1. Unit 2 Stoichiometry – Volumetric analysis
Advanced Higher
Unit 2
Stoichiometry – Volumetric analysis

2. Stoichiometry
Any reaction in which the substances react completely according to the mole ratios given by a balanced (stoichiometric) equation is called a quantitative reaction.
When a quantitative reaction takes place, if the quantity of one substance is known then another the unknown quantity of another can be determined.

3. Two chemical methods of quantitative analysis are
1. Volumetric analysis
2. Gravimetric analysis

4. Volumetric analysis
Volumetric analysis involves using a solution of accurately known concentration to determine the concentration of the other solution involved in the reaction. This is called a titration.

5. There are three main types of titration:
1. acid/base
2. complexometric
3. redox
To carry out a titration, one solution must be pre-prepared accurately using a standard flask and all volumes must be measured using a pipette or burette

6. Standard solutions
The pre-prepared solution of accurately known concentration is known as a standard solution.
A standard solution can only be prepared from a primary standard which must be readily available and have the following characteristics:
- a high purity
- be stable in air and in solution
- a reasonably high formula mass
- be readily soluble (usually in water)

7. The first three characteristics ensure that the sample is weighed out accurately.
The most important characteristic is the primary standard’s stability in air e.g. NaOH cannot be used because it absorbs both moisture and CO2 from the air.

8. Preparing standard solutions
Calculate the mass of solute required to make concentration of solution required.
Dissolve in small amount of deionised water.
Pour into standard flask through funnel.
Wash beaker into standard flask using wash bottle.
Make up to the line with deionised water.
Mix thoroughly by inverting flask.

9. Common primary standards
Acid/base titrations – oxalic acid and anhydrous (dry) sodium carbonate
Complexometric – EDTA
Redox – potassium iodate.

10. Equivalence point
Once a standard solution has been prepared it is titrated with a known volume of solution of unknown concentration.
The point at which the reaction is just complete is called the equivalence point.
An equivalence point must be observed either by a colour change in the reaction or by adding an indicator that changes colour at the equivalence point.
The colour change occurs when the end-point of the titration has been reached.

11. End Point of Titration
Acid/base : an indicator with a relevant pH range must be used to determine the end-point of the reaction e.g. phenolphthalein or methyl orange.

12. Complexometric : the colour changes at the end-point due to the formation of a transition metal complex. EDTA complexes with metal ions in a one to one ratio.
The indicator has to complex with the metal ion to give a visible colour that is different to the uncomplexed indicator. The indicator must also bond to the metal ion less well than EDTA so that when EDTA is added the indicator is displaced.
A colour change is observed when all the indicator is displaced.
Murexide is an excellent indicator for calcium and nickel ions.

13. Redox : potassium permanganate [potassium manganate (VIII)] is used as it acts as its own indicator.
It is decolourised in a redox reaction and therefore, the end-point occurs when a very pale pink colour due to an excess of manganate (VII) is observed.
One problem with this indicator is that the bottom of the meniscus in the burette is difficult to read because of the very dark colour of the solution. This is overcome by reading the scale at the top of the meniscus.

14. Titrations
Set up burette vertically and rinse with standard solution. Fill and record the top reading.
Rinse the pipette with the solution of unknown concentration. Fill to the graduation mark and add a known volume of the unknown to a conical flask.

15. Add a few drops of indicator to the conical flask.
Place the conical flask on a white tile to make the end-point easier to see.
Run the standard solution from the burette into the unknown while continually swirling the flask.
The burette jet must be just inside the flask to avoid missing the flask during swirling.

16. As the end-point is approached the solution is added drop-wise until the indicator shows a permanent colour change.
The reading on the burette is recorded again.
The titre is the second reading minus the first reading.

17. Normally one rough titration and two accurate titrations (within 0
Normally one rough titration and two accurate titrations (within 0.1cm3 of each other) are taken.
The two accurate titrations are then averaged to give the value used in the calculation to find the concentration of the unknown.

18. Volumetric Calculations
Example
10.0 cm3 of lithium hydroxide solution was neutralised by 16.7 cm3 of 0.1 mol l-1 phosphoric acid solution.
Calculate the concentration of the lithium hydroxide solution.
3LiOH + H3PO4 → Li3PO H2O

19. Using the equation:
ca x va = cb x vb
a b
0.1 x = Cb x 0.010
Cb = 3 x = mol l-1
0.010

20. Exercise Now try the exercise on page 6 of the Unit 2(a) booklet.
You may have to revise adding ion-electron equations together from the Higher course first.