2.1.4 (d) Making Standard Solutions & Carrying out Acid-Base Titrations

Syllabus

Acid–base titrations

(d) the techniques and procedures used when preparing a standard solution of required concentration and carrying out acid–base titrations.

What does this mean?

In Analytical Chemistry a titration is a reaction used to find information about the concentration or purity of another substance, eg by carefully adding an acid of a known concentration to a measured volume of alkali of unknown concentration, then calculating the moles of acid and using this to infer the moles of alkali in the measured volume.

This would be Quantitative analysis (looking a numbers) rather than Qualitative analysis (looking at changes of colour etc)

Standard Solutions

For the data to be any value it must be compared to a solution that we we have accurately prepared so that we know the concentration precisely.

In general it follows a 5 steps.

1 Weigh the solute

2 Dissolve the solute

3 Make up the solution to a known volume.

4 Homogenise

1. Weighing the solute

There is no point beginning to make a standard solution unless the solute you are about to dissolve is pure.

The solute is weighed on a plastic weighing boat.

This is first weighed full and then again after the solute has been transferred to a beaker so that the difference can be found.

This takes into account any solute that sticks to the weighing boat.

Mass of weighing boat and solute (g)

5.63

Mass of weighing boat after transfer (g)

0.85

Mass of solute transferred (g)

4.78

2 Dissolving the solute

The solute is transferred to a beaker because they are easy to stir.

The solute is dissolved in a small volume of distilled water.

This ensures that you are not adding any unaccounted for substances from tap water.

All the solute must dissolve or the concentration will be incorrect.

If it does not then add a little more distilled water.

Any solute that sticks to the sides of the beaker or to the glass rod can be dealt with providing that the volume of water we use is less than the volumetric flask it will be transferred to.

3 Making the solution up to a known volume.

A volumetric flask is a piece of glassware on which you'll find an engraved mark.

Making the solution up to this mark is more accurate than using a measuring cylinder.

However, it is important to wash the stirring rod you used into you beaker with more distilled water.

Then to wash the beaker into the flask.

This ensures that all the dissolved solute is in the flask.

If a funnel is used for this then the funnel must also be thoroughly washed.

Finally the flask is carefully made up to the mark so that the bottom of the meniscus touches the line as shown.

4. Homogenising the solution.

Currently most of the dissolved solute is in the bottom of the flask.

The top is mostly filled with the distilled water you made up the solution with at the end.

The concentration is not homogeneous - it is inconsistent from bottom to top.

This is no use for a titration.

You could wait for diffusion to equalise everything if you have a few days!

Better to put in the stopper and gently invert the flask a few times (with your thumb keeping the stopper in place).

Carrying out the Titration

Let us assume that we wish to know the concentration of a solution of Sodium Hydroxide (NaOH)

We will need a standard solution of an acid - say, Hydrochloric Acid (HCl)

We would take a 25cm³ volumetric pipette and fill it carefully to the mark with the NaOH solution.

This would be transferred to a clean conical flask.

A suitable indicator we be chosen - in this case Phenolphthalein or Methyl Orange.

The flask is then usually placed on a white tile to make any colour changes clearer.

The acid would be added to a burette.

It is generally better not to put bases in burettes as they are more likely to clag the tap.

It isn't necessary to fill the burette to precisely 0.00cm

Just keep careful note of the starting volume.

Carefully add the acid until the colour change is seen - the end-point.

Note the new volume reading on the burette.

You first attempt will probably not be accurate so you will be doing the titration at least twice more.

Make sure that there is enough acid left in the burette to carry out another titration and repeat until you have two concordant readings - two end-points that are within 0.1 cm³.

Tabulating your results.

This could pick up marks in an exam and is mandatory in PAGs

So, it is worth getting this right.

Your results should look something like the ones right:

Notice that the same number of decimal places is used throughout - even when the burette read 18.3 we record 18.30 to show that we were working to 2 dp.

You are never likely to use your rough titration in the calculation of the mean titre and Titration 2 was simply too different to titration 1 to stop at that point.

So titration 3 was carried out to see whether titration 1 or 2 was most accurate.

Since it is clearly 1 and 3 which are closest these are the only two used.

And the Mean Titre = (18.55 + 18.45)/2 = 18.50 cm³

Now we know the volume of the HCl used. We already knew its concentration.

So we can quickly find the number of moles used by

Moles = Concentration (mol dm^-3) x Volume (dm³)

We would need to think about the reaction.

HCl_(aq)+ NaOH_(aq) → NaCl_(aq) + H₂O_(l)

This allows us to check the stoichiometry.

In this case the mole ratio is 1:1

So however many moles of HCl were used they would have neutralised an equal number of moles of NaOH.

Since the NaOH was precisely 25cm³, (0.025dm³) we can now quickly calculate its concentration by:

Concentration = Moles / Volume (dm³)

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