1 An experiment is carried out to measure Ka for ethanoic acid, CH3CO2H. 100 cm3

of sodium ethanoate solution of concentration 1.0 mol dm -3 is placed into a beaker and 20 cm3 of hydrochloric acid of concentration 1.0 mol dm-3 is added followed in turn by the addition of one drop of universal indicator solution. The colour obtained is matched with the colour produced in a series of solutions of known pH. Three further 20 cm3 additions of the same acid are made and, after each addition, the pH is determined by the same technique.

The results obtained are:

After first addition of acid pH 5.0After second addition of acid pH 4.9After third addition of acid pH 4.6After fourth addition of acid pH 4.3

(a) Write an equation for the reaction between sodium ethanoate and hydrochloric acid.

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(b) With reference to the solution formed in the experiment, explain what is meant by a buffer solution.

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(c) Calculate for each of the four additions a value for the acid dissociation constant, Ka, for ethanoic acid.

(d) Suggest the limitations of this method for determining Ka for a weak acid.

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Chemistry H2 9746Tutor TuteeRevision Exercise 19: Integrated Questions

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(e) Hence, state the value of Ka you would quote as a result of this experiment.

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Practice in Chemistry Progressive Questions for AS and A level Pg 74 (22.17)2 An oxide of element of Group I, M2O, contains 83.0% of the metal. On further

heating in oxygen M2O reacts to give another oxide Y, which contains 71.0% of the metal.

(a)(i)

Calculate the relative atomic mass of M and hence identify it.

(ii) Write down the electronic configuration of M and use it to explain why M forms only a +1 ion.

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(b) Write down the equation for the first ionisation energy of M.

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(c)(i)

With the aid of an equation, describe the reaction of M with water.

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(ii) Explain why M reacts less readily with water than the element below it in Group I.

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(d)(i)

Calculate the empirical formula of the oxide Y and write down the formula of the ions which it contains.

(ii) Suggest why the oxide of lithium corresponding to Y is unstable compared to those of the other Group I metals.

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Practice in Chemistry Progressive Questions for AS and A level Pg 111 Q43 Nitrite ions can be determined quantitatively by titration with manganate(VII)

ions in acidic solution, according to the equation:

2MnO4– + 5NO2

– + 6H+ 2Mn2+ + 3H2O + 5NO3–

(a) Iron(II) ammonium sulfate (FeSO4.(NH4)2SO4.6H2O, molar mass = 392 g mol-1) can be used to standardise the solution of manganate(VII) ions.

(i) Write an ionic equation for the reduction of manganate(VII) ions to Mn2+ by iron(II) in acidic solution.

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(ii) Calculate the mass of iron(II) ammonium sulphate needed to prepare 250 cm3

of a solution of concentration 0.1 mol dm-3.

(b) Write down the two half equations for the overall reaction between manganate(VII) ions and nitrite ions in acidic solution.

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(c) In a typical experiment to determine the concentration of nitrite ions, 25.0 cm3

of 0.02 mol dm-3 solution of potassium manganate(VII) was acidified, heated to about 40oC and then titrated with a solution of sodium nitrite, of which 25.0 cm 3

were required to reach end-point.

(i) How might the end-point of the titration be most conveniently determined?

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(ii) Calculate the concentration, in mol dm-3, of nitrite ions in the solution.

(d) The aqueous Mn3+ ion is as powerful oxidising agent as manganate(VII), but it is rarely used because it readily disproportionate into solid MnO2 and Mn2+ ions.

(i) With reference to Mn3+, explain what is meant by disproportionation.

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(ii) Write a balanced equation for the disproportionation of the Mn3+ ion into MnO2

and Mn2+.

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(iii) State and explain how the tendency of the Mn3+ ion to disproportionate would be affected by changes in the pH of the reaction mixture.

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Practice in Chemistry Progressive Questions for AS and A level Pg 113 Q94 The use of the Data Booklet is relevant to this question.

Explain the following observations, giving equations where relevant.

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(a) The addition of aqueous sodium hydroxide to copper(II) sulphate solution forms a blue precipitate which is insoluble in excess sodium hydroxide, but dissolves in excess aqueous ammonia to give a deep-blue solution.

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(b) Iron wool glows when heated in dry chlorine, forming a blackish-green product. It behaves similarly when heated in hydrogen chloride gas, but the product is white. In the presence of water the products from both reactions have colours different from those obtained when no water is present.

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(c) Iron is protected from corrosion when placed in contact with a piece of zinc, but it corrodes more rapidly when placed in contact with a piece of tin.

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(d) Three chromium(III) complexes A, B and C having the molecular formula CrCl3(H2O)6 can be prepared. The results of adding aqueous silver nitrate to their aqueous solutions are described below. There are 2 possible structures for the cation in complex A.

Complex Number of moles of AgCl immediately precipitated per

mole of complexA 1B 2C 3

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(e) The number of stable oxidation states shown by the d-block elements increases from scandium to manganese but decreases markedly thereafter.

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(f) Manganese and iron both react readily with acids to form solutions of their respective +2 ions but manganese is the more powerful reducing agent.

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Synoptic Questions Q23 and Q245 The following question concerns the d-block elements, scandium to zinc.

(a) The highest oxidation state of the first six d-block elements are shown in the table below.

Element Sc Ti V Cr Mn FeHighest oxidation state

+3 +4 +5 +6 +7 +6

Comment on the factors affecting the range of oxidation states shown by these elements. Suggest a value for the highest oxidation state for cobalt, explaining your answer.

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(b) With reference to the Data Booklet, comment on and explain the stabilities of the Fe3+ and Mn3+ in aqueous solution.

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(c) A sample of steel (1.00 g) containing manganese was dissolved in nitric acid to give a solution containing manganese in oxidation state +2. All the manganese present was then oxidised to manganate(VII) by adding sodium bismuthate, NaBiO3. After the destruction of any excess bismuthate ion, the resulting purple solution required 36.0 cm3 of an iron(II) sulfate solution of concentration 0.10 mol dm-3 to reach an end point, the iron(II) being oxidised to iron(III).

(i) State the oxidation state of bismuth in NaBiO3.

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(ii) Given that the bismuthate ion is reduced to Bi3+, write an equation for the reaction in which it functions as an oxidising agent. Hence, write a balanced equation for the oxidation of manganese(II) to Mn(VII) by bismuthate in acidic solution.

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(iii) Suggest why the oxidation was carried out in acidic solution.

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(iv) Calculate the percentage by mass of manganese in the steel sample.

Synoptic Questions Q326 Borax, sodium borate Na2B4O7.10H2O, is, according to some chemists, the most

reliable substance with which to standardise a solution of an acid. It reacts with acid forming boric acid, which is so weak an acid that virtually none of it is dissociated into ions and therefore does not interfere with the titration if the correct indicator is chosen. This reaction can be summarised by the following equation:

B4O72- + 2H+ + 5H2O 4H3BO3

1.15 g of borax was dissolved in warm distilled water and made up to 100 cm3. 10.0 cm3 portions of borax solution were titrated with the sulfuric acid whose concentration has to be determined accurately. 20.0 cm3 of acid was required.

(a) Calculate the concentration of the borax solution in mol dm-3.

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(b) Calculate the concentration of sulphuric acid.

Synoptic Questions Q507 A geologist wanted to determine the percentage purity of a sample of limestone

(calcium carbonate) by reacting it with sulphamic acid, H3SO3N. However, he did not know how many hydrogen atoms had acidic properties and so decided to investigate this by reacting sulphamic acid with sodium carbonate solution.

He made up a solution of sulphamic acid containing 1.50 g in 250 cm3 of solution and titrated a 20.0 cm3 sample with 0.025 mol dm-3 sodium carbonate using a suitable indicator. 24.8 cm3 of sodium carbonate solution was required for complete reaction.

Calculate the number of moles of replaceable hydrogen atoms in one mole of sulphamic acid.

8 The substance known as ‘microcosmic salt’ is hydrated ammonium sodium hydrogen phosphate and has the formula NH4NaHPO4.xH2O. In order to find the value of x, a student weighed out 10.0 g of microcosmic salt and dissolved it in distilled water in 100 cm3 volumetric flask before making it up to the mark. She then took 20.0 cm3 of this solution and added the same volume of 1.0 mol dm -3

sodium hydroxide. The mixture was boiled gently until no more ammonia vapour could be detected. After cooling, the mixture was titrated with 0.50 mol dm-3 hydrochloric acid and the process was repated and the average of two concordant titres was obtained – 14.2 cm3.

(a) State a suitable indicator which can be used for the titration. Explain your

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answer.

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(b) Calculate the value of x.

Synoptic Questions Q53 9 10 cm3 of a saturated solution of potassium bromide was added to 30 cm3 of a

solution of copper(II) sulfate of concentration 0.20 mol dm-3. It turned from blue to green. The reaction involved was:

CuSO4(aq) + 4KBr(aq) ⇌ K2[CuBr4](aq) + K2SO4(aq)

(a) Explain the following observations.

(i) Solid sodium sulfate, Na2SO4, was added to 10 cm3 of the green solution. It turned blue.

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(ii) The remaining solution was then left to stand for some time until it became blue-green and was then divided into three portions A, B and C.

Portion A was warmed and it turned blue.Portion B was placed in an ice-bath and it turned green.Portion C was left at room temperature and it stayed blue-green.

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(b) From the results described in (a)(ii), deduce the nature of the energy change for the forward reaction shown in the equation above.

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(c) What colour changes, if any, would you expect to see if

(i) A large volume of sodium chloride solution was added to portion A after it had been warmed? Explain your answer.

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(ii) The solution was subsequently diluted? Explain your answer.

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Synoptic Questions Q5810 An aqueous solution of cobalt(II) chloride is pink; a result of its containing the

[Co(H2O)6]2+(aq) ion. On the addition of concentrated hydrochloric acid, the solution turns blue as a result of the reaction:

[Co(H2O)6]2+(aq) + 4Cl-(aq) ⇌ [CoCl4]2-(aq) + 6H2O(l)

Use is made of this product in colouring the drying agent, silica gel, which is used, for example, in desiccators. When the drying agent is active, it is blue. When it is saturated with water the colour changes to pink. It is then concerted back to blue on heating in an oven at 120oC. By discussing the effect of concentration changes on the essential reaction involved, explain these reactions.

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Synoptic Questions Q5911 100 cm3 of saturated calcium sulfate solution was passed through an ion

exchange column, which replaces Ca2+(aq) ions with H+(aq) ions as follows.

Ca2+(aq) + H2(resin)(s) 2H+(aq) + Ca(resin)(s)

The acid displaced required 20.0 cm3 of 0.05 mol dm-3 sodium hydroxide solution for complete neutralisation.

(a) Calculate the number of moles of H+(aq) that was displaced by the Ca2+(aq) ions.

(b) Calculate the number of moles of Ca2+(aq) ions that was present in the original calcium sulfate solution.

(c) What is the molar concentration of the calcium sulfate solution used in the experiment?

(d) Write an expression for the solubility product of calcium sulfate and calculate its value under the conditions of this experiment.

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(e) If the original solution is acidified with sulfuric acid, a white precipitate is formed. Name the solid which separates and explain why no precipitate is formed if hydrochloric acid is used in place of the sulphuric acid.

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Chemistry in Focus Pg 31 Q413 The concentrations of iron(II) and ethanedioate ions in solution can be

determined by titration against cerium(IV) ions. The ionic half equations are:

Fe2+(aq) ⇌ Fe3+(aq) + e-

C2O42-(aq) ⇌ 2CO2(g) + 2e-

Ce4+(aq) + e- ⇌ Ce3+(aq)

50.0 cm3 of 0.2 mol dm-3 cerium(IV) sulfate solution was added to 25 cm3 of a solution of iron(II) ethanedioate, FeC2O4. An excess of cerium(IV) ions remained and these were reduced to cerium(III) ions by titration with 0.1 mol dm -3 sodium ethanedioate solution, of which 30 cm3 were required.

(a) To start with, 250 cm3 of the 0.2 mol dm-3 cerium(IV) sulfate solution was prepared. Calculate the mass of solid cerium(IV) sulfate needed to make up the solution.

(b) Write down the balanced equations for the reaction of

(i) Cerium(IV) ions with ethandioate ions,

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(ii) Cerium(IV) ions with iron(II) ions.

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(c) Calculate the number of moles of ethanedioate ions needed to reduce the excess cerium(IV) ions.

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(d) How many moles of cerium(IV) ions were added to the iron ethanedioate solution?

(e) Hence, calculate the number of moles of cerium(IV) ions needed to react with iron ethanedioate.

(f) Calculate the concentration of iron(II) ethanedioate solution used in the experiment.

Chemistry in Focus Pg 37 Q1114(a)

The acid dissociation constant, Ka for ethanoic acid is 1.80 × 10–5 mol dm–3.

(i) Write an equation for the dissociation of aqueous ethanoic acid.

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(ii) Give the expression for the acid dissociation constant, Ka, for aqueous ethanoic acid.

(iii) Calculate the pH of 2.00 mol dm–3 ethanoic acid.

(b) When ethanoic acid and propan-1-ol are heated together in the presence of a little concentrated sulphuric acid an ester is produced.

(i) Draw the structural formula of the ester produced in this reaction.

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(ii) State the function of the sulphuric acid.

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(c) 6.00 g of ethanoic acid was mixed with 6.00 g of propan-1-ol and 0.00100 mol of concentrated sulphuric acid was added. The mixture was heated under reflux until no further reaction was detectable. The mixture was then cooled rapidly to room temperature and titrated with 1.00 mol dm–3 sodium hydroxide solution. 35.0 cm3 of the 1.00 mol dm–3 sodium hydroxide solution was required for complete neutralisation.

(i) Calculate the volume of sodium hydroxide solution that would have been needed if only the 0.00100 mol of concentrated sulphuric acid had been in the flask.

(ii) Calculate the volume of sodium hydroxide solution that must have been needed to react with the ethanoic acid remaining in the cooled mixture.

(iii) Calculate the amount (in moles) of ethanoic acid remaining.

(iv) Hence, calculate the amount (in moles) of propan-1-ol, ester and water in the final cooled mixture.

(v) Calculate the value of the equilibrium constant under these conditions. The volume of the reaction mixture should be taken as 100 cm3.

(d) The reverse of esterification is hydrolysis. In experiments to study the hydrolysis of this ester with aqueous sodium hydroxide, the following data were

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obtained.

Experiment [ester] / mol dm-3

[NaOH] / mol dm-3

Rate / mol dm-3

s-1

1 0.1 1.0 6.0 x 10-4

2 0.1 2.0 1.2 x 10-3

3 0.2 2.0 2.4 x 10-3

(i) Determine the order of the reaction with respect to [ester] and [NaOH]

(ii) Calculate the value of the rate constant, k, for this reaction, including its units.

A2 Chemistry rates a2.pdf Q5End of Paper

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