Year 12 Chemistry Stage 3
CACULATIONS ASSIGNMENT
Please read the following:

Complete these calculations. Correct answers but not worked solutions will be available.

The Validation Test will contain 4 questions. Two from this assignment and two unseen calculations.

You will not be allowed to use your solutions in the Validation Test
1. One method of preparing chlorine gas in the laboratory is to react hypochlorous acid
(bleach) with hydrochloric acid. The following reaction occurs:
HOCl(aq) + HCl(aq) → H
2
O(l) + Cl
2
(g)
In one such experiment, 150.0 mL of 0.502 mol L -1 hydrochloric acid was added to 500.0 mL of a bleach solution which contained 4.00 g of HOCl per litre.
(a) Determine the limiting reagent in this reaction.
(b) Calculate the volume of chlorine gas obtainable from the above experiment if it is measured at 25 o C and 94.0 kPa.
2. In an industrial accident at a chemical factory, 2.00 x 10
5
L of hydrochloric acid solution spills into a pond which is holding 5.00 x 10
6
L of pure water. The pH of the pondwater is tested after the spillage and is found to be 0.716.
(a) Calculate the concentration of the original hydrochloric acid solution.
(b) It is decided to neutralise the hydrochloric acid in the pond by adding a suitable quantity of pure solid hydrated sodium carbonate (Na
2
CO
3
.10H
2
O) to the pond. Calculate the mass of the hydrated sodium carbonate needed.

3.
Iron is often found in the earth's crust as a hydrated iron oxide (Fe
2
O
3
.xH
2
O).
A 0.668 gram sample of this ore was dissolved in excess sulfuric acid and all the iron was converted into an iron(II)sulfate solution. This solution was titrated against a 0.050 mol L
-1 acidified potassium permanganate solution and exactly 25.00 mL of the purple solution was needed to complete the titration. Find the value of x in the hydrated formula above.
Equations:
Fe
2+ 
Fe
3+
+ e
-
MnO
4
-
+ 8H
+
+ 5e

Mn
2+
+ 4H
2
O
4. An analytical chemist is given the task of determining the percentage purity of a sample of potassium permanganate. She takes a representative sample of the compound weighing 2.350 g , dissolves it in water and makes the volume up to 250.0 mL in a volumetric flask. After thorough mixing, the permanganate solution is transferred to a clean dry burette ready for titration.
She then prepares a solution of oxalic acid by dissolving 5.15 g of pure H
2
C
2
O
4
.2H
2
O in water and making the volume up to 250.0 mL in another volumetric flask.
The chemist then takes four 20.0 ml aliquots of the oxalic acid solution in four separate conical flasks, adds sulfuric acid catalyst to each, warms the solutions to about 80 o
C and then performs four titrations by adding the permanganate solution to the conical flasks until the first permanent trace of pink
(excess MnO
4
-
(aq)) appears.
The results of the four titrations are shown in the Titration Table below:
Volume of KMnO
4
(aq) added to 20.0 mL oxalic acid
Titre 1 Titre 2 Titre 3 Titre 4
Initial reading (mL) 1.60 0.05 2.35 1.05
Final reading (mL) 25.80 22.50 24.85 23.45
Added volume (mL)
Average titre = .......................mL
From the above data, calculate the percentage purity of the potassium permanganate sample.

5. Analysing an organic compound 13 marks
A certain organic compound is known to contain only carbon, hydrogen and oxygen.
The compound was analysed as follows.

A 2.149 g sample was burned and the carbon dioxide produced was bubbled through a barium hydroxide solution, producing 11.27 g of barium carbonate
(BaCO
3
).
CO
2
+ Ba(OH)
2

BaCO
3
+ H
2
O

The mass of water produced by burning of the sample was 0.7721 g

The compound was found to have a molecular weight of 150.1 a) What is the empirical formula of the compound?
[ You may do this by finding the masses of carbon, hydrogen and oxygen in the sample ] b) What is the molecular formula of the compound? c) The compound is also known to be a carboxylic acid; that is, containing one COOH group. Write the molecular formula in the form of C
X
H
Y
O
Z
COOH (giving values for
X, Y and Z).
6. Sodium pyrosulfate (Na
2
S
2
O
7
) is produced by the following series of reactions:
S(s) + O
2
(g)  SO
2
(g)
2SO
2
(g) + O
2
(g)  2SO
3
(g)
SO
3
(g) + H
2
O(l)  H
2
SO
4
(aq)
H
2
SO
4
(aq) + Na
2
CO
3
(s)  2NaHSO
4
(aq) + H
2
O(l) + CO
2
(g)
2NaHSO
4
(aq)  Na
2
S
2
O
7
(s) + H
2
O(l)
In the commercial production of Na
2
S
2
O
7
17.5 Kg of sulphur is consumed in the production of
50 kg of Na
2
S
2
O
7
. Calculate the percentage efficiency of the process.

7. A sample of distilled and purified hydrogen peroxide solution was taken to a quality control laboratory to check the concentration of the hydrogen peroxide in the solution.
A 10.0 mL sample of the product was diluted to 500.0 mL in a volumetric flask.
Acidified aliquots of 10.0 mL of this diluted solution were then titrated against a standard 0.112 mol L
–1
potassium permanganate solution.
The burette readings obtained are shown in the table below.
Trials (mL)
Titration result
1 2 3 4
Final reading 19.32 37.73 18.84 37.54
Initial reading 0.03 18.98 0.14 18.84
Titre
(a) Write a balanced redox equation for the reaction between the hydrogen peroxide and permanganate ion.
(b) Determine the average titre value.
(c) Calculate the concentration of the hydrogen peroxide, in mol L
–1
, in the original sample
(d) What mass of hydrogen peroxide is in 100 mL of this original solution?
(e) What is the mass of 100 mL of the original peroxide solution if it has a density of
1.29 g mL
–1
?
(f) The concentration of hydrogen peroxide from the production process should be 70 % w/w (or 70 % by mass). This means that in 100 g of hydrogen peroxide solution there should be 70 g H
2
O
2
. From your answers to (d) and (e), calculate the concentration of the hydrogen peroxide solution as a % w/w and state whether the production process is operating appropriately.

8.
9.
An unknown compound contains sulphur, oxygen and chlorine. When mixed with water the compound yields a mixture of sulfuric and hydrochloric acids. A 0.504 g sample of the compound was dissolved in water and split into equal aliquots. The first aliquot, when treated with excess barium nitrate yielded 0.4671 g of barium sulfate.
The second aliquot was titrated with 0.250 mol L -1 NaOH solution and required 24.0 mL for complete neutralisation.
Determine the empirical formula of the compound.
When you suffer from an acid stomach the fluid in your stomach can contain HCl at a concentration of 2.5 x 10 -4 mol L -1 .
To reduce the concentration of acid an antacid tablet can be taken. If the tablet contains 5.00mg of aluminium hydroxide and 5.00 mg of magnesium hydroxide and the contents of your stomach has a volume of 1.50 L of fluid, calculate: a) the concentration of hydrogen ions b) the concentration of hydroxide ions c) the pH
10. A textile company uses a solution of sodium hypochlorite to bleach sheets. An industrial chemist employed by the company is asked to investigate the change in the concentration of the bleaching solution after sheets have been soaked in this solution for a time of 2 hours. The original solution is known to contain 30.0 g of sodium hypochlorite per litre of solution.
The chemist removes a sample of the bleaching solution after two hours of bleaching and tests it with sodium thiosulfate solution to determine the final concentration of hypochlorite ions. The reaction between hypochlorite ion and thiosulfate ion in an acidified solution may be represented as follows:
OCl (aq) + 2S
2
O
3
2(aq) + 2H + (aq)

S
4
O
6
2(aq) + Cl (aq) + H
2
O(l)
Using a starch-iodine indicator to moderate the equivalence point of the reaction, it is found that an average of 47.5 mL of 0.250 mol L -1 sodium thiosulfate solution is needed to react with 20.0 mL samples of the final bleaching solution.
(a) Calculate the concentration of sodium hypochlorite in mol L -1 in the original bleaching solution.
(b) Calculate, from the given chemical equation the concentration of hypochlorite ions in the final bleaching solution.
THE END