works well, since the pH jump near the equivalence point is so steep that even a slightly mismatched indicator will still change color very close to the true equivalence point. For a weak acid/strong base titration (equivalence point above pH 7) or a strong acid/weak base titration (equivalence point below pH 7), the indicator must be chosen more carefully, since the pH change near the equivalence point is less steep, making a poorly matched indicator noticeably less accurate.
Phenolphthalein changes color in the roughly pH 8.2-10 range, making it a good choice for a weak acid/strong base titration (equivalence point above 7), but a poor choice for a strong acid/weak base titration (equivalence point below 7), where an indicator like methyl orange (changing color in a lower, more acidic pH range) would be more appropriate.
Calc
Calculating unknown concentration from titration data
The core calculation in any titration uses the stoichiometry at the equivalence point: moles of acid = moles of base (for a 1:1 reacting ratio; the ratio must be adjusted for acids or bases that release more than one proton or hydroxide per formula unit). Since moles = concentration × volume, and the volume of titrant needed to reach the equivalence point is measured directly from the burette, the unknown concentration can be calculated directly: (concentration of unknown) × (volume of unknown) = (concentration of titrant) × (volume of titrant used), rearranged to solve for whichever concentration is unknown.
If 25.0 mL of an unknown HCl solution is titrated to the equivalence point by exactly 30.0 mL of 0.100 M NaOH, then moles of NaOH used = 0.100 M × 0.0300 L = 0.00300 mol, which equals the moles of HCl originally present (1:1 ratio), giving [HCl] = 0.00300 mol / 0.0250 L = 0.120 M.
🔬 Applied Scenario — Titration in Quality Control and Research
Titration remains one of the most widely used quantitative techniques in chemistry, valued for its precision and relatively simple equipment requirements.
A
Determining the concentration of a commercial product. Titration is routinely used in quality control to verify the actual concentration of acidic or basic components in commercial products — such as verifying the acetic acid concentration in vinegar, or checking that a cleaning product's active base concentration matches its label.
B
Determining an unknown weak acid's identity via its Ka. Because the pH at the halfway point of a weak acid titration equals that acid's pKa, researchers can use titration data to experimentally determine an unknown weak acid's Ka value, which can then help identify the unknown compound by comparing it against known reference values.
C
Environmental water testing. Titration is used to measure the acidity or alkalinity (buffering capacity) of natural water samples, an important measurement in assessing water quality and its ability to resist pH changes from pollutants or acid rain.
D
Standardizing a titrant solution before use. Because a titration's accuracy depends entirely on precisely knowing the titrant's concentration, chemists first perform a standardization titration — titrating the prepared titrant solution against a precisely weighed primary standard (a highly pure, stable reference compound) — to confirm the titrant's exact concentration before using it to titrate an actual unknown sample.
⚠️ Most Common Titration Mistakes
The equivalence point is NOT always pH 7 — this is the single most tested and most commonly missed concept in the entire titration topic. Students frequently assume every titration's equivalence point lands at neutral pH. Only the strong acid/strong base case actually produces pH 7 at the equivalence point; weak acid/strong base and strong acid/weak base titrations produce equivalence points above or below 7 respectively, due to hydrolysis of the resulting conjugate base or conjugate acid.
The endpoint (where the indicator changes color) and the equivalence point (defined by stoichiometry) are related but not identical. Students sometimes use these terms interchangeably as if they mean exactly the same thing. The equivalence point is a theoretical, stoichiometrically defined point; the endpoint is the practical, visually observed point where the indicator changes color, and a well-chosen indicator will have its endpoint occur very close to, but not necessarily perfectly at, the true equivalence point.
The midpoint of a weak acid titration (halfway to the equivalence point) is where pH = pKa — not the equivalence point itself. Students sometimes confuse the halfway point (a buffering region where pH equals pKa) with the equivalence point (a separate point further along the curve, where the acid has been fully converted to its conjugate base).
✓ Quick Self-Test
1. What is the equivalence point of a titration, defined precisely?
2. Why does a strong acid/strong base titration have its equivalence point at exactly pH 7, while a weak acid/strong base titration does not?
3. What is the difference between the endpoint and the equivalence point of a titration?
4. What pH value does the midpoint (halfway to the equivalence point) of a weak acid titration correspond to, and why is that useful?
5. If 40.0 mL of an unknown NaOH solution is titrated to the equivalence point by 20.0 mL of 0.200 M HCl, what is the concentration of the unknown NaOH solution?
Answers:
1. The equivalence point is the point in a titration at which the moles of acid exactly equal the moles of base that have reacted, based on the balanced neutralization equation — it is defined by stoichiometry, not by a specific pH value.
2. In a strong acid/strong base titration, the salt formed at the equivalence point (for example, NaCl from HCl + NaOH) has ions that do not undergo any meaningful hydrolysis with water, so the solution remains neutral (pH 7). In a weak acid/strong base titration, the weak acid is fully converted into its conjugate base at the equivalence point, and that conjugate base does hydrolyze water, producing a basic solution (pH above 7).
3. The equivalence point is a theoretical point defined by stoichiometry (moles of acid = moles of base). The endpoint is the practical, observed point during the actual titration where a chemical indicator changes color — a well-chosen indicator's endpoint occurs very close to the true equivalence point, but the two are conceptually distinct.
4. The midpoint (halfway to the equivalence point) of a weak acid titration corresponds to pH = pKa of the weak acid being titrated. This is useful because it allows the weak acid's Ka to be determined experimentally directly from titration data, without needing separate equilibrium measurements.
5. At the equivalence point, moles of HCl = moles of NaOH. Moles of HCl = 0.200 M × 0.0200 L = 0.00400 mol. Since this equals the moles of NaOH originally present, [NaOH] = 0.00400 mol / 0.0400 L = 0.100 M.