1. Higher Chemistry Unit 3 – Chemistry in Society Section 18 – Volumetric Analysis

2. Learning Outcomes/Success Criteria Unit 3 Higher Chemistry
3.18 – Volumetric Analysis
Learn that many chemical reactions are reversible
Learning Outcomes/Success Criteria Unit 3 Higher Chemistry
I can carry out acid/base and redox titrations
I can explain the terms standard solution, end point and indicator.
I can use a balanced equation to calculate the quantity of an unknown reactant using information from a titration experiment
I know that redox titrations involving potassium permanganate are self-indicating

3. Titrations
Titration is a technique for measuring the concentration of a solution. A solution of known concentration is used to work out the unknown concentration of another solution.
Redox titrations involve solutions of reducing and oxidising agents.
At equivalence-point of a redox titration precisely enough electrons have been removed to oxidise all of the reducing agent. 3 3

4. Standard Solutions
Standard solutions are solutions of known concentrations. They are made in volumetric glassware and can be made from a solid or from another solution.
Most solutions are aqueous i.e. the solvent is water
When preparing aqueous solutions deionised water (or distilled water) is added why?
Using deonised water in the chemistry laboratory is important because the ions found in water can affect your experiments. Incorporation of these ions, even if in only small amounts, can cause your experiments to receive false results — or to not work at all.

5. Calculations Involving Concentration
How many moles are in 100cm3 of sodium hydroxide concentration 0.1 mol l-1?
0.01 moles

6. Making a Standard Solution
Finally, fill the solvent to the line.
weigh out the solid.
Next, dissolve the solid in a small amount of solvent.

7. Making a Standard Solution
What mass of solid should be dissolved in 100cm3 of water to prepare a 0.1 mol l-1 solution of copper(II) sulphate?
n = C x V/1000
n = 0.1 x 100/1000
n= 0.01 moles
m = n x gfm
m = 0.01 x 159.6
m = 1.596g
1 mole  g
0.01 moles  =(0.01/1)x159.6
=1.596g

8. Calculations involving concentration
How many moles are in 100cm3 of sodium hydroxide concentration 0.1 mol l-1?
0.01 moles

9. Vitamin C - Redox Titration
Vitamin C (ascorbic acid) is an important component of our diet. In its absence the protein, collagen, cannot form fibres properly and this results in skin lesions and blood vessel fragility.
Although vitamin C occurs naturally in many fruits and vegetables many people take vitamin C tablets to supplement their intake.
Vitamin C can undergo a redox reaction with iodine in which the vitamin C is oxidised and the iodine molecules are reduced
I2(aq) + 2e I-(aq) 9

10. Redox Titrations, Vitamin C
I2 (aq) + 2e-  2I - (aq)
C6H8O6 
C6H6O H+ (aq) e-
reduction
oxidation 
Blue/Black (in the presence of starch)
I2 (aq) + C6H8O6
colourless
C6H6O H I- (aq)
Iodine, those concentration is known (in the burette)
acts as an oxidising agent.
Vitamin C, the unknown concentration (in the conical flask)
is a reducing agent.
Starch is added to show when the end-point is reached.

11. Procedure – Making a Standard Solution
1. Use the mortar and pestle to grind up 1 vitamin C tablet. Then add the powdered vitamin C tablet to a medium sized beaker.
2. Add some deionised water (approximately 50 cm3) to the beaker and stir the mixture until the tablet has dissolved.
3. Carefully add the resulting solution to the 250 cm3 standard flask. Rinse out the beaker several times with water and add the washings to the flask.
4. Add water to the standard flask to bring the volume of the solution up to the graduation mark on the neck.
5. Stopper the flask and invert it several times to make sure the solution is thoroughly mixed. 11

12. After rinsing the pipette with a little of the vitamin C solution, pipette 25 cm3 of it into
the conical flask. Add a few drops of starch solution to the vitamin C solution in the
conical flask.
After rinsing the burette with a little iodine solution, fill the burette with the iodine
solution.
Note the initial burette reading. Since the solution has a dark colour, it is difficult to
see the bottom of the meniscus. Take the burette reading from the top of the
meniscus.
Add the iodine solution slowly from the burette whilst gently swirling the solution in the
conical flask. Initially you will see a blue/black colour as the iodine reacts with the
starch but this will rapidly disappear as the iodine reacts with the vitamin C.
Near the end-point of the titration the colour disappears more slowly. At this point add
the iodine solution drop by drop until the solution just turns a blue/black colour and
remains so. This is the end-point of the titration i.e. all the vitamin C has reacted. Note
the final burette reading. Repeat the titrations until concordant results are obtained. 12

13. Redox Titration 1 2 3 4 What to do: z
Carefully fill the burette with iodine solution. 1 z
Carefully pipette exactly 20 ml of vitamin C solution into the conical flask. Add a Couple drops of starch solution. 2 3
Add the iodine until a permanent
blue/black colour appears in the conical flask.
A rough titration is done first to give a rough equivalence-point (end-point), then repeated more accurately to give concordant results (+or-0.1cm3). 4 13 13

14. Average of the First and Second Run titres
Redox Titration
 Results
Trial
First Run
Second Run
1st Burette Reading
cm3
2nd Burette Reading
Iodine titre
Average of the First and Second Run titres

15. Calculation
Knowing the average volume and concentration of the iodine solution used in the redox titration, calculate the number of moles of iodine can be calculated.
With the result from step (a) and the balanced equation for the redox reaction, we can work out the number of moles of vitamin C in 25 cm3 of the vitamin C solution. This can be scaled up to find the number of moles of vitamin C in 250 cm3 of the vitamin C solution.
Your final answer in step (b) will, of course, be equal to the number of moles of vitamin C in the tablet. Using this result and the mass of one mole of vitamin C (176 g) we can finally work out the mass of vitamin C in the tablet. 15

16. Volumetric Titrations
H2SO4 + 2NaOH → Na2SO4 + 2H2O
20 cm3 of sodium hydroxide solution was neutralised by 15cm3 of 0.1mol l-1 sulphuric acid.
Calculate the concentration of the sodium hydroxide solution.
0.15 mol l-1 16

17. Volumetric titrations
What volume of potassium hydroxide, KOH, concentration 2 mol l-1 is
required to neutralise 50cm3 of H2SO4, concentration 0.5 mol l-1 ?
0.05 Litres 17

18. Volumetric Titrations
Examples for you to try
1) What volume of sulphuric acid, concentration 2 mol l-1 is required to neutralise 25cm3 of KOH, concentration 4 mol l-1?
2) If 50cm3 of KOH solution is neutralised by 17.8cm3 of H2SO4 (2 mol l-1), what is the concentration of the alkali?
3) What volume of HCl, concentration 1.0 mol l-1 is required to neutralise 100cm3 of NaOH solution concentration 0.5moll-1?
25 cm3
1.42mol l-1
50cm3 18

19. The Chemistry of Fizzy Sweets
A gelatine sweet incorporates soluble flavourings and sweeteners
A common ingredient is citric acid which gives the sharp, citrus taste.
A role of the analytical chemist is to check that the gum contains the correct amount of citric acid confirming that the plant is working efficiently.
How do chemists check what it contains?

20. The Reaction C6H8O7 + 3NaOH C6H5O7Na3+3H2O
Citric acid is neutralised by sodium hydroxide
Phenolphthalein indicator is colourless in acid but changes to pink once the acid is used up and the alkali is in excess

21. What to do Fill a burette with 0.001 mol l-1 NaOH
Accurately weigh out approximately 2g of jelly sweet.
Add to 75cm3 warm water and stir until the jelly and the soluble citric acid are dissolved.
Transfer this to a 250 cm3 volumetric flask, together with the washings and make up to the mark.
Fill a burette with mol l-1 NaOH
Pipette 20cm3 of the citric acid into a conical flask and add a few drops of phenolphthalein indicator
Titrate against the NaOH until a pink colour is seen. Repeat until concordant results are obtained.

22. Calculation Citric acid : NaOH 1 mol : 3 mol
Calculate the percentage mass of citric acid in the sweet.
Citric acid : NaOH
1 mol : 3 mol

23. Redox Titrations
The same procedures can be used with a balanced redox equation
Permanganate ions react with hydrogen peroxide in acidic solution.
2MnO4- + 6H+ + 5H2O2 → 2Mn2+ + 8H2O + 5O2
25cm3 of hydrogen peroxide solution reacted with 16cm3 of permanganate solution,
concentration 0.1 mol l-1.
Calculate the concentration of the hydrogen peroxide solution.
0.16 mol l-1 23

24. Redox Titrations Try the following example
The overall equation for the reaction of I2 with SO32- ions is:
I2 + SO H2O → 2I- + SO H+
Calculate the volume of iodine solution (concentration 0·5mol 1-1) needed to completely react with 50cm3 of sodium sulphite solution of concentration 0·2mol 1-1.
2. The overall equation for the reaction of Fe3+ ions with I- ions is:
2Fe I- → I Fe2+
Calculate the volume of iodide solution (concentration 0·2mol 1-1) needed to completely react with 25cm3 of iron (III) nitrate solution of concentration 0·1mol 1-1.
20cm3
12.5cm3 24

25. Redox Titrations
Try the following example
a) Write the equation for the overall reaction between acidified dichromate ions and bromide ions.
b) Calculate the volume of bromide solution, concentration 0.5 mol l-1 that would react completely with 50cm3 of 0.2 mol l-1 solution of acidified dichromate.
2Cr2O H+ + 6Br- → 2Cr3+ + 7H2O + 3Br2
120cm3 25

26. Redox Titrations a) Write the ion-electron equation for the
Try the following example
a) Write the ion-electron equation for the
i) reduction of dichromate ions
ii) oxidation of sulphite ions
b) Write the equation for the overall reaction between dichromate ions and sulphite ions.
c) Calculate the volume of dichromate solution 0.5 mol l-1 that would react completely with 30cm3 of 0.25 mol l-1 solution of sulphite ions
2Cr2O H+ + 6e- → 2Cr3+ + 7H2O
SO H2O → SO H+ + 2e-
2Cr2O H+ + 3SO H2O → 2Cr3+ + 7H2O + 3SO H+
5cm3 26

27. Redox Titrations
A 50·0cm3 sample of contaminated water containing chromate ions was titrated and found to require 27·4 cm3 of 0·0200 mol l–1 iron(II) sulphate solution to reach the end-point. The redox equation for the reaction is: 3Fe2+(aq) + CrO42–(aq) + 8H+(aq) → 3Fe3+(aq) + Cr3+(aq) + 4H2O(l) Calculate the chromate ion concentration, in mol l–1, present in the sample of water. Show your working clearly.
Concentration of CrO42- = 0·00365 (mol l-1 ) 27

28. MnO4- + 8H++ 5Fe2+ → Mn2+ + 4H2O + 5Fe3+
Redox Titrations
MnO4- + 8H++ 5Fe → Mn2+ + 4H2O + 5Fe3+
8.25g of an iron (II) salt was dissolved in 250 cm3 of pure water cm3 aliquots were
pipetted from this stock solution and titrated with a standard solution of 0.02 moll-1
acidified potassium permanganate solution.
The titration values obtained were cm3, cm3 and cm3.
Calculate the average titration volume.
(ii) What is meant by a standard solution?
(iii) Calculate the number of moles of potasssium permaganate which reacted with the 25.0cm3 aliquot.
(iv) Calculate the number of moles of iron (II) in the 25cm3 aliquot.
(v) Calculate the number of moles of iron in the original sample.
(vi) Calculate the total mass of iron in the original sample of the iron(II) salt.
(vii) calculate the % iron in the salt.
(viii) Why is the potassium permanganate solution acidified? 28

29. Redox Titrations
Titration is a technique for measuring the concentration of a solution. A solution of known concentration is used to work out the unknown concentration of another solution.
Redox titrations involve solutions of reducing and oxidising agents.
At equivalence-point of a redox titration precisely enough electrons have been removed to oxidise all of the reducing agent.

30. Redox Titrations
Aim : To estimate the iron(II) content of an iron tablet, a tablet is first dissolved in distilled water. This solution is then titrated against previously standardised potassium manganate(VII) solution. The reaction is represented by the equation:
5 Fe2+ (aq) + 8H+ (aq) + MnO4- (aq)  5 Fe3+ (aq) + Mn2+ (aq) + 4H2O(l)
purple
colourless

31. Redox Titrations 1. Find the mass of one iron tablets.
5 Fe2+ (aq) + 8H+ (aq) + MnO4- (aq)  5 Fe3+ (aq) + Mn2+ (aq) + 4H2O(l)
purple
colourless
1. Find the mass of one iron tablets.
2. Crush the tablet in a mortar and pestle. Transfer all the ground material to a beaker where it is dissolved in about 100 cm3 of distilled water.
3. All of this solution (including washings) is transferred to a 250 cm3 volumetric flask and the solution made up to the mark with deionised water. The volumetric flask should be stoppered and inverted several times. This is the solution containing iron(II) ions. 31

32. Redox Titration 1 2 3 4 What to do:
Carefully fill the burette with potassium permanganate 0.001moll-1 2
Carefully pipette exactly 25 ml of iron (II) sulphate into the conical flask. 3
Add the permanganate until a permanent purple colour appears in the conical flask.
A rough titration is done first to give a rough equivalence-point (end-point), then repeated more accurately to give concordant results. 4

33. Questions relating to the experiment
In this experiment why is dilute sulfuric acid added to the permanganate solution?
Why are burette readings taken from the top of the meniscus?
How is the end-point of the titration detected?
Why is a rough titration carried out?
Why is more than one titration carried out subsequently?

34. 2016 Higher Chemistry Q20 A

35. c) Calculations concerning reactions which involve solutions
Learning intention
Learn how to calculate quantities of reactant or product for reactions involving solutions, using the concentration.

36. Neutralisation Reactions/titrations
What volume hydrochloric acid 1.0 mol l-1 is needed to neutralise
50 cm3 of potassium hydroxide solution concentration 0.25 mol l-1?
12.5cm3

37. Calculations for you to try.
Calculate the concentration of potassium hydroxide (KOH ) if 14.8 cm3 is required to neutralise 20cm3 of 0.1 mol/l nitric acid (HNO3).
Calculate the volume of 0.15 mol/l sulphuric acid(H2SO4) if it is neutralised by 25 cm3 of 0.25mol/l sodium hydroxide (NaOH).
0.135 mol/l
20.83 cm3

38. d) Excess Learning intention
Learn how to calculate how much of a particular reactant is in excess from the balanced equation. 38

39. Graphs and Rates of Reaction
e.g. Zn + 2HCl  ZnCl2 + H2
Zn in excess
Vol H2
cm 3
time /s
2mol l-1 HCl 20oC
2mol l-1 HCl 40oC
Faster, but same amount of
gas produced
1 mol l-1 HCl 20oC
Half the gas produced
HCl is limiting reagent 39 39

40. Vol H2 cm 3 time /s Calculations for you to try.
The graph below was obtained when 1.0g of powdered zinc was added to excess hydrochloric acid 1.0 mol l-1, copy the graph and sketch a line to show what you would expect if the reaction was repeated using
a) 2.0 mol l-1 HCl and 1.0g Zn
b) 1.0 mol l-1 HCl and 0.75g Zn
Vol H2
cm 3
time /s 40

41. Word Definition Standard Solution Indicator End-Point Titration
Glossary
Word
Definition
Standard Solution
Indicator
End-Point
Titration