Volumetric analysis

The technique of volumetric analysis is used to determine the concentration of an unknown solution.

This is achieved by reacting a measured volume of a solution of known concentration (called a standard solution,

which is often prepared from a primary standard) with a measured volume of the solution of unknown solution. The

solutions are mixed in a titration until they have just reacted and no reactants remain in excess.

Major pieces of equipment used in a titration:

Definitions

Molarity the concentration of a solution expressed in mol.L-1 (M)

Burette used to deliver variable volumes accurately

Titre the volume delivered from the burette

Pipette used for accurate transfer of fixed volumes e.g. 20.00 mL

Aliquot the known volume delivered from the pipette

Titrationthe process of mixing the reactants in a volumetric analysis until they are present in the exact mole ratio represented in the balanced equation for the reaction. This is a procedure for accurately determining the concentration of an unknown solution

Standard solution a solution of accurately known concentration

Primary standard a substance which can be used to make up a standard solution

funnel

conical flask

burette volumetric flask

pipette

1. Preparing a standard solution

In order to make up a standard solution, a primary standard can be used. This is a chemical that must have the

following properties:

1. Be readily available in pure form

2. Have a known formula (so molar mass can be calculated)

3. Be easy to store without deteriorating or reacting with the atmosphere

The primary standard must be accurately weighed and made up to a known volume in a standard flask. This allows

the molarity to be accurately calculated. The solution that has been prepared is the standard solution.

Step 1: Adding the solid

Weigh solids plus weighing bottle and lid.

Transfer solids to a 100mL beaker and dissolve in a small volume of distilled water.

Reweigh the empty weighing bottle and lid. The mass difference equals the mass of the crystals added to the

standard flask.

Step 2: Adding most of the water

Transfer dissolved solids to flask through the funnel. Thoroughly rinse the beaker with distilled water, adding the

rinses to the flask.

Add de-ionised water to the flask through the funnel until the flask is about ¼ full.

Swirl the flask to dissolve the solute.

Keeping adding water to the flask, but only fill to about 1 cm below the etched mark.

Step 3: Making up to volume

Add the final amount of water using a dropping pipette, until the bottom of the meniscus

sits on the etched line.

Add the lid and invert the flask to ensure that the solution has an even distribution of

solute

The standard solution is now ready for use. Some of the solution should be poured into a

clean dry beaker for measuring out an aliquot by pipette.

Once a standard solution has been prepared (or one has been supplied), the titration can begin.

2. Use of the burette

1. Rinse with distilled water then rinse again with the solution with which it is to

be filled.

2. Fill the burette with the solution and run a little solution out into the beaker to

remove air bubbles (remember to remove the funnel).

3. Read the burette down from the top to two decimal places (estimate the second

decimal place). This is the initial reading.

3. Use of the pipette

1. Rinse with distilled water then with the solution with which it is to be filled.

2. Place the tip of the pipette well below the surface of a fresh sample of

the solution.

3. Draw up the solution until the bottom of the meniscus sits on the

etched line.

4. Deliver the aliquot into a conical flask that has been rinsed with

distilled water. Do not try and deliver the last drop of solution from

the pipette. The volume of this drop has been taken into account.

A titration is finished when just enough of one chemical has been added to completely react with the other. When

this occurs, the mole ratios according to the balanced equation will be correct, and the equivalence point has been

reached (neither reactant is in excess).

How can we determine when the equivalence point has been reached?

Indicators are chemicals that change colour when the chemical environment changes. In the case of acid base

reactions, indicators change colour when the pH of the solution changes. Therefore, they can be used to ‘see’ the

equivalence point. However, because indicators required a slight excess of one reactant before they change

colour, they only give an approximation to the equivalence point. This is known as the end point of the titration.

4. Add indicator to the solution in the conical flask.

5. Titrate until the end point has been reached. When you get close to the end point, add the solution drop wise

from the burette and continually swirl the flask. Use a white tile to help observe the colour change. Stop titrating

when the colour change is permanent for 30 seconds.

6. Read the burette to two decimal places. This is the final reading. Repeat the titration procedure until you have

three concordant results . Concordant results are usually within 0.1 mL of each other, although this may vary for

some experiments.

Summary

In volumetric analysis, an accurately measured volume of one solution is placed in a conical flask by means of a

pipette. The second solution is run into the flask from the burette, until the chemicals have just reacted with each

other. This is typically determined by the use of an indicator. This procedure is called a titration.

The indicator shows the end point of the reaction. Typically one reactant is very slightly in excess. Ideally the end

point is determined so that it is as close as possible to the equivalence point, which is the point where neither

reactant is in excess.

Mass of A

Moles of A

Mass of B

Moles of BChemical equationMole ratio

Molar mass of A

Molar mass of B

Mmn A

known of coeff.unknown of coeff.

Mnm B

ReminderSummary: Mass–mass stoichiometry1. Balance the equation for the reaction.2. Convert the known masses of the substances to moles.3. Use the balanced equation to set up the appropriate mole ratios.4. Use the appropriate mole ratios to calculate the number of moles of the desired reactant or product.5. Convert from moles back to grams if required by the problem.

Question

A student performed an experiment to determine the concentration of a sodium hydroxide solution. She was given a

standard solution of 0.1048 M hydrochloric acid to use in the titration. She used a volume of 20.00 mL of sodium

hydroxide, and she achieved the following results:

1 2 3 4 5

Final volume (mL) 22.70 42.84 21.18 41.48 21.10

Initial volume (mL) 2.08 22.70 1.00 21.18 0.86

Titre (mL) 20.62 20.14 20.18 20.30 20.24

Calculate the concentration of the sodium hydroxide solution.

Step 1: Write a balanced equation

Step 2: Determine the average titre using 3 concordant results.

Step 3: Calculate the average amount (in mol) of HCl used.

Step 4: Calculate the amount (in mol) of NaOH required to neutralise the HCl (from the balanced equation).

Step 5: Calculate the molar concentration of the NaOH solution.

Question involving a dilution

A 50.0 mL sample of ethanoic acid (CH3COOH) was diluted to 250.0 mL in a volumetric flask. A 20.00 mL aliquot of this solution required the addition of 27.98 mL of 0.134 M sodium hydroxide solution in order to be neutralised.

a. Write an equation for the neutralisation reaction.

b. What is the molarity of the diluted ethanoic acid?

c. What is the molarity of the original undiluted ethanoic acid?

Questions – refer to Ch 20 of your text