1.9 Titrations

Section A

For each of the questions only ONE of the lettered responses (A-D) is correct.

Select the correct response in each case and mark its code letter by connecting the dots as illustrated on the answer sheet.

1. A sample of 11.6 g of hydrated sodium carbonate Na₂C0₃.xH₂0 was dissolved in water and the solution made up to 1000cm³. It was found that 25.0 cm³ of the solution required 29.6 cm³ of 0.05 mol dm^-3 sulphuric acid for neutralisation. Which one of the following is the value of x?

A 2

B 5

C 8

D 10

2. Which one of the Indicators listed is appropriate for the titration stated?

0.1 M base 0.1 M acid Indicator

A sodium carbonate hydrochloric acid phenolphthalein

B ammonia solution hydrochloric acid phenolphthalein

C sodium hydroxide sulphuric acid methyl orange

D ammonia solution ethanoic acid solution methyl orange

3. A solution of sodium sulphide, Na₂S, contains 7.8 g/l. What is the concentration, in mol/l of sodium ions?

A 0.05

B 0.10

C 0.20

D 1.00

4. In order to determine the concentration of an aqueous solution of a weak acid, a portion of the solution was pipetted into a conical flask and titrated against aqueous sodium hydroxide, using phenolphthalein as indicator. In a subsequent experiment, the concentration of the acid found in this way was found to be too low.

Which one of the following could explain this low value?

A Overshooting the end point of the titration.

B Rinsing the conical flask with distilled water immediately before pipetting the acid into it.

C Rinsing the pipette with distilled water immediately before filling it with acid.

D Rinsing the burette with distilled water immediately before filling it with the aqueous sodium hydroxide.

5. When 0.3 g of an impure sample of ammonium sulphate ,(NH₄)₂SO₄, was boiled with an excess of aqueous sodium hydroxide, the ammonia evolved required 15 cm³ of 0.10 M hydrochloric acid for neutralization.

Assuming that none of the impurities reacted with the sodium hydroxide. Which one of the following is the percentage purity of the ammonium sulphate?

A 5.0

B 16.5

C 33.0

D 66.0

6. 25 cm³ of potassium iodate (V) solution were added to excess potassium iodide solution dissolved in sulphuric acid. The iodine liberated required 30 cm³ of 0. 50 M Na₂S₂0₃. Which one of the following is the concentration of the potassium iodate solution?

A 0.1 M

B 0.2 M

C 0.4 M

D 0.5 M

Section B

1. Sodium hydroxide absorbs water and carbon dioxide from the atmosphere. Hence a standard solution of sodium hydroxide cannot be prepared by direct weighing. Instead, a solution of sodium hydroxide is prepared and titrated with the weak acid, oxalic acid, as the following method describes:

Weigh out about 1.2 g of sodium hydroxide pellets, place in a beaker, dissolve them in distilled water and make up to 250 cm³ in a standard flask. Weigh out accurately 1.600 g of oxalic acid dihydrate crystals, dissolve them in water and transfer to a 250 cm³ standard flask. Rinse the original container, add the washings to the flask and make up to the mark. Titrate 25.0 cm³ portions of the acid solution with the alkali.

The reaction taking place is:

2NaOH + (COOH)₂.2H2O (COONa)₂ + 4H₂O

COOH

oxalic acid

(a) Explain the following terms.

(i) Standard solution

[1]

(ii) Dihydrate

[1]

(iii) Weak acid

[1]

(b) Draw the ions present in sodium oxalate. The covalent bonds should not be shown.

` [2]

(i) State a suitable indicator for the titration.

[1]

(ii) Describe the colour change taking place in the flask.

[2]

(d) The volume of alkali needed to neutralise the acid is 26.0 cm³. Calculate the molarity of the standard sodium hydroxide solution.

mol dm^-3 [4]

2. Commercial concentrated sulphuric acid is not pure H₂SO₄; a small amount of water is present. If the concentrated acid is added to water a diluted solution is produced in an exothermic process.

The percentage of H₂SO₄ in the commercial acid may be determined as follows:

Accurately weigh a small glass stoppered weighing bottle and add from a clean dry measuring cylinder 0.7-0.8 cm³ of concentrated sulphuric acid. Re-stopper the bottle and reweigh accurately. Place about 100 cm³ of water in a 250 cm³ volumetric flask.

Wash the acid into the flask. Dilute nearly to the mark and after about an hour make the solution up to the mark. Titrate three 25.0 cm³ portions against 0. 1 M sodium hydroxide.

(a) At room temperature the density of concentrated sulphuric acid is

1.84 g cm ^–3.

(i) Suggest why the acid is weighed rather than using its volume and density to determine its mass.

[1]

(ii) Suggest why the diluted sulphuric acid is left for about one hour before it is made up to the mark in the volumetric flask.

[1]

(b) Write the equation for the reaction of sulphuric acid with sodium hydroxide.

[2]

[1]

(ii) Explain why three titrations are carried out.

[1]

(d) Calculate the percentage of H₂SO₄ in the commercial acid from the following results:

mass of weighing bottle = 10.124 g

mass of weighing bottle + sulphuric acid = 11.415 g

initial reading/cm³ final reading/cm³

first titration 1.2 27.9

second titration 10.1 35.9

third titration 20.5 46.4

% [5]

4. In a textbook, the instructions for an experiment to find the value of x in the formula Na₂CO₃.xH₂O in a sample of hydrated sodium carbonate crystals were as follows:

Weigh out accurately about 5 grams of soda crystals, being careful to select well-defined translucent crystals rather than those covered in powder. Dissolve these in a beaker of water, transfer to a 250 cm³ flask, shake well and make up to the mark. Pipette 25.0 cm³ of this solution into a conical flask, add a few drops of methyl orange and titrate with standard 0.2 M hydrochloric acid until the first permanent colour change is observed. Repeat to obtain two results within 0. 1 cm³ of each other.

A student who follows these instructions weighs out 4.90 g of the crystals and finds that an accurate titre of 17.7 cm³ is required for neutralisation.

(a) What is meant by standard 0.2 M hydrochloric acid?

[1]

(b) (i) What colour change will be observed with the methyl orange?

[2]

(ii) Why did the instructions suggest the use of translucent crystals rather than those covered in powder?

[1]

(ii) Write an equation for the reaction between the hydrochloric acid and the sodium carbonate solution.

[1]

(iii) Calculate the concentration of the sodium carbonate solution.

[1]

(iv) Calculate the value of x.

[2]

5. Hydrogen peroxide has largely replaced chlorine as an oxidising agent

in the paper industry.

The concentration of solutions of hydrogen peroxide may be found by oxidising excess potassium iodide solution and titrating the iodine produced with sodium thiosulphate solution. The redox half-equations are:

H₂O₂ + 2H⁺ + 2e^- 2H₂O

2I^- I₂ + 2e^-

If 10 cm³ of a solution of hydrogen peroxide is made up with distilled water to 1 dm³ and a 25 cm³ sample reacted with excess iodide, it is found that 30 cm³ of 0.2 mol dm^-3 sodium thiosulphate solution is needed to react completely with the iodine produced.

(a) Name the indicator used to make the end-point clearer and describe the colour change.

Indicator [1]

Colour change from to [2]

(b) Calculate the molarity of the original hydrogen peroxide solution.

Molarity [4]

(c) Hydrogen peroxide decomposes according to the equation:

2H₂O₂ 2H₂O + O₂

Calculate the volume of oxygen, at 20^oC and 1 atm pressure, which would be produced from the complete decomposition of 1 0 cm³ of the original hydrogen peroxide solution.

[2]

(i) An organic acid, butanoic acid, is found in rancid butter and sweat and also has an unpleasant smell. It is a soluble, weak acid similar to ethanoic acid (vinegar). Assuming all the glassware is clean and starting from a standard 0. 1 mol dm^-3 solution of sodium hydroxide, suggest how the concentration of an aqueous solution of butanoic acid could be found.

[6]

(ii) Suggest a suitable indicator.

[1]

6. Ethanoic acid may be prepared by the following method.

Concentrated sulphuric acid (6 cm³) is added to water (1 0 cm³). Sodium dichromate, Na₂Cr₂O₇, (10 g) is added to this solution together with some "boiling chips". The flask is fitted with a reflux condenser and the solution is boiled. A mixture of ethanol (4 cm³) and water (10 cm³) is added, a little at a time, to the reaction mixture through the condenser. When all the aqueous ethanol has been added the flask is refluxed for 30 minutes. The ethanoic acid is obtained by distillation collecting the fraction boiling between 115 and 120 ^oC.

(a) In this preparation sodium dichromate is reduced to chromium(III).

(i) Deduce the oxidation number of chromium in sodium dichromate.

[1]

(ii) Balance the half equation for the reduction of dichromate to chromium(III) ions.

Cr₂O₇^2- + H⁺ Cr³⁺ + H₂O [2]

(b) Sodium dichromate oxidises the ethanol to ethanoic acid.

CH₃CH₂OH + 2[O] CH₃COOH + H₂O

The distillate from the reaction is a mixture of ethanoic acid (40% by mass) and water. Calculate the percentage yield of ethanoic acid if the density of ethanol is 0.79 g cm^-3 and 10 g of distillate is obtained.

[4]

(i) Calculate the percentage of ethanoic acid in vinegar if a 10 cm³ sample required 5.60 cm³ of 1.0 M sodium hydroxide for complete neutralisation.

(Assume the density of vinegar is 1.0 g cm^-3)

[3]

(ii) Giving practical details outline how you would titrate a sample of vinegar stating the indicator used.

[6]

(d) Ethanoic acid has a carboxylic acid group, -COOH, which contains a double bond between the carbon atom and one of the oxygen atoms.

(i) Using the outer electrons only, draw a dot and cross diagram for the -COOH group.

[2]

(ii) The double bond is polarised. State the polarity of the bond and use the concept of electronegativity to explain your answer.

[2]

7. The concentration of the iodate (V) ion in solution can be determined by reaction with excess acidified, potassium iodide solution and titrating the liberated iodine with standard sodium thiosulphate solution.

KIO₃ + 5KI + 6HCl 3I₂ + 3H₂0 + 6KCl

I₂+ 2Na2S₂0₃ 2NaI + Na₂S₄0₆

(a) The standard t