1. Speaking volumes Titration

2. Question Layout Usually Question 1
One of three practical questions of which you must do at least two
Four parts
Using Apparatus
Indicator, colour change
Special conditions
Calculations

4. Part 1 – Using Apparatus Burette Clamp it vertically
Add solution drop-wise when near end point
KMnO4 - read from the top of the meniscus (or from the bottom with a light behind)
Don't put NaOH in burette it may react with glass of burette or block tap [not really valid now].
Rinse with deionised water to wash out any impurities
Then with the solution it is going to contain to wash out the deionised water.
Fill using a funnel and remove it as drops may fall from it or it may dip into the liquid giving a false level.
Remove the air bubble from the tip (jet) by opening tap quickly
Read from the bottom of meniscus - with eye level with this point.

5. Pipette Rinse with deionised water to wash out any impurities
Then with the solution it is going to contain to wash out the deionised water.
Fill using pipette filler - the solution may be poisonous or caustic.
Read from the bottom of meniscus
meniscus level with the ring on the stem
eye level with this.
Empty into the conical flask and touch tip against the side.
DON'T BLOW it is calibrated to allow for the drop at the tip.

6. Conical Flask Rinse out with deionised water only.
Place on white tile (to see colour change more easily
Mix continuously by swirling contents.
Add only a few drops of indicator (They are weak acids or bases and may upset the results)
Wash down drops on the side of the flask with deionised water. (This won’t effect amount of reactant in flask or change the result.)

7. Volumetric Flask Long thin neck to make it accurate.
Read from bottom of meniscus at eye level.
Make sure it is at room temperature - it is calibrated at 20oC.
Mix by inverting 20 times to make sure the solution is homogeneous - long thin neck makes this necessary.

8. Titration Use the correct indicator
Only drops of indicator [weak acid-base]
One rough and 2 accurate titres
The two accurate should agree within 0.1 cm3
Average the 2 accurate titres
Mix well by swirling (NOT shaking)
Add from burette drop by drop near the end point [Point at which reaction is complete - shown by colour change
Identify the standard solution [one you are given the concentration of] - for the calculations that are to follow.

9. Part 2 – Indicator choice & colour change
Indicator choice with reason
Colour change(s)

10. Part 2 – Indicator choice & colour change
SASBANY (Inuit for don’t eat yellow snow)
Strong Acid Strong Base - Any indicator
SAWBMO (Inuit for don’t eat polar bears liver)
Strong Acid Weak Base - Methyl Orange
WASBPH
Weak Acid Strong Base - Phenolphthalein
WAWBNONE
Weak Acid Weak Base – None
Indicator should change colour at a pH where pH changes rapidly during the titration

11. Bronsted-Lowry – acid - good proton donor:
Strong versus Weak
Strong:
Arrhenius - dissociates fully in aqueous solution
Bronsted-Lowry – acid - good proton donor:
base – good proton accepter
Weak:
Arrhenius - dissociates slightly in aqueous solution
Bronsted-Lowry – acid - poor proton donor:
base – poor proton accepter
Strong Acid
HCl
H2SO4
HNO3
Weak Acid
HCOOH
CH3COOH
H2CO3
Weak Base
NH4OH
Na2CO3
Strong Base
NaOH
KOH

12. Titration Curve Choose Indicator that Changes colour
Where graph is vertical
Strong base changes pH rapidly from 6.5 to 10.5
As does phenolphthalein
So use phenolphthalein
For weak acid with strong base
Titration Curve
Phenolphthalein
Methyl Orange

13. Indicator Colours Indicator Acid Colour Alkaline Methyl Orange Pink
Litmus
Red
Blue
Phenolphthalein
Colourless
Purple

17. Part 3 - Special conditions
Some titrations need special conditions
We will deal with each of these as we describe the titration

18. 1. Preparing a Standard Solution
Standard Solution is one whose concentration is known accurately
Primary Standard is one that can be made up directly using a measured amount of pure solid.
To be suitable for use as a primary standard a substance must be
Pure
Very soluble
Have a high molecular mass
And can be dissolved in water to make a solution of known concentration.

19. Secondary Standard: Make up a solution and then standardise this solution using a primary standard.
Standardise: means to find the concentration of using titration
Can't use MnO4- as Primary Standard - because it can’t be got pure.
Can’t use iodine [I2] because it sublimes.
Can’t use KOH or NaOH they absorb CO2 and moisture from the atmosphere

20. Preparing a Standard Solution of Sodium Carbonate
Weigh out an accurate mass of anhydrous crystals on a clock glass
Place in beaker of deionised water
Rinse clock glass into beaker using deionised water
Stir until fully dissolved
Pour into volumetric flask using a funnel
Rinse beaker and funnel into flask
Fill till bottom of meniscus is level with mark (eye also level with both)
Stopper and invert 20 times to make homogeneous
Long neck makes it accurate and makes mixing necessary

22. 2. HCl – NaOH titration Indicator – Any – strong acid/strong base
Put NaOH into the flask
Do with indicator to get average titre
This tells you exactly how much HCl to add to neutralise the NaOH
Do without indicator
Evaporate water to obtain salt crystals
Allow crystals to dry in air or in warm place e.g over a radiator

24. 3. Concentration of Ethanoic Acid in Vinegar using NaOH Standard Solution
Dilute by factor of 10 [25 cm3 made up to 250 cm3 in volumetric flask]
Why? Acid is too concentrated
Indicator – Phenolphthalein (WASBPH)
Pink in alkaline – colourless in acid
Vinegar used as a condiment or to make cellulose acetate film
If you have diluted don’t forget to multiply your answer by 10 at the end

25. 4. Standardise a KMnO4 Solution
KMnO4 in burette – read from top of meniscus
Standard solution – Ammonium iron(II)sulfate
Indicator – none as MnO4 acts as own indicator
Sulphuric acid added to ensure reaction goes to completion
If not enough get brown precipitate of MnO2
Autocatalysis – first drop of MnO4- goes colourless slowly
Next drops faster as product of reaction (Mn2+) catalyses reaction
End point when solution becomes permanent pink

26. 5. Amount of Iron in an Iron Tablet
Crush tablets (using pestle in mortar to speed up solution) in dilute H2SO4
Add 1st lot of sulfuric acid as crushing tablets to stop atmospheric oxygen oxidising Fe2+ to Fe3+
Add second lot of H2SO4 to conical flask for titration to ensure reaction goes to completion
If not enough get brown precipitate of MnO2
End point when solution becomes permanent pink

27. 6. Degree of Hydration of Sodium Carbonate
Weigh out accurate mass of hydrated crystals (with water of crystallisation)
Dissolve in deionised water
Make up to mark in (250 cm3) volumetric flask
Invert 10 times to make homogeneous
Pipette (20 cm3) into conical flask
Add methyl orange indicator (yellow in base) SAWBMO
Titrate with previously standardised solution of HCl
End point when turns from yellow to red

28. Calculate the molarity of Na2CO3 (anhydrous)
Calculate No moles of Na2CO3 in 20cm3 [× by 20/1000*M]
Work out mass of anhydrous (without water of crystallisation) Na2CO3 [moles × by 106]
Work out mass of water of crystallisation present in crystals [mass of 1st crystals – mass of anhydrous Na2CO3)
Calculate ratio:

29. 7. Iodine – Thiosulphate Titration
Iodide ions are used to help iodine dissolve
Starch Indicator
Add when solution straw coloured
Otherwise will complex with iodine and ruin result
Goes blue-black (due to starch)
End point is when it turns colourless (when all the iodine is gone)

30. 8. % of Hypochlorite in Bleach
Dilute by factor of 10 (25 cm3 into 250cm3)
Concentrated bleach can burn
Pipette 20 cm3 diluted bleach into conical flask
Add (20 cm3) dilute H2SO4
Add (10cm3) 0.5 M KI to help iodine dissolve
Brown colour of Iodine appears
ClO I H+ = Cl- + I2 + H2O
Titrate with standard solution of Na2S2O3
When straw coloured add starch indicator

31. Blue-black colour formed
End point when it goes colourless
Calculate concentration of ClO-
Eq1 ClO- = I2 and
Eq2 I2 = 2S2O3
so 1ClO- = 2S2O3

32. 9. Determine the Hardness of Water
Hardness always measured as ppm of CaCO3
Indicator: Eriochrome Black
Colour change: wine-red ⇒ blue
Buffer solution keeps pH =10 so indicator will work properly
Unboiled water = permanent + temporary hardness
Boiled = Permanent hardness only
Temporary hardness = unboiled – boiled value

33. 10. Dissolved Oxygen – Winkler Method
Wet bottle to stop bubbles of air sticking to sides
Fill and stopper under water
Add conc. MnCl2 and Alkaline KI using dropper under water surface to prevent reaction with O2 in air
Mn2+ + OH- = Mn(OH)2 [white ppt.]
Small amounts: so as not to change volume of sample
Stopper without trapping air bubbles
Invert 20 times
2Mn(OH)2 + O2 = 2MnO(OH)2 [brown]
If White Precipitate this means no dissolved oxygen so dilute with fully oxygenated water

34. Add conc. H2SO4
Brown pcpt. dissolves as it becomes I2
Place 50 cm3 in conical flask
Titrate using standard solution of thiosulphate
Starch Indicator
Add when solution straw coloured
Otherwise will complex with iodine and ruin result
Goes blue-black (due to starch)
End point is when it turns from blue-black to colourless (when all the iodine is gone)
Place second sample in dark cupboard (to prevent photosynthesis which would produce O2)
For 5 days at 20OC (to allow organic matter time to decay and use up oxygen) and retest for dissolved O2
Difference in ppm dissolved O2 is BOD

35. Part 4 Calculations Calculate molarity Convert to other units
Balanced equation will be given
Assign (a) and (b) to each of the two reactants
Fill in the following table
Volume Va = Vb =
Molarity ( one will be ?) Ma = Mb =
No. Moles in eqn. na = nb =
Use equation

36. Example

38. nb= 1 na = 2 a b a = S2O32- b= I2 Va = Vb= Ma = Mb= na = nb=
Molarity: mol l-1 (6) or
Va Ma = Vb Mb
na nb 20 25
 20 x M = 25 x (3)
M = (3) ?
0.05
Grams l-1: 31 g l-1 (6) or
Conc. In g l-I = molarity * molar mass
0.125 x 248* (3) = 31 g l-1 (3) 2 1
* Addition must be shown for error to be treated as a slip
[Na2S2O3.5H2O = 2* * * (2 * ) = 248]
Slip loses only 1 mark

39. Solving from First Principles
For last two years
It has not been possible to solve using VaMa /na = VbMb /nb equation alone
Some work from first principles has been required

40. Now calculate the molarity of the iodine solution
Sodium thiosulfate is a reducing agent that reacts with iodine according to the following balanced equation. I2 + 2S2O32– → 2I– + S4O62–
To determine the concentration of a sodium thiosulfate solution, a student titrated it against 25.0 cm3 portions of a standard iodine solution.
Given that there were 6.35 g of iodine (I2) in 500 cm3 of the iodine solution, calculate
the number of moles of iodine in each 25.0 cm3 portion,
the number of moles of sodium thiosulfate required to reduce this quantity of iodine,
the concentration of the sodium thiosulfate solution in moles per litre, cm3 of which were required to reduce 25.0 cm3 of the iodine solution,
the concentration of the sodium thiosulfate solution in grams per litre of its crystals (Na2S2O3.5H2O). (18)
(i) Find the number of moles of standard solution you have – in this case iodine Moles = Mass/RMM (I2) = 6.35 / (127×2) = 0.025
Now calculate the molarity of the iodine solution
I.e. No of moles in 1 litre (1000 cm3) = 0.025×1000/volume = 0.025×1000/500 = 0.05
Therefore 25 cm3 will contain 25/1000 of 0.05 = mole
[Once you know the molarity you can answer (iii) using VaMa /na = VbMb /nb but not (ii)]
(ii) From the equation the number of moles of thiosulfate is twice the number of moles of iodine I2 + 2S2O32– = ×2 =
(iii) cm3 of thiosulfate contains moles therefore 1 cm3 contains /17.85
so 1 litre (1000cm3) contains 1000 × /17.85 = M
(iv) conc. = molarity × Molar mass = * 248 = = 35g/l

41. NB. Dilution by factor of 250/25 =10
To determine the concentration of ethanoic acid in a sample of vinegar, 25.0 cm3
of the vinegar were diluted to 250 cm3 and then the diluted vinegar was titrated with a previously standardised solution which contained 1.20 g of sodium hydroxide in 500 cm3
ofsolution. On average, cm3 of the diluted vinegar were required to neutralise 25.0 cm3
of this sodium hydroxide solution.
The equation for the titration reaction is: CH3COOH + NaOH  CH3COONa + H2O
NB. Dilution by factor of 250/25 =10
(i) Calculate molarity of NaOH = Mass/RMM (NaOH) = 1.2/40 × 1000/500 = 0.06 (3)
Moles of NaOH = vol/1000 × molarity = 25/1000 × 0.06 = (3) (ii) From equation the number of moles of ethanoic acid in cm3 = moles of NaOH
18.75 cm3 = (3) in 1cm3 = /18.75 = (3)
(d) (i) 1 cm3 = mol so 1 litre (1000cm3) contains 1000 × = 0.08 mol (3)
Diluted by factor of 10 so moles per litre in undiluted vinegar =008 × 10 = 0.8 (3)
(ii) Mass of ethanoic acid in 1 litre of 0.8 molar = 0.8 × 60 = 48 g/L (3)
In 100 cm3 = 48/10 = 4.8% (w/v) (3)
Calculate
(i) the number of moles of sodium hydroxide in each 25.0 cm3 portion,
(ii) the number of moles of ethanoic acid per cm3 of diluted vinegar. (12)
(d) Find the concentration of ethanoic acid in the original vinegar
(i) in terms of moles per litre,
(ii) as a percentage (w/v). (9)

42. Conversion Table (Mole Map)
Mass
No of Particles
Volume
÷ Molar Mass
÷ 6 x 1023
÷ 22.4 if dm3
or if cm3
Moles
x 22.4 if dm3
or if cm3
x Molar Mass
x 6 x 1023
Mass
No of Particles
Volume

43. END