Doc Brown's Advanced A Level Chemistry Revision Notes

Some Advanced A Level Practical Exercises and Calculations involving volumetric or gravimetric analysis

 Index [(a) to (i) are quoted from AQA AS/A level chemistry]

(a) How to determine the concentration of ethanoic acid in vinegar

(b) How to determine the mass of calcium carbonate in an indigestion tablet

(c) How to determine the formula mass of an MHCO3 hydrogencarbonate

(d) How to determine the molecular mass of a dibasic acid eg succinic acid

(e) How to determine the mass of aspirin in an aspirin tablet

Check out what is available? Study the different examples then try the Quizzes!(f) How to determine the yield for the conversion of magnesium to magnesium oxide

(g) How to determine the formula mass of a hydrated salt (eg magnesium sulfate) by heating to constant mass

(h) How to determine the percentage conversion of a Group 2 carbonate to its oxide by heat

(i) How to determine the number of moles of water of crystallisation in a hydrated salt by titration


(a) How to determine the concentration of ethanoic acid in vinegar

You can find the concentration of ethanoic acid in vinegar by titrating vinegar with a standard solution of sodium hydroxide.

Since ethanoic acid is a weak acid and sodium hydroxide a strong base you would use phenolphthalein indicator.

All you need to know/do is covered on ....

All the basics of doing acid-alkali titrations and simple calculations (including example 12.6 and APPENDIX 2.2)


(b) How to determine the mass of calcium carbonate in an indigestion tablet

You can find the mass of calcium carbonate in an indigestion tablet by dissolving the tablet in excess acid - provided from a precise volume of known concentration e.g. 0.1 mol/dm3 hydrochloric acid.

You would then have to do a back-titration to determine the quantity of unreacted acid.

You then have to subtract the unreacted acid from the total acid at the start to get to moles of acid that reacted with the carbonate.

From the moles of reacted acid you calculate the moles of calcium carbonate in the tablet.

See Q5 on purity of magnesium oxide calculation - involves back titration

How to do acid-alkali titration calculations, diagrams of apparatus, details of procedures


(c) How to determine the formula mass of an MHCO3 hydrogencarbonate

You can find the formula mass of a hydrogen carbonate of a group 1 metal by titrating a known mass with standard hydrochloric solution.

An appropriate quantity of the carbonate is dissolved in deionised water and made up to 250 cm3 in a standard volumetric flask.

25 cm3 aliquots are then titrated with 0.1 molar hydrochloric acid using methyl orange indicator.

See Q8 on titrating group 1 hydrogencarbonates

How to do acid-alkali titration calculations, diagrams of apparatus, details of procedures


(d) How to determine the molecular mass of a dibasic acid eg succinic acid

You can find the molecular mass of a dibasic acid by titration with standard sodium hydroxide solution

See Q15 on titration question on determining molecular mass of a weak monobasic organic acid

See Q16 on titration question on determining molecular mass of a weak dibasic organic acid

How to do acid-alkali titration calculations, diagrams of apparatus, details of procedures


(e) How to determine the mass of aspirin in an aspirin tablet

You can find the mass of aspirin in an aspirin tablet by titration with standardised sodium hydroxide solution.

You can titrate the crushed and accurately weighed powder directly, but dissolve in a little alcohol before adding more water and phenolphthalein indicator.

See Qs 15, 17 and 20 calculations on aspirin/organic carboxylic acid purity

How to do acid-alkali titration calculations, diagrams of apparatus, details of procedures


(f) How to determine the yield for the conversion of magnesium to magnesium oxide

You can find the yield of magnesium oxide from heating magnesium in air in a crucible.

2Mg(s)  +  O2(g)  ==>  2MgO(s)

Puzzled on this one. You readily convert magnesium to magnesium oxide by heating a known mass of magnesium ribbon in a crucible (plus lid, and everything accurately weighed).

You have to heat, cool and reweigh the crucible + lid several times until the total mass is constant.

There should be no unreacted magnesium, but you do get a little magnesium nitride (Mg3N2) formed (especially with the lid on which restricts oxygen, but stops MgO 'smoke' escaping from the crucible!).

From the weighings you can compare the actual yield of MgO and the theoretical amount of MgO, as long as nothing else is formed!

This experiment is fraught with error in my experience!

See Mg ===> MgO reacting mass calculation but I'll work in moles!

  • This reaction can be carried out in a school or college laboratory using a ceramic crucible and lid.

  • The crucible & lid are weighed (m1), a piece of magnesium ribbon added and the crucible weighed again (m2).

  • The difference in weights gives the mass of magnesium (m2 - m1 = m3).

  • The crucible is placed on a clay pipe triangle resting on a tripod above a bunsen burner.

  • The crucible is heated so the magnesium reacts with the oxygen in air, removing the lid to allow air to come in.

  • The crucible is cooled and reweighed with the lid on (m4).

  • m4 - m2 = m5 = mass of oxygen gained.

  • Because the conversion might not be 100% you cannot assume m4 - m1 = mass of MgO

  • Strictly speaking, the procedure should be repeated until no further gain in weight is observed.

  • The gain in weight is due to the solid magnesium combining with oxygen gas from the air to give the solid magnesium oxide.

  • You can then calculate the mass or moles of magnesium oxide formed and compare with the theoretical prediction.

Example of calculation

Relative masses: Mg = 24, O = 16, MgO = 40

initial Mg + O2  ==>  MgO formed + Mg unreacted

I'll just ignore the weighings and get to the heart of the calculation!

Suppose 1.92g of magnesium ribbon yields 3.10g of white residue - MgO

The mass of oxygen gained by Mg = 3.10 - 1.92 = 1.18g

therefore mol oxygen = 1.18/16 = 0.07375

and from the equation ...

mol O atoms = mol MgO = 0.07375, so mol MgO formed = 0.07375

initial mol Mg atoms = maximum mol MgO = 1.92/24 = 0.080

percentage conversion Mg to Mg O = 100 x 0.07375 / 0.080 = 92.2%

BUT this calculation assumes some Mg in the residue.

This is an example of gravimetric analysis - analysis from 'weighings'!


(g) How to determine the formula mass of a hydrated salt (eg magnesium sulfate) by heating to constant mass

 You can find the formula and formula mass of some hydrated salts by heating a given quantity of it to constant mass.

See Q25 gravimetric calculation on heating hydrated sodium sulfate and Q31 heating hydrated copper sulfate

See example of calculation of water of crystallisation and formula mass of copper(II) sulfate


(h) How to determine the percentage conversion of a Group 2 carbonate to its oxide by heat

You can find the percentage conversion of a group 2 carbonate to its oxide by heating an accurately known mass of the carbonate in a crucible (+ lid).

You need accurate mass weighings of (m1) crucible + lid, (m2) crucible + lid + carbonate, (m3) final mass of crucible + lid + oxide residue

(m2) - (m1) gives mass of carbonate, from this calculate theoretical yield of oxide by reacting mass calculation

(m3) - (m1) gives mass of oxide, compare with calculated mass above (maybe!?, see final comment)

Example of method results and calculation

  • This reaction can be carried out in a school or college laboratory using a ceramic crucible and lid.

  • The crucible & lid are weighed (m1), a few grams of group 2 carbonate and the crucible weighed are again (m2).

  • The difference in weights gives the initial mass of group 2 carbonate (m2 - m1 = m3).

  • The crucible & lid are placed on a clay pipe triangle resting on a tripod above a bunsen burner.

  • The crucible & lid are heated so the group 2 carbonate decomposes.

  • The crucible & lid are cooled and reweighed (m4).

  • Strictly speaking, the procedure should be repeated until no further loss in weight is observed.

  • The loss in mass is due to the formation of carbon dioxide gas.

  • Thermal decomposition: MCO3(s)  ===>  MO(s)  +  CO2(g)

  • The mass of residue = m4 - m1 = m5 but you can't assume this is all MO.

  • The mass of carbon dioxide lost = m4 - m2 = m6 (a more crucial measurement).

  • BUT it isn't that simple. m5 may include undecomposed group 2 carbonate.

    • Magnesium carbonate MgCO3 readily decomposes at red heat in a crucible and for MgO observed mass residue should be similar to that theoretically calculated from the reacting mass calculation.

    • But the stability of MCO3 increases down the group, so CaCO3 (limestone) requires 900oC so is only partially decomposes in the crucible.

    • So you need a bit on elaborate reacting mass calculation so that the weighings can sort out conversion with a residual mixture of both MO and MCO3

Example calculation - m = mass 1-6 in grams, you need to follow the logic very carefully from above and onwards!

m3 MCO3 (initial)  ==>  m5 [MO formed + MCO3 undecomposed] residue  +   m6 CO2 lost

m3 we know and m6 we no from the weighings.

So how do we sort out the mixture of m5?

Well from m6 we get moles of CO2 which equals moles of MO

We can then compare moles MO formed with theoretical moles for 100% decomposition, hence calculate the % conversion.

Atomic masses: Ca = 40, C = 12, O = 16, formula masses: CaCO3 = 100, CaO = 56, CO2 = 44

I'll now outline the calculation with limestone, but ignoring the 'real' weighings, that's hardly the difficult part of the calculation!

Suppose after heating 5.00g of calcium carbonate (limestone) the residue in the crucible is 3.20g

Mass of CO2 lost = 5.00 - 3.20 = 1.80g,

so mole CO2 formed = 1.80/40 = 0.045

and from CaCO3  ===>  CaO  +  CO2 (1 : 1 : 1 mole ratio)

we can argue that mol CaO formed = 0.045

and initial mol CaCO3 = maximum theoretical mol CaO = 5.00/100 = 0.05

Therefore % conversion of carbonate to oxide = 100 x 0.045 / 0.05 = 90%

So, it wasn't too bad a calculation after all, was it?!

This is another example of gravimetric analysis - analysis from 'weighings'!


(i) How to determine the number of moles of water of crystallisation in a hydrated salt by titration

If its a hydrated carbonate you can titrate an accurately known mass with standardised hydrochloric acid.

See Q30 on titrating hydrated sodium carbonate


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