Acid Base Titration Explained Step by Step

Acid-base titration is a laboratory method used to determine the concentration of an unknown acid or base by reacting it with a solution of known concentration (the titrant) until the reaction is complete. You add the titrant gradually from a burette to a measured volume of the unknown solution, and the point at which the acid and base are exactly neutralized is called the equivalence point. The concentration of the unknown is then calculated from the volume of titrant used and the stoichiometry of the reaction.

Formula

The core relationship in acid-base titration comes from the balanced chemical equation. For a monoprotic acid and a monoprotic base (like HCl and NaOH), the reaction is 1:1, so the formula is:

C_aV_a = C_bV_b

• C_a = concentration of the acid (mol/L)
• V_a = volume of the acid solution (mL or L)
• C_b = concentration of the base (mol/L)
• V_b = volume of the base solution (mL or L)

If the acid and base react in a different ratio (e.g., H₂SO₄ with NaOH is 1:2), you must include the mole ratio from the balanced equation.

Worked Example

Suppose you titrate 25.0 mL of hydrochloric acid (HCl) with 0.100 mol/L sodium hydroxide (NaOH). It takes 20.0 mL of NaOH to reach the equivalence point. What is the concentration of the HCl?

1. Write the balanced equation: HCl + NaOH → NaCl + H₂O (1:1 ratio)
2. Use the formula: C_aV_a = C_bV_b
3. Plug in the values: C_a × 25.0 mL = 0.100 mol/L × 20.0 mL
4. Solve: C_a = (0.100 × 20.0) / 25.0 = 0.0800 mol/L

The concentration of the HCl is 0.080 mol/L. The calculator above performs the same calculation using your inputs.

What You Need for a Titration

Performing an acid-base titration requires specific glassware and reagents. The essential items are:

• Burette – a graduated tube that delivers the titrant drop by drop, allowing precise volume measurement.
• Pipette – used to transfer a known volume of the unknown solution into a flask.
• Erlenmeyer flask – holds the unknown solution during the titration.
• Indicator – a substance that changes color at or near the equivalence point (e.g., phenolphthalein, methyl orange).
• Titrant – a solution of known concentration, usually a strong acid or base.
• Analyte – the solution of unknown concentration.

The choice of indicator is critical – it must change color at the pH of the equivalence point, not before or after.

How to Perform a Titration Step by Step

1. Rinse the burette with the titrant and fill it, making sure there are no air bubbles at the tip.
2. Use a pipette to transfer a precise volume of the unknown solution into the flask.
3. Add 2–3 drops of indicator to the flask.
4. Record the initial volume of titrant in the burette.
5. Slowly add titrant while swirling the flask. Near the endpoint, add drop by drop.
6. Stop when the indicator color changes permanently (the endpoint).
7. Record the final volume and calculate the volume of titrant used.
8. Repeat the titration at least three times and average the results for accuracy.

Choosing the Right Indicator

The equivalence point of a titration is not always at pH 7. The pH at equivalence depends on the strength of the acid and base involved. Indicators are weak acids or bases themselves that change color over a specific pH range. The correct indicator is one whose color-change range overlaps the steep part of the titration curve (the near-vertical section).

• Strong acid + strong base: equivalence at pH 7 – phenolphthalein or bromothymol blue works.
• Weak acid + strong base: equivalence above pH 7 – phenolphthalein (pH 8.2–10.0) is suitable.
• Strong acid + weak base: equivalence below pH 7 – methyl orange (pH 3.1–4.4) is suitable.
• Weak acid + weak base: no sharp pH change – titration is not practical with a simple indicator; use a pH meter.

Common Mistakes and How to Avoid Them

Several errors can ruin titration results. The most frequent are:

• Using the wrong indicator – leads to an endpoint that does not match the equivalence point.
• Not rinsing the burette with titrant – residual water dilutes the titrant, causing errors.
• Reading the burette incorrectly – always read at eye level, from the bottom of the meniscus.
• Adding titrant too quickly – overshooting the endpoint forces you to start over.
• Forgetting to account for the mole ratio – using C_aV_a = C_bV_b when the reaction is not 1:1 gives a wrong answer.
• Not using a blank or correcting for indicator volume – the indicator itself consumes a tiny amount of titrant, but this is usually negligible.

Frequently Asked Questions

What is the difference between endpoint and equivalence point?
The equivalence point is the exact moment when the moles of acid equal the moles of base according to the balanced equation. The endpoint is when the indicator changes color. Ideally they are the same, but a poorly chosen indicator can make the endpoint occur before or after the true equivalence point.

Why do I need to standardize the titrant?
Standardization is the process of determining the exact concentration of a titrant solution by titrating it against a primary standard (a highly pure substance of known mass, like potassium hydrogen phthalate). This is necessary because many solutions absorb moisture or degrade over time, so their actual concentration may differ from the label.

Can I titrate a weak acid with a weak base?
It is possible in theory, but the pH change at the equivalence point is so gradual that no simple indicator can detect it accurately. For such titrations, a pH meter is used to monitor the pH continuously and plot a titration curve.

What is a titration curve?
A titration curve is a graph of pH (y-axis) versus volume of titrant added (x-axis). It shows how the pH changes slowly at first, then jumps sharply near the equivalence point, and then levels off. The midpoint of the steep vertical section is the equivalence point.

Why do I need to do multiple trials?
Repeating the titration and averaging the results reduces random errors such as reading the burette slightly off or adding a drop too much. A typical good titration set includes three trials that agree within 0.1 mL of each other.

Can I use this method for real-world samples like vinegar or aspirin?
Yes. Vinegar contains acetic acid, which can be titrated with a strong base. Aspirin contains acetylsalicylic acid, which can be titrated with a strong base after dissolving it in a suitable solvent. The same principles apply, but the sample may need preparation to isolate the acid of interest.

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in v | Volume of acid (mL) | mL | 25.0 | 0
in c | Concentration of base (mol/L) | mol/L | 0.100 | 0
in vb | Volume of base used (mL) | mL | 20.0 | 0
out ca | Concentration of acid (mol/L) | mol/L | (c * vb) / v | 3