[Maximum number: 1]

Sample X is added to water and made up to a total volume of $200 \mathrm{~cm}^{3}$ . This gives a solution of $0.100 \mathrm{~mol} \mathrm{dm}^{-3} \mathrm{HCl}$ .

What is X ?

A

$10 \mathrm{~cm}^{3}$ of $1.00 \mathrm{~mol} \mathrm{dm}^{-3} \mathrm{HCl}$

B

$30 \mathrm{~cm}^{3}$ of $0.90 \mathrm{~mol} \mathrm{dm}^{-3} \mathrm{HCl}$

C

$50 \mathrm{~cm}^{3}$ of $0.40 \mathrm{~mol} \mathrm{dm}^{-3} \mathrm{HCl}$

D

$100 \mathrm{~cm}^{3}$ of $0.30 \mathrm{~mol} \mathrm{dm}^{-3} \mathrm{HCl}$

[Maximum number: 1]

Which contains the largest number of hydrogen atoms?

A

0.10 mol of pentane

B

0.20 mol of but-2-ene

C

1.00 mol of hydrogen molecules

D

$6.02 \times 10^{23}$ hydrogen atoms

[Maximum number: 3]

The elements in Group 17 are known as the halogens.

(a)

The concentration of NaClO in bleach S is $x \mathrm{gdm}^{-3}$ .

NaClO reacts with $\mathrm{H}_{2} \mathrm{O}_{2}(\mathrm{aq})$ as shown.

\mathrm{H}_{2} \mathrm{O}_{2}(\mathrm{aq})+\mathrm{NaClO}(\mathrm{aq}) \rightarrow \mathrm{H}_{2} \mathrm{O}(\mathrm{l})+\mathrm{NaCl}(\mathrm{aq})+\mathrm{O}_{2}(\mathrm{~g})

A $5.00 \mathrm{~cm}^{3}$ sample of S completely reacts with $\mathrm{H}_{2} \mathrm{O}_{2}(\mathrm{aq})$ . The volume of $\mathrm{O}_{2}(\mathrm{~g})$ produced is $24.0 \mathrm{~cm}^{3}$ under room conditions.

Assume that only the NaClO in S reacts with $\mathrm{H}_{2} \mathrm{O}_{2}(\mathrm{aq})$ .
Calculate x. Show your working.

x=

$\mathrm{g} \mathrm{dm}^{-3}$

[ 3 ]

[Maximum number: 1]

In the Periodic Table, the p block contains elements whose outer electrons are found in the p subshell.

(a)

Sulfur-containing compounds, such as $\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{SH}$ , are found in fossil fuels, and produce $\mathrm{SO}_{2}$ when they are burned.

[ 1 ]

(i)

Write the equation to show the complete combustion of $\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{SH}$ .

[ 1 ]

[Maximum number: 5]

A0.17 g sample of a Group 14 chloride, $\mathrm{XCl}_{4}$ , reacted with water to produce an oxide, $\mathrm{XO}_{2}$ , and HCl.
equation 1

\mathrm{XCl}_{4}(\mathrm{~s})+2 \mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \rightarrow \mathrm{XO}_{2}(\mathrm{~s})+4 \mathrm{HCl}(\mathrm{aq})

The HC lproduced was absorbed in $100 \mathrm{~cm}^{3}$ of $0.10 \mathrm{~mol} \mathrm{dm}^{-3}$ sodium hydroxide solution (an excess).
In a titration, the unreacted sodium hydroxide solution required $30.0 \mathrm{~cm}^{3}$ of $0.20 \mathrm{~mol} \mathrm{dm}^{-3}$ hydrochloric acid for complete neutralisation.

(a)

Calculate the amount, in moles, of hydrochloric acid used in the titration to neutralise the unreacted sodium hydroxide solution.
amount = mol

[ 1 ]

(b)

Calculate the amount, in moles, of sodium hydroxide neutralised in the titration.
amount = mol

[ 1 ]

(c)

Calculate the amount, in moles, of sodium hydroxide that reacted with the HCl produced by the reaction in equation 1 .
amount = mol

[ 1 ]

(d)

Calculate the amount, in moles, of HCl produced by the reaction in equation 1.
amount = mol

[ 1 ]

(e)

Calculate the amount, in moles, of $\mathrm{XCl}_{4}$ in the original 0.17 g sample.

amount =mol [1]

[ 1 ]

[Maximum number: 2]

An experiment was carried out to determine the percentage of iron in a sample of iron wire.

(a)

A3.35 g piece of the wire was reacted with dilute sulfuric acid, in the absence of air, so that all of the iron atoms were converted to iron(II) ions. The resulting solution was made up to $250 \mathrm{~cm}^{3}$ .

[ 2 ]

(i)

Calculate the amount, in moles, of dichromate(VI) ions used in the titration.

amount =mol

(ii)

Calculate the amount, in moles, of iron in the 3.35 g piece of wire.

amount =mol [1]

[ 1 ]

(iii)

Calculate the mass of iron in the 3.35 g piece of wire.

mass =g

(iv)

Calculate the percentage of iron in the iron wire.

percentage =% [1]

[ 1 ]

(a)

$10 \mathrm{~cm}^{3}$ of a gaseous hydrocarbon, $\mathrm{C}_{x} \mathrm{H}_{y}$ , was reacted with $100 \mathrm{~cm}^{3}$ of oxygen gas, an excess. The final volume of the gaseous mixture was $95 \mathrm{~cm}^{3}$ .

This gaseous mixture was treated with concentrated, aqueous sodium hydroxide to absorb the carbon dioxide present. This reduced the gas volume to $75 \mathrm{~cm}^{3}$ .

All gas volumes were measured at 298 K and 100 kPa .

[ 6 ]

(i)

Calculate the volume of carbon dioxide produced by the combustion of the hydrocarbon.

\text { volume of } \mathrm{CO}_{2} \text { produced }=\ldots \ldots \ldots \ldots \ldots \ldots . \mathrm{cm}^{3}

[ 1 ]

(ii)

Calculate the volume of oxygen used up in the reaction with the hydrocarbon.

\text { volume of } \mathrm{O}_{2} \text { used }=\ldots \ldots \ldots \ldots \ldots \ldots . . \mathrm{cm}^{3}

[ 1 ]

(iii)

Use your answers to (b)(ii) and (b)(iii), together with the initial volume of hydrocarbon, to balance the equation below.

\ldots \ldots \ldots \ldots . \mathrm{C}_{x} \mathrm{H}_{y}+\ldots \ldots \ldots \ldots . \mathrm{O}_{2} \rightarrow \ldots \ldots \ldots \ldots . \mathrm{CO}_{2}+z \mathrm{H}_{2} \mathrm{O}

[ 1 ]

(iv)

Deduce the values of x, y and z in the equation in (iv).
x=
y=
z=

[ 3 ]

[Maximum number: 12]

Carbon dioxide, $\mathrm{CO}_{2}$ , makes up about 0.040 % of the Earth's atmosphere. It is produced by animal respiration and by the combustion of fossil fuels.

In animal respiration, oxygen reacts with a carbohydrate such as glucose to give water, carbon dioxide and energy.

The typical daily food requirement of a human can be considered to be the equivalent of 1.20 kg of glucose, $\mathrm{C}_{6} \mathrm{H}_{12} \mathrm{O}_{6}$ .

You should express all of your numerical answers in this question to three significant figures.

(a)

(i)

Use your equation to calculate the amount, in moles, of CO2 produced by one person in one day from 1.20 kg of glucose.

(ii)

On the day on which this question was written, the World population was estimated to be $6.82 \times 10^{9}$ .

Calculate the total mass of $\mathrm{CO}_{2}$ produced by this number of people in one day. Give your answer in tonnes. [1 tonne $=1.00 \times 10^{6} \mathrm{~g}$ ]

[ 5 ]

(b)

When fossil fuels are burned in order to give energy, carbon dioxide and water are also produced.

The hydrocarbon octane, $\mathrm{C}_{8} \mathrm{H}_{18}$ , can be used to represent the fuel burned in motor cars. A typical fuel-efficient motor car uses about $4.00 \mathrm{dm}^{3}$ of fuel to travel 100 km .

[ 5 ]

(i)

The density of octane is $0.700 \mathrm{~g} \mathrm{~cm}^{-3}$ .

Calculate the amount, in moles, of octane present in $4.00 \mathrm{dm}^{3}$ of octane.

(ii)

Calculate the mass of $\mathrm{CO}_{2}$ produced when the fuel-efficient car is driven for a distance of 100 km .

[ 5 ]

(c)

Calculate how many kilometres the same fuel-efficient car would have to travel in order to produce as much $\mathrm{CO}_{2}$ as is produced by the respiration of $6.82 \times 10^{9}$ people during one day. Use your answer to (a)(iii).

[ 2 ]

[Maximum number: 4]

Zinc is an essential trace element which is necessary for the healthy growth of animals and plants. Zinc deficiency in humans can be easily treated by using zinc salts as dietary supplements.

(a)

A simple experiment to determine the value of x in the formula $\mathrm{ZnSO}_{4} \cdot x \mathrm{H}_{2} \mathrm{O}$ is to heat it carefully to drive off the water.

\mathrm{ZnSO}_{4} \cdot x \mathrm{H}_{2} \mathrm{O}(\mathrm{~s}) \rightarrow \mathrm{ZnSO}_{4}(\mathrm{~s})+x \mathrm{H}_{2} \mathrm{O}(\mathrm{~g})

A student placed a sample of the hydrated zinc sulfate in a weighed boiling tube and reweighed it. He then heated the tube for a short time, cooled it and reweighed it when cool. This process was repeated four times. The final results are shown below.

[ 7 ]

(i)

Why was the boiling tube heated, cooled and reweighed four times?

(ii)

Calculate the amount, in moles, of the anhydrous salt produced.

(iii)

Calculate the amount, in moles, of water driven off by heating.

(b)

For many people, an intake of approximately 15 mg per day of zinc will be sufficient to prevent deficiencies.

Zinc ethanoate crystals, $\left(\mathrm{CH}_{3} \mathrm{CO}_{2}\right)_{2} \mathrm{Zn} \cdot 2 \mathrm{H}_{2} \mathrm{O}$ , may be used in this way.

[ 4 ]

(i)

What mass of pure crystalline zinc ethanoate ( $M_{\mathrm{r}}=219.4$ ) will need to be taken to obtain a dose of 15 mg of zinc?

(ii)

If this dose is taken in solution as $5 \mathrm{~cm}^{3}$ of aqueous zinc ethanoate, what would be the concentration of the solution used?
Give your answer in $\mathrm{mol} \mathrm{dm}^{-3}$ .

[ 4 ]

[Maximum number: 6]

Compound A is an organic compound which contains carbon, hydrogen and oxygen.
When 0.240 g of the vapour of A is slowly passed over a large quantity of heated copper(II) oxide, CuO , the organic compound A is completely oxidised to carbon dioxide and water. Copper is the only other product of the reaction.

The products are collected and it is found that 0.352 g of $\mathrm{CO}_{2}$ and 0.144 g of $\mathrm{H}_{2} \mathrm{O}$ are formed.

(a)

In this section, give your answers to three decimal places.

[ 6 ]

(i)

Calculate the mass of carbon present in $0.352 \mathrm{~g}^{\circ}$ of $\mathrm{CO}_{2}$ .

Use this value to calculate the amount, in moles, of carbon atoms present in 0.240 g of A.

(ii)

Calculate the mass of hydrogen present in $0.144 \mathrm{~g} \mathrm{~g}^{\circ} \mathrm{H}_{2} \mathrm{O}$ .

Use this value to calculate the amount, in moles, of hydrogen atoms present in 0.240 g of A.

(iii)

Use your answers to calculate the mass of oxygen present in 0.240 g of A.

Use this value to calculate the amount, in moles, of oxygen atoms present in 0.240 g of A.

[ 6 ]