Titration Virtual Lab: Reacting Volumes and pH Curves

This titration virtual lab covers two school practicals in one simulation. Part A matches the AQA GCSE required practical on titration: students add sodium hydroxide from a burette to hydrochloric acid until the indicator changes, record the titre and calculate the acid concentration. Part B follows the AP Chemistry lab on how acid strength shapes a titration curve: students titrate a strong and a weak acid, compare the curves, read the pKa at half-equivalence and test two indicators. Every drop updates the pH and the curve, so students see why the end point is so sharp.

Curriculum links

• AQA GCSE Chemistry 4.4.2.5 (titrations) and its required practical on finding the reacting volumes of a strong acid and a strong alkali.
• AP Chemistry Unit 8, topic 8.5 Acid–Base Titrations (SAP-9.E), and the AP Chemistry lab on titration curves of acids with different strengths and concentrations.
• The same titration appears in IB Chemistry and A-level Chemistry.

Simulic is not affiliated with or endorsed by the College Board or AQA.

Before the lab (5 min)

Ask students to commit to a prediction:

"20 mL of 0.10 M HCl and 20 mL of 0.10 M CH₃COOH are each titrated with 0.10 M NaOH. Which one needs more NaOH to reach the equivalence point?"

Expect many votes for HCl ("it's the stronger acid"). Leave the question open.

Method in the simulation

Part A: reacting volumes (GCSE)

1. Set Acid to HCl, C(acid) to 0.08 M, V(acid) to 25 mL, C(NaOH) to 0.10 M and Indicator to phenolphthalein. Pin these on the class link.
2. Rough titration: press +1 mL until the flask turns pink. Note the volume.
3. Press Restart. Add +1 mL until you are 1 mL short of the rough titre, then press +1 drop until the first pink color. Record the burette reading.
4. Calculate: moles of NaOH = 0.10 × titre (in L), moles of HCl = moles of NaOH, concentration of HCl = moles ÷ 0.025 L.

Part B: strong vs weak acid curves (AP)

5. Set HCl, 0.10 M, 20 mL, NaOH 0.10 M, phenolphthalein. Record the pH at each volume in the table.
6. Switch Acid to CH₃COOH and repeat.
7. Switch Indicator to methyl orange for each acid. Read where the simulation says the color changes.

V(NaOH) (mL)	0	5	10	15	19	20	21	25
pH, HCl
pH, CH₃COOH

Expected results

Part A. At 20.00 mL the flask is still colorless at pH 7.00. The next drop (20.05 mL) takes the pH to 10.05 and the flask turns pink. With a titre of 20.05 mL: n(NaOH) = 0.10 × 0.02005 = 0.002005 mol, and c(HCl) = 0.002005 ÷ 0.025 = 0.080 mol/L. Repeat titrations give the same reading, because the simulation has no reading errors.

Part B. Readings from the simulation:

V(NaOH) (mL)	0	5	10	15	19	20	21	25
pH, HCl	1.00	1.22	1.48	1.85	2.59	7.00	11.39	12.05
pH, CH₃COOH	2.88	4.27	4.75	5.22	6.02	8.72	11.39	12.05

• Both acids reach equivalence at 20.00 mL: equal moles of acid need equal moles of NaOH.
• The weak acid starts higher, rises gently through a buffer region, and has its equivalence point at pH 8.72, not 7.
• At half-equivalence (10.00 mL) the weak acid's pH is 4.75, which matches pKa = 4.74.
• Phenolphthalein changes color at 20.00 mL (HCl) and 19.99 mL (CH₃COOH). Methyl orange works for HCl (19.98 mL) but changes at 6.22 mL for CH₃COOH, an error of about −69%.

Questions for students

1. (Prediction, asked again after the lab) Which acid needs more NaOH to reach the equivalence point?
2. In Part B, which is the independent variable?
3. In Part A, what is the concentration of the HCl?
4. What is the pH of the CH₃COOH solution at half-equivalence, and what does it tell you?
5. Why is methyl orange a poor indicator for the CH₃COOH titration?

Answers for teachers: (1) Both the same: 20.00 mL. (2) The type of acid, strong or weak. (3) 0.080 mol/L (accept 0.078–0.082). (4) 4.75, equal to the pKa of ethanoic acid (accept 4.72–4.78). (5) Its color change (pH 3.1–4.4) falls in the buffer region, far before the steep jump, so it changes at about 6 mL instead of 20 mL.

Common misconceptions

• "A strong acid needs more alkali." Strength changes the pH, not the moles of acid. Equal moles need equal volumes.
• "The equivalence point is always at pH 7." Only for a strong acid with a strong base. Ethanoate ions make the weak acid's equivalence solution basic.
• "The end point and the equivalence point are the same thing." The end point is where the indicator changes. It only matches the equivalence point when the indicator's range sits inside the jump.

Extension

• Concentration and the jump: set both concentrations to 0.01 M and compare the size of the pH jump with the 0.10 M curve. Ask which indicator still works.
• Burette limits: set C(acid) to 0.5 M and V(acid) to 50 mL. The simulation warns that the equivalence volume is over 50 mL. Students redesign the titration so the titre is 10–25 mL.

FAQ

Can the acid concentration be hidden as a true unknown?

No. The simulation shows the acid concentration and the equivalence volume. Use Part A for technique and the calculation, and keep real unknowns for the wet lab.

My class is GCSE only. Which questions should I use?

Questions 1, 3 and the Part A method. Replace questions 4 and 5 with your own, for example "Why do we add the NaOH drop by drop near the end point?"

Why does the pH jump from 7.00 to 10.05 in one drop?

At equivalence there is no excess acid or base. One extra drop of 0.10 M NaOH is a large excess of OH⁻ in a solution with almost no H⁺, so the pH rises by about three units.

Related simulations and guides

Strong vs Weak Acids – Same Concentration, What Differs? Acids, bases and indicators – titration drop by drop

For the buffer region in more depth, see the buffer virtual lab. For pH basics, see the pH scale lesson plan. For building quizzes around this simulation, see add simulations to quiz questions.