Titration Lab
Add base from a virtual burette to an acid solution and watch the pH change in real time. Find the equivalence point from the sigmoidal pH curve, observe indicator color changes, and explore the difference between strong and weak acid titrations.
Guided Experiment: Finding the Equivalence Point
If you add base to an acid solution, at what volume of base do you predict the pH will equal 7? How will the pH curve look near that point?
Write your hypothesis in the Lab Report panel, then click Next.
Controls
Results
pH vs Volume Added
The red dot shows the current pH. The dashed vertical line marks the equivalence point.
Data Table
(0 rows)| # | Trial | Base Added(mL) | pH | Indicator Color |
|---|
Reference Guide
Acid-Base Titration
Titration is a technique to find the concentration of an unknown solution by reacting it with a standard solution of known concentration.
A base (NaOH) is added drop by drop from a burette into an acid (HCl or acetic acid) until the moles of acid equal the moles of base — the equivalence point.
pH Curve Shape
Strong acid vs. strong base: the pH stays low until near the equivalence point, then jumps sharply from about 3 to 11 over just a few drops.
Weak acid vs. strong base: the jump is less steep due to the buffer region. The pH starts higher and the equivalence point occurs above pH 7.
Henderson-Hasselbalch Equation
In the buffer region of a weak acid titration, the pH is governed by the Henderson-Hasselbalch equation.
At the half-equivalence point, [A-] = [HA], so the log term = 0 and pH = pKa. This makes it easy to read the pKa directly from the titration curve.
Indicators
Indicators are weak acids that change color when they lose or gain a proton. Choose an indicator whose transition range spans the equivalence point.
| Indicator | Range | Colors |
|---|---|---|
| Methyl Orange | 3.1 – 4.4 | Red → Yellow |
| Bromothymol Blue | 6.0 – 7.6 | Yellow → Blue |
| Phenolphthalein | 8.2 – 10.0 | Colorless → Pink |