Advanced Level Quantitative Chemistry: Volumetric titration calculations 1 (non-redox)
Advanced Level Chemistry Revision on Volumetric Titrations
GCE A Level AS-A2 IB Chemistry Volumetric Analysis: Acid-base and other non-redox volumetric titration quantitative calculation questions
PART 1 Questions 1 to 20
These pages are designed to give you problem solving practice in volumetric titration calculation questions - most involve some kind of volumetric analysis where you titrate exact volumes of solutions, an accurately weighed mass. PART 1 of A Level an chemistry volumetric titration analysis worksheet of structured questions: Worked out titration questions - Q1-8 and Q13-14 & 19 based on acid-base titrations (acid-alkali, oxide, hydroxide, carbonate and hydrogencarbonate) and Q15-18 based on alkali (NaOH)-organic acid titrations e.g. standardising sodium hydroxide solution or like Q20 about aspirin analysis. Q9 includes useful exemplars for coursework on how much to use in titrations including EDTA, Q10-12 are on silver nitrate-chloride ion titrations, further Q's will be added in the future. Appendix 1. information on EDTA structure and function of EDTA in titrations.
If you find these useful or spot a silly error please EMAIL query? comment
The non-redox titration Questions
I've tried to quote the data to the appropriate significant figures and associated 'trailing zeros'.
Q1 A solution of sodium hydroxide contained 0.250 mol dm-3. Using phenolphthalein indicator, titration of 25.0 cm3 of this solution required 22.5 cm3 of a hydrochloric acid solution for complete neutralisation.
Q2 A solution made from pure barium hydroxide contained 2.74 g in exactly 100 cm3 of water. Using phenolphthalein indicator, titration of 20.0 cm3 of this solution required 18.7 cm3 of a hydrochloric acid solution for complete neutralisation. [atomic masses: Ba = 137, O = 16, H = 1)
Q3 4.90g of pure sulphuric acid was dissolved in water, the resulting total volume was 200 cm3. 20.7 cm3 of this solution was found on titration, to completely neutralise 10.0 cm3 of a sodium hydroxide solution. [atomic masses: S = 32, O = 16, H = 1)
Q4 100 cm3 of a magnesium hydroxide solution required 4.5 cm3 of sulphuric acid (of concentration 0.100 mol dm-3) for complete neutralisation. [atomic masses: Mg = 24.3, O = 16, H = 1)
Q5 Magnesium oxide is not very soluble in water, and is difficult to titrate directly.
Q6 2.00 dm3 of concentrated hydrochloric acid (10.0 M) was spilt onto a laboratory floor. It can be neutralised with limestone powder. [atomic masses: Ca = 40, C = 12, O = 16)
Q7 A 50.0 cm3 sample of sulphuric acid was diluted to 1.00 dm3. A sample of the diluted sulphuric acid was analysed by titrating with aqueous sodium hydroxide. In the titration, 25.0 cm3 of 1.00 mol dm-3 aqueous sodium hydroxide required 20.0 cm3 of the diluted sulphuric acid for neutralisation.
Q8 A sample of sodium hydrogencarbonate was tested for purity using the following method. 0.400g of the solid was dissolved in 100 cm3 of water and titrated with 0.200 mol dm-3 hydrochloric acid using methyl orange indicator.
23.75 cm3 of acid was required for complete neutralisation. [Ar's: Na = 23, H = 1, C = 12, O = 16]
Q9 This question involves theoretical calculations to do with 'how much to weigh out' for titrations and a common requirement to show development in coursework projects. They involve reagents such as pure anhydrous sodium carbonate, standardised hydrochloric acid and EDTA titrations (theory).
Atomic masses: O = 16, H = 1, Na = 23, C = 12, Ca = 40, P = 31.0
9(a)(i) Write out the equation, complete with state symbols for the reaction between hydrochloric acid and sodium carbonate.
9(b)(i) The simplified molecular structure of 2-ethanoylhydroxybenzoic acid ('Aspirin') is CH3COOC6H4COOH.
9(c) Pure calcium carbonate can be used to make a standard calcium ion solution to practice a complexometric titration of calcium ions with EDTA or determine the molarity of the EDTA reagent.
See Appendix 1. for theoretical information on EDTA structure and function in titrations (advisable to read).
Q10 25.0 cm3 of seawater was diluted to 250 cm3 in a graduated volumetric flask.
Q11 0.12 g of rock salt was dissolved in water and titrated with 0.100 mol dm-3 silver nitrate until the first permanent brown precipitate of silver chromate is seen.
Q12 5.00 g of a solid mixture of anhydrous calcium chloride(CaCl2) and sodium nitrate (NaNO3) was dissolved in 250 cm3 of deionised water in a graduated volumetric flask. A 25.0 cm3 aliquot of the solution was pipetted into a conical flask and a few drops of potassium chromate(VI) indicator solution was added.
Q13 A bulk solution of hydrochloric acid was standardised using pure anhydrous sodium carbonate (Na2CO3, a primary standard).
13.25 g of sodium carbonate was dissolved in about 150.0 cm3 of deionised water in a beaker.
The solution was then transferred, with appropriate washings, into a graduated flask, and the volume of water made up to 250 cm3, and thoroughly shaken (with stopper on!) to ensure complete mixing.
Q14 For this question, the relevant formula mass and equation are in the answers to Q13.
Q15 (a) Describe a procedure that can used to determine the molecular mass of an organic acid by titration with standardised sodium hydroxide solution. Indicate any points of the procedure that help obtain an accurate result and explain your choice of indicator.
0.279g of an organic monobasic aromatic carboxylic acid, containing only the elements C, H and O, was dissolved in aqueous ethanol. A few drops of phenolphthalein indicator were added and the mixture titrated with 0.100 mol dm-3 sodium hydroxide solution.
It took 20.5 cm3 of the alkali to obtain the first permanent pink. [at. masses: C = 12, H = 1 and O = 16]
Q16 Using the method outlined in the answer to Q15(a), 0.103g of a dibasic/diprotic non-aromatic carboxylic acid required 19.85 cm3 of a standardised sodium hydroxide solution for complete neutralisation. If the concentration of the alkali was 0.0995 mol dm-3. [at. masses: C = 12, H = 1 and O = 16]
Q17 The % purity of an organic acid can be determined by the procedure outlined in the answer to Q15(a).
Q18 Using the method described in the answer to Q15(a), sodium hydroxide solution can be standardised. 0.250 g of very pure benzoic acid (C6H5COOH) was titrated with a solution of sodium hydroxide of unknown molarity. If 22.5 cm3 of the alkali was required for neutralisation, calculate ...
Q19 The solubility of calcium hydroxide in water can be measured reasonably accurately to 3sf by titrating the saturated solution with standard hydrochloric acid.
Q20 ASPIRIN ASSAY ANALYSIS This question follows on in some respects from Q9b which I'd forgotten I'd already written, apologies for some repetition!
2-ethanoylhydroxybenzoic acid (acetylsalicylic acid), known commercially as aspirin, can be analysed by titration with standard sodium hydroxide solution when a sample of it is dissolved aqueous alcohol (a mixture of ethanol and water) and using phenolphthalein indicator (pKind = 9.3, useful range pH 8.3-10). In the pharmaceutical industry, aspirin is manufactured by reacting 2-hydroxybenzoic acid (salicylic acid) with ethanoic anhydride. Prior to this reaction, 2-hydroxybenzoic acid is manufactured by reacting carbon dioxide with phenol, the mixture is heated under pressure sodium hydroxide in the so called Kolbe Reaction. Aspirin, therefore, always contains a small percentage of 2-hydroxybenzoic acid as an impurity!
(a) Give the equation for the Kolbe synthesis of 2-hydroxybenzoic acid.
(b) Give the equation for the formation of aspirin from 2-hydroxybenzoic acid.
(c) Give the molecular formulae and calculate the molecular masses of 2-hydroxybenzoic acid and aspirin.
(d) Why must ethanol be added to the water prior to doing the titration?
(e) Five samples of aspirin were titrated with commercially purchased precisely 0.1000 mol dm-3 (0.1000M) sodium hydroxide solution and the results are given below.
In each case calculate the titre/mass and work out its average value for the five titrations.
(f) Assuming that only aspirin was titrated (though not true), from the average titre/mass figure calculate the 'theoretical' % purity of the aspirin by the following sequence:
(g) Why is the theoretical % purity based on this titration method always likely to be over 100%?, ignoring any titration errors - which does not necessarily explain why via this method of analysis you will always tend to get >100%, especially if you do the titration very accurately!
(h) Assuming that 2-hydroxybenzoic acid is titrated with NaOH on a 1 : 1 molar basis, calculate the % of this impurity in the aspirin by the follow sequence:
(i) Suppose for the sake of argument, there was an error of 0.1 cm3 on the titration value which is likely to be the biggest source of error. Obviously there are errors associated with the NaOH molarity, the weighing, burette reading.
Appendix 1. Information on EDTA structure and function
EDTA is an acronym abbreviation for the old name EthyleneDiamineTetraAcetic acid and is used in equations.
It is a hexadentate ligand i.e. it can donate 6 electron pairs to form 6 dative-covalent bonds and binds strongly with many metal cations Mn+ where n is usually 2 or 3.
Multi-dentate ligands are called chelating* agents, because the two ligand bonds become part of a five membered ring system. (*The chelate is from the Greek word meaning a crab's claw)
The solid EDTA 'off the shelf' used in analysis is usually the disodium dihydrate salt, which has the structure
The full unionised structure is (HOOCCH2)2NCH2CH2N(CH2COOH)2 which we could abbreviate to H4EDTA since theoretically four hydrogens from the four carboxylic acid groups are ionisable.
of which H2EDTA2- is the most prominent chelating species in solutions of pH10 in titrating calcium ions though the complex is actually formed by the combination of a metal ion and the EDTA4- ion.
The theory behind the titration of calcium ions with EDTA reagent is a bit complicated and the titration should be carried out in the presence of magnesium ions, usually included in the EDTA volumetric reagent, but if not, they must be in the mixture being titrated. This may seem to prelude an incorrect titration for calcium since magnesium ions reacting with EDTA, but it doesn't (see explanation later).
Hydrated/hexa-aqua metal ions like aqueous calcium and magnesium ions (M2+) give the following reaction with EDTA reagent,
which, for theoretical explanation of the titration and subsequent calculations, is best simplified to
Both calcium and magnesium EDTA complexes are strongly formed i.e. virtually 100% to the right BUT the Kstab for the formation of the EDTA-calcium ion complex is greater than that for the EDTA-magnesium ion complex i.e. the calcium ion complex is more stable and calcium ions will displace magnesium ions from their EDTA complex.
The indicators used e.g. Eriochrome Black T (represented in a 'free' anionic form as HIn2-) weakly complexes with ions such as the magnesium ion.
Both metal ions form a weak complex with the indicator at the start of the titration, but the indicator is displaced by the stronger binding EDTA (ligand displacement reactions), but much more slowly from the calcium complex than the magnesium complex i.e.
and this means without the presence of magnesium ions, the end-point is sluggish giving an inaccurate with just the calcium ions present, because (eq 6) is too slow.
Therefore the order of complex ion stability is [CaEDTA]2-(aq) > [MgEDTA]2-(aq) > [MgIn]-(aq) and this order of stability is crucial to the success of the titration as the ensuing argument will show.
In the EDTA solution are the Mg-EDTA complex ion plus excess uncomplexed EDTA ions. As the EDTA reagent is run into the calcium ion solution the calcium ion-EDTA complex is formed by reactions (eq 7) or (eq 8).
In (eq 8) the magnesium ion is displaced from its EDTA complex on a 1 : 1 molar basis by the calcium ion and then the free magnesium ions form a red complex with the blue indicator (eq 4) below. This continues as long as there are still Ca2+ ions to titrate and magnesium ions to be displaced i.e. no blue colour is seen yet.
Continued addition of EDTA eventually converts all the 'free' calcium ions into their EDTA complex ion via (eq 7 or 8), BUT, the 1st drop of excess EDTA after the calcium ions are all complexed then releases the blue form of the indicator via the fast reaction (eq 5),
so giving the sharp end point from red to blue at pH10.
In the case of analysing a mixture of calcium and magnesium ions in the same mixture, one method is to analyse for Ca2+ as above (VCa). Obviously, you do NOT add magnesium ions to the EDTA or the mixture being titrated if you wish to estimate the total (Mg2+ + Ca2+). You add a known excess of standardised EDTA solution (Vexcess) and then back titrate with another standardized M2+ ion solution (Vback) and the end point is blue to red.
I DO MY BEST TO CHECK MY CALCULATIONS, as you yourself should do, BUT I AM HUMAN! AND IF YOU THINK THERE IS A 'TYPO' or CALCULATION ERROR PLEASE EMAIL ME ASAP TO SORT IT OUT!
ILLUSTRATIONS OF ACID-ALKALI TITRATIONS and SIMPLE STARTER CALCULATIONS for A level students, but originally written for GCSE/IGCSE/O level students
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