🧪 pH Titration — Acid-Base ChemistryChemistry🇺🇦 Українська
Titration System
Reaction type
Analyte concentration (M)
0.10 M
Titrant concentration (M)
0.10 M
Volume of analyte (mL)
25 mL
Titrant Added
0.0 mL
Indicator
Current State
pH—
mL added0.0
Equivalence at—
pKa / pKb—
Buffer region—
% Neutralised0%
Presets
Strong acid/strong base: steep sigmoidal curve,
equivalence at pH 7. Weak acid/strong base: half-equivalence at pH = pKa
(Henderson-Hasselbalch). Buffer region ±1 pH unit around pKa. Phenolphthalein changes pH 8.2–10.
Litmus pH 6–8. Methyl orange pH
3.1–4.4.
Titration beaker
pH titration curve — mL of titrant added vs pH
About Acid-Base Titration
Acid-base titration is the process of adding a solution of known concentration (the titrant) to a solution of unknown concentration (the analyte) until the chemical reaction between them is complete. The endpoint is detected by a sharp change in pH near the equivalence point — the point at which stoichiometric amounts of acid and base have been mixed. This simulator calculates pH using exact equilibrium equations for strong acid/strong base, weak acid/strong base (Henderson-Hasselbalch in the buffer region), and weak base/strong acid systems.
Add titrant drop by drop or use the auto-fill feature to trace the full titration curve. The beaker changes colour according to the selected indicator, and the graph plots pH against volume added — making it easy to identify the equivalence point, the half-equivalence point (where pH = pKa for a weak acid), and the flat buffer region either side of it.
Frequently Asked Questions
What causes the sharp jump in pH at the equivalence point?
For a strong acid/strong base titration, the solution transitions from having an excess of H⁺ ions to an excess of OH⁻ ions within a fraction of a millilitre near the equivalence point. Because pH is a logarithmic scale (pH = −log[H⁺]), even a tiny excess of base can shift the pH from 3 to 11 — a change of eight pH units from adding less than 0.1 mL of titrant.
What is the Henderson-Hasselbalch equation?
The Henderson-Hasselbalch equation relates pH to the pKa of a weak acid and the ratio of its conjugate base to acid concentrations: pH = pKa + log([A⁻]/[HA]). In a buffer region this ratio changes slowly, so pH changes slowly too. At the half-equivalence point exactly half the acid has been neutralised, [A⁻] = [HA], and pH = pKa — a handy way to measure a weak acid's pKa experimentally.
Why is the equivalence point of a weak acid / strong base titration above pH 7?
At the equivalence point, all the weak acid (e.g., acetic acid, pKa = 4.76) has been converted to its conjugate base (acetate). Acetate is itself a weak base that partially hydrolyses water: CH₃COO⁻ + H₂O ⇌ CH₃COOH + OH⁻. This produces a slight excess of OH⁻, making the solution basic. For acetic acid / NaOH, the equivalence point pH is typically around 8.9.
Which indicator should I choose for a strong acid / strong base titration?
Any indicator whose colour-change range falls within the steep pH jump (roughly pH 4–10 for HCl/NaOH) will work. Phenolphthalein (pH 8.2–10) and litmus (pH 6–8) are both suitable. For weak acid / strong base titrations the equivalence point is above pH 7, so phenolphthalein (which turns pink above pH 8.2) is the preferred choice.
What makes a good buffer, and what is its capacity?
A buffer works best when pH ≈ pKa, meaning the weak acid and its conjugate base are present in roughly equal amounts. Buffer capacity — the moles of strong acid or base a litre of buffer can absorb before the pH shifts by 1 unit — is greatest at pH = pKa and falls off sharply more than 1 pH unit away. This is why physiological buffers (blood pH 7.4, bicarbonate pKa ≈ 6.1) operate at the edge of their effective range, using CO₂ exhalation as a secondary control.
How does a pH indicator actually work?
A pH indicator is itself a weak acid (HIn) whose protonated and deprotonated forms (In⁻) absorb different wavelengths of light. At low pH, [HIn] dominates and the solution has one colour; at high pH, [In⁻] dominates and the colour changes. The transition midpoint occurs at pH ≈ pKa(HIn). Phenolphthalein has pKa ≈ 9.1, so it changes colour at around pH 8.2–10.
What is the difference between the endpoint and the equivalence point?
The equivalence point is the theoretical point at which moles of acid equal moles of base — it is a calculated quantity. The endpoint is the experimentally observed colour change of the indicator. Because the indicator's pKa may not exactly match the equivalence-point pH, there is always a small titration error. Choosing an indicator whose range brackets the equivalence-point pH minimises this error.
Can I use this simulation to find the pKa of an unknown acid?
Yes — in practice, perform a titration and find the half-equivalence volume (half the volume needed to reach the equivalence point). At that volume, exactly half the acid has been neutralised and the Henderson-Hasselbalch equation gives pH = pKa directly. In this simulator, look for the labelled "pH = pKa" marker on the titration curve for weak acid systems.
Why does the simulation use millimoles instead of moles?
Analytical titrations are carried out on the millilitre scale in laboratory glassware, so concentrations in mol/L (M) and volumes in mL give amounts in millimoles (mmol). For example, 25 mL of 0.1 M acetic acid contains 2.5 mmol — a conveniently sized quantity. The simulator computes n_analyte = C_A × V_A and n_titrant = C_B × V_B in mmol throughout.
What happens beyond the equivalence point?
After the equivalence point, the titrant is in excess. For a strong base titrant, excess OH⁻ determines the pH: pH = 14 + log([OH⁻]_excess). Each additional millilitre of base raises the pH steeply at first, then more slowly as the logarithmic scale compresses. The curve flattens into a gently rising plateau far beyond the equivalence point as the solution simply becomes more alkaline.
About this simulation
This simulator plots a real acid–base titration curve by solving pH from equilibrium chemistry rather than approximating it. Depending on the Reaction type you pick, it uses the strong-acid/strong-base excess-ion equation, or the Henderson–Hasselbalch equation in the buffer region for a weak acid or weak base, switching automatically to hydrolysis maths right at the equivalence point. A beaker canvas and a live pH-vs-volume graph update together as you add titrant, so you can watch the exact moment the colour changes.
🔬 What it shows
A colour-changing beaker and a pH curve that plots mL of titrant added against pH from 0 to 14. The graph marks the equivalence point (yellow dot), shades the buffer region ±1 pH unit either side of the pKa for weak-acid/base systems, and tints a band showing your chosen indicator's colour-change range.
🎮 How to use
Pick a Reaction type (strong/strong, weak acid/strong base, or weak base/strong acid), set Analyte and Titrant concentration and Volume of analyte with the sliders, then add titrant with + Drop, the Titrant Added slider, or ▶ Auto-fill. Switch indicator with the Phenolphthalein / Litmus / BTB / Methyl Orange buttons, or jump straight to a scenario with the Presets list.
💡 Did you know?
Phenolphthalein is colourless below pH 8.2 because its acidic form barely absorbs visible light, but above pH 10 its fully deprotonated form becomes strongly conjugated and turns bright pink — the same molecule is also a mild laxative and was once an ingredient in over-the-counter medicines.
Frequently asked questions
Why does the pH axis jump so steeply near the equivalence point?
Because pH is a logarithmic measure of hydrogen-ion concentration, a tiny change in the moles of excess acid or base near the equivalence point produces a huge swing in pH. For the strong acid/strong base preset, adding less than 0.1 mL of titrant either side of the 25 mL equivalence point can move the reading from pH 4 to pH 10.
What is happening in the buffer region on the graph?
Between roughly zero mL and the equivalence volume, the weak acid/base system contains a mix of the weak acid and its conjugate base, and pH follows pH = pKa + log([base]/[acid]). Because this ratio changes only gradually as titrant is added, the curve flattens out — this flat stretch is the buffer region, centred on the point where pH equals pKa.
Why is the equivalence point above pH 7 for the acetic acid preset?
At the equivalence point every molecule of acetic acid has been converted into acetate, which is itself a weak base and reacts slightly with water to release hydroxide ions. That leftover OH⁻ pushes the solution to about pH 8.9 rather than neutral pH 7, which is why the simulator switches to a hydrolysis calculation exactly at the equivalence volume.
How does the simulator decide the beaker colour?
Each indicator has its own colour-response function keyed to pH: phenolphthalein fades from clear to pink between pH 8.2 and 10, litmus shifts red-to-blue between pH 6 and 8, bromothymol blue (BTB) shifts yellow-to-blue between pH 6 and 7.6, and methyl orange shifts red-to-orange between pH 3.1 and 4.4. The beaker canvas recomputes the fill colour every time pH updates.
What do the bubbles near the equivalence point mean?
They are a visual cue, not a real chemical effect — the simulator draws small bubble marks in the beaker whenever the titrant volume is within about 15% of the calculated equivalence volume, to help you notice you are approaching the steep part of the curve before the colour change happens.