Titration calculations

With all titrations, it’s not just about the swirling and colour change — the real skill lies in the calculations. Learn to titrate and calculate the unknown concentration like a pro. #Chemistry #TitrationSkills
 

Mastering Titration: The Art of Measuring with Precision

Titration is one of those essential practical skills every chemistry student learns — often accompanied by a lot of swirling, careful drop-counting, and the occasional sigh of frustration when you overshoot the end point by just one drop.

But titration isn't just about handling a burette with a steady hand. To complete the picture, you also need to master the titration calculations — turning that colour change into meaningful data. In this blog, we’ll explore both the practical and mathematical sides of titration.

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What is Titration?

Titration is a laboratory technique used to determine the concentration of an unknown solution by reacting it with a solution of known concentration.

The most common example is an acid-base titration. Here’s the basic idea:

• You have an acid of known concentration in a burette.

• You place a base (or vice versa) of unknown concentration in a flask.

• You slowly add one to the other until neutralisation occurs — this is known as the end point, usually signalled by a colour change thanks to an indicator like phenolphthalein or methyl orange.

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Equipment You’ll Need

• Burette (to deliver the titrant)

• Pipette (to measure a fixed volume of the analyte)

• Conical flask

• White tile (makes the colour change easier to spot)

• Indicator (to signal the end point)

• Clamp stand and funnel

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The Practical Steps

1. Rinse and fill the burette with the solution of known concentration.

2. Pipette a fixed volume of the unknown solution into the conical flask.

3. Add a few drops of indicator.

4. Titrate by adding the known solution from the burette slowly while swirling the flask until the indicator changes colour.

5. Record the volume used — this is your titre.

6. Repeat for concordant results (within 0.1 cm³ of each other).

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The Titration Calculations

Once you have your average titre (volume added), you can work out the unknown concentration using this key equation:

Moles = Concentration × Volume

Important: Volumes must be in dm³ (1 dm³ = 1000 cm³).

Step-by-Step Example:

Let’s say:

• You titrated 25.0 cm³ of sodium hydroxide (NaOH).

• You used 22.4 cm³ of 0.100 mol/dm³ hydrochloric acid (HCl) to neutralise it.

• The balanced equation is:
  HCl + NaOH → NaCl + H₂O

Step 1: Calculate moles of HCl

$$\text{Moles of HCl} = 0.100 \, \text{mol/dm}^3 \times \frac{22.4}{1000} = 0.00224 \, \text{mol}$$

Step 2: Use the mole ratio

From the equation, HCl and NaOH react 1:1.
So, moles of NaOH = 0.00224 mol

Step 3: Calculate concentration of NaOH

$$\text{Concentration} = \frac{\text{Moles}}{\text{Volume (in dm}^3)} = \frac{0.00224}{25.0/1000} = 0.0896 \, \text{mol/dm}^3$$

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Common Pitfalls

• Not converting cm³ to dm³ — always divide by 1000.

• Forgetting mole ratios — they’re essential if the reaction isn’t 1:1.

• Overshooting the end point — go drop by drop near the colour change.

• Using the first (rough) titre in your average — only include concordant results in your final average.

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Final Thoughts

Titration is both a science and an art — a delicate balance between careful experimental work and sharp calculation. Master both, and you'll not only impress your examiner but also gain a deeper understanding of how chemists measure things with such precision.

And remember: it’s not just about the pretty pink flash in the flask — it’s what you do with those numbers afterwards that really counts.