Acid–Base Titration and the Titration Curve
ChemistryAcids, Bases & TitrationAges 17–18Virtual lab
Use it with your class: students open it from a link or a QR code, with no account needed.
Next in Acids, Bases & Titration
Add NaOH solution drop by drop from a burette into the acid in an Erlenmeyer flask. The pH-vs-volume graph updates with every drop, showing the sharp pH jump at the equivalence point. Change the strong/weak acid, the concentrations, and the indicator to see why titrating a weak acid needs phenolphthalein rather than methyl orange. A pH-probe lab view adds a weak base (NH₃ with HCl), a weak acid with a weak base, and polyprotic H₃PO₄, citric acid and Na₂CO₃: calibrate the probe with buffers, record a results table, find end points from ΔpH/ΔV and read pKa at half-equivalence. A primary-standard practical makes up oxalic acid or sodium carbonate in a 250 mL volumetric flask and titrates NaOH or HCl to concordant titres to find its molarity and strength in g/L.
Lesson: Acid–base titration practical
What it shows
In an acid–base titration, a solution is added from a burette until the equivalence point, where acid and base have reacted in exact proportion. Plotting pH against volume gives a curve with a steep jump there. A strong acid with a strong base jumps around pH 7; a weak acid gives a buffer region, where pH = pKa at half-equivalence, and a basic equivalence point. Polyprotic acids can show several jumps, and a suitable indicator changes color inside the jump. A primary standard, such as oxalic acid or anhydrous sodium carbonate, is weighed and made up to a known volume to find the exact concentration of another solution.
How to use
In Classic titration, choose the Acid and the Indicator, add NaOH with +1 mL and +1 drop, and watch where the color changes. pH-probe lab adds more systems: calibrate the probe with buffers, press Record after each addition and find the end point from ΔpH/ΔV. In Primary standard, set Mass weighed, press Transfer and dissolve, add water to the mark, press Stopper and mix, then use New titration and Record until two titres agree, and press Check.
Parameters you can change
- Acid being titrated HCl (strong acid), CH₃COOH (weak acid, Ka = 1.8×10⁻⁵)
- Concentration of the flask solution 0.01–0.5 M
- Volume of the flask solution 5–50 mL
- Concentration of the burette solution 0.01–0.5 M
- Indicator Phenolphthalein (8.2–10), Methyl orange (3.1–4.4), Methyl red (4.4–6.2), Bromothymol blue (6.0–7.6), Thymolphthalein (9.3–10.5)
- View Classic titration (HCl or CH₃COOH with NaOH), pH-probe lab (more systems, results table), Primary standard practical (make up a standard solution, concordant titres)
- Titration system (lab view) HCl + NaOH (strong acid, strong base), CH₃COOH + NaOH (weak acid, strong base), NH₃ + HCl (weak base, strong acid), CH₃COOH + NH₃ (weak acid, weak base), H₃PO₄ + NaOH (triprotic acid), Citric acid + NaOH (triprotic acid), Na₂CO₃ + HCl (two equivalence points)
- Primary standard (primary-standard practical) Oxalic acid dihydrate, titrated with NaOH, Anhydrous sodium carbonate, titrated with HCl
- Mass of primary standard weighed 0.2–5 g
Questions to explore
- Why is the pH at the equivalence point above 7 when CH₃COOH is titrated with NaOH?
- Why does methyl orange change color too early in the titration of CH₃COOH?
- Why is methyl red a better indicator than phenolphthalein when NH₃ is titrated with HCl?
Common questions
What is the difference between the equivalence point and the end point of a titration?
The equivalence point is where acid and base have reacted in exact proportion, while the end point is where the change is actually seen, usually an indicator's color change. A good indicator changes color inside the pH jump, so the two almost coincide. With a pH probe, the end point is taken where ΔpH/ΔV is largest.
Why does pH equal pKa at half-equivalence?
At half-equivalence, half of the weak acid has been converted into its conjugate base, so [HA] = [A⁻]. The Henderson–Hasselbalch equation, pH = pKa + log([A⁻]/[HA]), then reduces to pH = pKa because log 1 = 0. For CH₃COOH, with Ka = 1.8×10⁻⁵, the pH at this point is about 4.74.
Why can't NaOH be used as a primary standard?
Solid NaOH is not a primary standard because it absorbs water and CO₂ from the air, so a weighed mass does not give an exact amount of NaOH. A primary standard must be pure, stable and easy to weigh, as oxalic acid dihydrate and anhydrous sodium carbonate are. An NaOH solution is therefore standardised by titration against one of them.
Teaching ideas
All guides- Free Online Science Labs for Middle and High SchoolFree online science labs for middle and high school: 28 virtual labs in physics, chemistry, biology and earth science, what students do and the lesson plans.
- Buffer Virtual Lab: Buffer pH and Buffer CapacityA buffer virtual lab: add acid and base to water and to buffers, then show that the base:acid ratio sets the pH and the total concentration sets the capacity.
- Titration Virtual Lab: Reacting Volumes and pH CurvesA titration virtual lab: find the end point with phenolphthalein, calculate an acid concentration, then compare strong and weak acid pH curves and read pKa.
- Titration Unknowns Virtual Lab: Vinegar, Juice and H₂O₂A titration virtual lab with hidden unknowns: find the acid in vinegar and fruit juice with NaOH, and the % H₂O₂ with KMnO₄, from concordant titres.
