Q.1 Volumetric Questions
Go over all volumetric from 2002-2014
2012
2002
2008
2006
2014
Q.1
2003
2009
2007
2011
2004
2010
2005
2013
Preparation of a standard solution of sodium carbonate
Standardisation of a hydrochloric acid solution
Hydrochloric acid / Sodium hydroxide titration and use
of this reaction to make salt NaCl
Determination of the concentration of ethanoic acid in
vinegar
Determination percentage water of crystallisation in a
sample of hydrated sodium carbonate
A potassium manganate (VII) / ammonium iron(II)
sulphate titration
Determination of the amount of iron in an iron tablet
An iodine \ thiosulphate titration
Determination of the percentage of hypochlorite in
bleach
Estimation of total hardness using
ethylenediaminetetraacetic acid. [edta]
Estimation of dissolved oxygen by redox titration
Some of the exam questions are below
the volumetric tips & expts. The rest are
in the Exam Papers
Volumetric Tips + 10 Experiments
Titration Precautions
Pipette
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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 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 surface.
DON'T BLOW it is calibrated to allow for the drop at the tip.
Burette
o Rinse with deionised water and then with the solution it is going to contain.
o Fill using a funnel and remove it as drops may fall from it or it may dip into
the liquid giving a false level.
o Remove the air bubble from the tip by opening tap quickly
o Read from the bottom of meniscus - with eye level with this point.
o KMnO4 - read from the top of the meniscus (or from the bottom with a light
behind)
o Don't put NaOH in burette it may react with glass of burette or block tap
[not really valid now].
Conical Flask
o Rinse out with deionised water only.
o Place on white tile - to see colour change more easily.
o Mix continuously.
o Add only a few drops of indicator
(They are weak acids or bases and may upset the results)
o Wash down drops on the side of the flask with deionised water.
(This won’t affect amount of reactant in flask or change the result.)
Volumetric Flask
o Long thin neck to make it accurate.
o Read from bottom of meniscus at eye level.
o Make sure it is at room temperature - it is calibrated at 20oC.
o Mix by inverting 20 times to make sure the solution is homogeneous - long
thin neck makes this necessary.
Titration
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Use the correct indicator [SAWWAP]
Only 3 - 4 drops of indicator
One rough and 2 accurate titres
two accurate should be within 0.1 cm3
Average the 2 accurate titres
Mix well by swirling
Add from burette drop by drop near the end point.
Point at which reaction is complete - shown by colour change
o Identify the standard solution [one you are given the concentration of] - for
calculations that are to follow.
Other Precautions
MnO4o must be acidified (with dil. sulphuric acid)
o Otherwise reaction stops at MnO2 [brown precipitate.] instead of going on to the
colourless Mn2+.
o It acts as its own indicator going from purple to colourless.
o The Mn2+ acts as an autocatalyst. (One of the products of the reaction catalyses the
reaction)
o Read from the top of the meniscus as it is too deeply coloured to see through it
clearly
Iodine / thiosulphate [I2 / S2O32-]
o Use starch indicator
o Blue/black while iodine (I2) is present - colourless when iodine has gone i.e. I2 all
been turned to iodide (I-).
o Don't add the starch indicator too early - it may complex with iodine and ruin the
result.
o Add it when the iodine solution is straw coloured. (e.g. Winkler Method)
Indicator Choice
SAMWAP
o Strong Acid Strong Base - Any indicator
o Strong Acid Weak Base - Methyl Orange
o Weak Acid Strong Base - Phenolphthalein
o Weak Acid Weak Base - None
Standards
o Standard Solution is one whose concentration is known accurately
o Primary Standard is one that can be made up directly using a measured amount of
pure solid.
o Secondary Standard Make up a solution and then standardise this solution using a
primary standard.
o Standardise means to find the concentration of using titration
o Can't use MnO4- as Primary Standard - it can’t be got pure.
o Can’t use iodine [I2] because it sublimes.
o Can’t use KOH or NaOH they absorb CO2 and moisture from air
o Can’t use conc. H2SO4 absorbs water from air
Making a salt e.g. NaCl
o Use HCl and NaOH
o Do titration with indicator
o Get volumes
o Do without indicator
o Evaporate water leaving pure salt
Indicator Colours
o Methyl Orange Acid Red Alkali Orange/yellow
o Phenolphthalein Acid clear Alkaline Purple/pink
Exp 4.1 Preparing a Standard solution of Sodium Carbonate
 Weigh out 2.65 g of anhydrous Na2CO3 on a clock
glass [[23x2] + 12 +[16x3]/10 *250/1000]
 Place in beaker with 100 ml deionised water with
washings
 Stir to dissolve
 Place into 250 ml volumetric flask with washings
 Make up to mark with deionised water
 On level surface with bottom of meniscus and eye
level with mark
 Invert 20 times to make homogeneous
 Label
Exp 4.2 To standardise a solution of HCl using a standard solution of
Na2CO3
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Pipette 25 cm3 dil. HCl into a 250cm3 volumetric flask
Make up to the mark with deionised water (dilutes acid to a factor of 10)
Mix well to make homogeneous
Pipette 20ml HCl into a conical flask
Add methyl orange indicator [SAMWAP] and note the colour
Add Na2CO3 from burette
Mix constantly
Wash any drops from side with deionised water [doesn’t affect amount
acid in flask]
Note volume as indicator changes colour (red → orange)
Do one rough and two accurate titres and take the average of two accurate
Use VaMa/na = VbMb/nb to solve
Multiply answer by 10 to allow for dilution
Exp 4.3
To determine the percentage of Ethanoic
acid in vinegar
 Pipette 20 ml NaOH into conical flask,
 Add a few drops of phenolphthalein indicator
(SAMWAP), turns pink.
 Put Ethanoic acid in burette (remove air bubble).
 Record initial volume
 Add to NaOH in flask
 When pink turns colourless note volume (adding
dropwise near end).
 Do one rough and two accurate and take the average of two
accurate
Exp 4.4 To determine the percentage of water of crystallisation in
hydrated
sodium carbonate (washing soda)
 weigh out 5g washing soda crystals [ Na2CO3]
 dissolve in 100cm3 deionised water, stir
 transfer to 250 cm3 flask with washings, make solution up to mark, invert
20 times
 pipette 25cm3 of Sodium carbonate solution into conical flask
 add 3 drops of methyl orange indicator goes yellow [SAMWAP]
 fill burette with HCl solution
 titrate: colour change from yellow to pink
 do one rough and 3 accurate – average of 2 accurate
 Calculation - use: VaMa = VbMb
na
nb
 Mass of anhydrous Na2CO3 = M x 106
 mass of water = mass of crystals - mass of anhydrous Na2CO3 =
 % anhydrous Na2CO3 = mass of anhydrous Na2CO3
% water =
mass of water
mass of crystals
mass of crystals
 % anhydrous Na2CO3 : % water
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106
18
Exp 4.5 Potassium manganate(VII) / Ammonium iron(II) sulphate titration
 Ammonium iron(II) sulphate primary standard solution
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KMnO4 in burette – read from top of meniscus
Pipette 25 ml ammonium iron(II) sulphate into conical flask
Add 10 ml dil. H2SO4 to make sure doesn’t stick at brown pptte. of MnO2
Mn2+ autocatalyst - product of reaction catalyses reaction
Acts as own indicator
End point when pink colour remains permanently
Calculation - use: VaMa = VbMb
na
nb
12+
Equation: MnO4 + 5 Fe + 8 H+ = Mn2+ + 5 Fe3+ + 4 H2O
Exp 4.6 To determine the amount iron in iron tablet.
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Find the mass of the five iron tablets.
Crush with pestle and mortar in dilute H2SO4 [to stop Fe2+ - Fe3+]
Transfer the paste with washings to beaker
Stir to dissolve the paste.
Transfer solution into 250cm3 volumetric flask with washings
Make up to mark with deionised water. Eye/meniscus level with mark.
Invert 20 times.
Pipette 25cm3 of tablet solution into conical flask.
Add 20cm3 of dilute sulphuric acid to ensure full reaction
Titrate with KMnO4 from burette until a permanent pink colour remains
Equation MnO41- + 5 Fe2+ + 8 H+ = Mn2+ + 5 Fe3+ + 4 H2O
Mn2+ autocatalyst - product of reaction catalyses reaction
One rough and two accurate. Average of 2 accurate
Calculation - use: VaMa = VbMb
na
nb
Mass = molarity * RMM = M * 152 =
Multiply by 1000 to give mg
Divide by 5 to give mass of each tablet
Exp 4.7 An Iodine/Thiosulphate Titration
 Make up 0.05M solution of iodine
 By reacting 0.017 M potassium iodate with excess potassium iodide
 The iodide then reacts with the sodium thiosulphate, which is added from
a burette.
 I2 + 2 S2 O32- = S4O62- + 2I Weigh 6.25g of sodium thiosulphate crystals onto clock glass
 Transfer crystals to beaker containing 100cm3 of de-ionised water
 Stir and when dissolved transfer to 250cm3 volumetric flask with
washings
 Make up to mark with de-ionised water. Invert 20 times.
 Pipette 25cm3 of potassium iodate
 Use graduated cylinder to add 20cm3 of dilute sulphuric acid, followed by
10cm3 of 0.5M potassium iodide solution to the conical flask. Note a
reddish/brown colour of liberated iodine.
 Titrate. Until pale yellow. Add a few drops of starch indicator - blue
colour is observed.
 Continue adding thiosulphate solution drop by drop until blue colour
disappears, to give colourless solution. Note the titration figure.
 One rough and two accurate titres. Average two accurate
 Vo * Mo / no
=
Vr * Mr / nr
Exp 4.8 To determine the percentage (w/v) of sodium hypochlorite NaOCl
in bleach
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Pipette 25 ml concentrated bleach into 250 ml volumetric flask
Fill up to the mark with deionised water, invert it 20 times.
Diluted by factor 10
Pipette 25 ml bleach into conical flask, add 20 ml sulphuric acid and 10
ml of 0.5M potassium iodide (KI) solution. Note reddish/brown colour
Titrate until a pale yellow colour
Add a few drops of starch indicator, note blue colour
Add thiosulfate drop by drop until colour change from blue to colourless.
Note titration figure.
Do one rough and two accurate readings. Take average of two accurate
Calculate molarity of sodium hypochlorite NaClO3
Vo * Mo / no
=
Vr * Mr / nr
Multiply by 10 for dilution
Calculate mass of hypochlorite in 1 litre [mass = molarity * RMM of
hypochlorite]
Divide by 10 to get mass in 100g
Exp 9.3 Estimation of Total Hardness of water using Ethylenediaminetetraacetic acid
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Place 50 ml of hard water in conical flask
Place 0.01 M EDTA in burette
Add 1 cm3 of buffer solution [pH 10] to keep alkaline so indicator works properly
Add 5 drops of Erichrome black – gives wine red colour
Add EDTA from burette until solution turns blue
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Take 1 rough and 2 accurate titres - average 2 accurate
1 cm3 of 0.1 M EDTA ≡ I mg CaCO3
Multiply average titre value by 20 to find mg L-1 [ p.p.m.]
Unboiled = permanent + temporary
Boiled water = permanent only
Temporary = Unboiled value – boiled value
Exp 9.4 To measure the amount of Dissolved Oxygen by Winkler Method
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Rinse bottle to stop bubbles forming
Fill and stopper under water
add 1cm3 of conc. manganese(II) sulphate [so not to upset volume]
add 1cm3 of conc. alkaline KI [so not to upset volume]
Brown precipitate forms – if white precipitate forms no oxygen present.
Add 1cm3 of conc. H2SO4.
Solution goes brown colour due to the liberated iodine.
Put 100 cm3 of iodine solution into conical flask.
0.005 M Sodium thiosulphate in burette is standard solution
Titrate until a pale straw coloured
Add a few drops of starch indicator
Continue titrating until the blue/black colour disappears.
Do one rough and 2 accurate – average 2 accurate
Calculate the concentration of O2 in water expressing your results in p.p.m.
Let dissolved oxygen = a and Let thiosulphate = b
Ratio of dissolved oxygen to thiosulphate = 1:4 therefore na = 4 and nb = 1
Calculate molarity of the oxygen
Multiply by 32 to convert to grams per L
Multiply by 1000 to convert grams to mg/L [p.p.m.]
Titration
2002
1. Vinegar is a solution of ethanoic acid (acetic acid). Some bottles of vinegar are labelled “White
Wine Vinegar”.
(a) What compound in white wine is converted to ethanoic acid in vinegar? What type of chemical
process converts this compound to ethanoic acid? (8)
The concentration of ethanoic acid in vinegar was measured as follows: A 50 cm3 sample of vinegar
was diluted to 500 cm3usingdeionised water. The diluted solution was titrated against 25 cm3 portions
of a standard 0.12 M sodium hydroxide solution, using a suitable indicator.
(b) Describe the procedure for accurately measuring the 50 cm3 sample of vinegar and diluting it to
500 cm3. (12)
(c) Name the piece of equipment that should be used to measure the ethanoic acid solution during the
titration. State the procedure for washing and filling this piece of equipment in preparation for the
titration. Name a suitable indicator for this titration. (15)
The titration reaction is
CH3COOH + NaOH → CH3COONa + H2O
After carrying out a number of accurate titrations of the diluted solution of ethanoic acid against the
25 cm3 portions of the standard 0.12 M sodium hydroxide solution, the mean titration figure was
found to be 20.5 cm3.
(d) Calculate the concentration of ethanoic acid in the diluted vinegar solution in moles per litre and
hence calculate the concentration of ethanoic acid in the original sample of vinegar. Express this
concentration in terms of % (w/v). (15)
2003
1. Iron tablets may be used in the treatment of anaemia. To analyse the iron(II) content of
commercially available iron tablets a student used four tablets, each of mass 0.360 g, to make up 250
cm3 of solution in a volumetric flask using dilute sulfuric acid and deionised water.
About 15 cm3 of dilute sulfuric acid was added to 25 cm3 portions of this iron(II) solution and the
mixture then titrated with a 0.010 M solution of potassium manganate(VII), KMnO4.
(a) Why was it important to use dilute sulfuric acid as well as deionised water in making up the
solution from the
tablets? (5)
(b) Describe in detail the procedure for making up the 250 cm3 solution from the tablets. (18)
(c) Why was more dilute sulfuric acid added before the titrations were commenced? (6)
(d) How was the end-point detected? (3)
The titration reaction is described by the equation
MnO4 + 5Fe2+ + 8H+ = Mn2+ + 5Fe3+ + 4H2O
(e) In the titrations the 25 cm3 portions of the iron(II) solution made from the tablets required 13.9 cm3
of the
0.010 M KMnO4 solution. Calculate
(i) the concentration of the iron(II) solution in moles per litre
(ii) the mass of iron(II) in one tablet
(iii) the percentage by mass of iron(II) in each tablet. (18)
2004
1. In an experiment to determine the total hardness of a water sample containing both calcium and
magnesium ions, a solution of the reagent edta (ethylenediaminetetraacetic acid) in the form of its
disodium salt (represented by Na2H2Y) was titrated against a sample of the water using a suitable
indicator. The reaction between the ions (represented by M2+n the hard water and the edta reagent
may be represented as
M2++ H2Y2- = MY2- + 2H+
(a) Name a suitable indicator for this titration. What colour change is observed at the end point of the
titration using this indicator? (8)
(b) Describe the correct procedure for rinsing the burette and filling it with edta reagent. (15)
(c) The addition of a small quantity of another solution to the water in the conical flask is essential
before
commencing the titration. What solution must be added and what is its purpose? (6)
(d) In the experiment it was found that 100 cm3 portions of the water required an average titre of 8.10
cm3 of 0.010
M edta solution. Calculate the total hardness in (i) moles per litre, (ii) grams per litre expressed in
terms of
CaCO3 and (iii) p.p.m. expressed in terms of CaCO3 (15)
(e) A whitish deposit is often found on the insides of kettles in hard water districts. If some of this
deposit is
scraped into a test tube and dilute hydrochloric acid is added a reaction is observed. Write a
balanced equation
for this reaction. (6)
2005
1. In an experiment to measure the concentration of dissolved oxygen in a river water sample, a bottle
of
water was filled from the river and it was analysed immediately. The experiment was carried out as
follows:
A few cm3 each of concentrated manganese(II) sulfate (MnSO4) solution and alkaline potassium
iodide (KOH/KI) solution were added to the water in the bottle. The stopper was carefully replaced
on the bottle and the bottle was shaken to ensure mixing of the reagents with the water. A brownish
precipitate was produced. The stopper was removed from the bottle and a few cm3 of concentrated
sulfuric acid (H2SO4) were added carefully down the inside of the neck of the bottle using a dropper.
The precipitate dissolved and a golden-brown solution was produced. The concentration of iodine (I2)
in this solution was found by titrating it in 50 cm3 portions against a standard (0.01 M) sodium
thiosulfate (Na2S2O3) solution.
(a) Why was it necessary to analyse the sample of river water immediately? (5)
(b) In making the additions to the sample, why should the solutions used be concentrated? (6)
(c) Describe how the additions of the concentrated solution of manganese(II) sulfate (MnSO4) and
alkaline potassium iodide (KOH/KI) to the bottle of river water should be carried out. What essential
precaution should be taken when replacing the stopper of the bottle after these additions are made? (9)
(d) Describe clearly the procedure for using a pipette to measure exactly 50 cm3 portions of the iodine
(I2) solution into the titration flask. (9)
(e) What indicator is used in this titration? State when the indicator should be added to the titration
flask and describe the colour change observed at the end point. (9)
(f) The titration reaction is described by the following equation.
2S2O3 + I2 2I¯ + S4O6
Calculate the concentration of the iodine solution in moles per litre given that 6.0 cm3 of the 0.01 M
sodium thiosulfate (Na2S2O3) solution were required in the titration for complete reaction with 50 cm3
portions of the iodine solution. (6)
(g) For every 1 mole of oxygen gas (O2) in the water sample 2 moles of iodine (I2) are liberated in this
experiment. Hence calculate the concentration of dissolved oxygen in the water sample in p.p.m. (6)
2006
1. An experiment was carried out to determine the percentage water of crystallisation and the degree
of water of crystallisation, x, in a sample of hydrated sodium carbonate crystals (Na2CO3.xH2O). An
8.20 g sample of the crystals was weighed accurately on a clock glass and then made up to 500 cm3 of
solution in a volumetric flask. A pipette was used to transfer 25.0 cm3 portions of this solution to a
conical flask. A previously standardised 0.11 M hydrochloric acid (HCl) solution was used to titrate
each sample. A number of accurate titrations were carried out. The average volume of hydrochloric
acid solution required in these titrations was 26.05 cm3.
The titration reaction is described by the equation:
Na2CO3 + 2HCl = 2NaCl + CO2 + H2O
(a) Identify a primary standard reagent which could have been used to standardise the hydrochloric
acid solution. (5)
(b) Name a suitable indicator for the titration and state the colour change observed in the conical flask
at the end point. Explain why not more than 1 – 2 drops of indicator should be used. (12)
(c) (i) Describe the correct procedure for rinsing the burette before filling it with the solution it is to
deliver.
(ii) Why is it important to fill the part below the tap of the burette? (12)
(d) From the titration figures, calculate the concentration of sodium carbonate (Na2CO3) in the
solution in (i) moles per litre, (ii) grams per litre. (9)
(e) Calculate the percentage water of crystallisation present in the crystals and the value of x, the
degree of hydration of the crystals. (12)
Helpful Hints
Definitions
These can be worth up to 15% of the marks on any given paper so it is well worth
spending time learning them. All you hve to do is regurgitate them. There is the
added benefit that knowing them makes understanding topics easier.
2007
1. A solution of sodium thiosulfate was prepared by weighing out a certain mass of crystalline sodium
thiosulfate (Na2S2O3.5H2O) on a clock glass, dissolving it in deionised water and making the solution
up carefully to 500 cm3 in a volumetric flask. A burette was filled with this solution and it was then
titrated against 25.0 cm3 portions of previously standardised 0.05 M iodine solution in a conical flask.
The average titre was 20.0 cm3. The equation for the titration reaction is 2S2O3 + I2 → 2I¯ + S4O6
(a) Sodium thiosulfate is not a primary standard. Explain fully the underlined term. (8)
(b) Describe how the crystalline thiosulfate was dissolved, and how the solution was transferred to the
volumetric flask and made up to exactly 500 cm3 (15)
(c) Pure iodine is almost completely insoluble in water. What must be added to bring iodine into
aqueous solution?
(3)
(d) A few drops of freshly prepared starch solution were added near the end point as the indicator for
this titration. What sequence of colours was observed in the conical flask from the start of the
titration until the end point was reached? (12)
(e) Calculate the molarity of the sodium thiosulfate solution and its concentration in grams of
crystalline sodium thiosulfate (Na2S2O3.5H2O) per litre. (12)