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FIJI SEVENTH FORM CERTIFICATE EXAMINATION 2013

CHEMISTRY PAPER 1

Time Allowed: Three Hours

(An extra ten minutes is allowed for reading this paper.)

INSTRUCTIONS

1. Write all your answers in the Answer Book provided.

2. Write your Index Number on the front page and inside the back flap of the Answer Book.

3. If you need more paper, ask the supervisor for extra sheets. Tie these inside the Answer Book at the appropriate places.

4. You may use a calculator, provided it is silent, battery-operated and non-programmable. In Sections B and C, working should be shown in the questions involving calculations.

5. There are three sections in this paper. All the sections are compulsory.

Note : A Periodic Table of Elements is provided for your use during this examination. Both the atomic number and the relative atomic mass for each element are given.

SUMMARY OF QUESTIONS

Section Guidelines Total

Mark

Suggested Time

A

There are twenty multiple-choice questions. All the questions are compulsory.

20 36 minutes

B There are ten questions.

All the questions are compulsory.

30 54 minutes

C There are five questions.

All the questions are compulsory.

50 90 minutes

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SECTION A

[20 marks] The multiple-choice questions in this section are all compulsory.

Each question is worth 1 mark.

INSTRUCTIONS FOR MULTIPLE – CHOICE QUESTIONS

1. In your Answer Book, circle the letter which represents the best answer. If you change your mind, put a line through your first choice and circle the letter of your next choice.

For example:

2. If you change your mind again and like your first answer better, put a line through your second circle and tick ( ) your first answer.

For example:

3. No mark will be given if you circle more than one letter for a question.

1. The electron configuration of an iron atom is best represented by

A. 1s2 2s22p6 3s23p6 4s0 3d8 B. 1s2 2s22p6 3s23p6 4s1 3d7 C. 1s2 2s22p6 3s23p6 4s2 3d6 D. 1s2 2s22p6 3s23p6 4s2 3d5 4p1

2. The set of quantum numbers describing the electrons in the outermost shell of a sodium atom is

A. n = 3, l = 0, ml = 0, ms = ±

2 1

B. n = 3, l = 1, ml = +1, ms = ±

2 1

C. n = 3, l = 2, ml = +2, ms = ±

2 1

D. n = 3, l = 3, ml = +3, ms = ±

2 1

12 A B C D

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3. An element X in Group II of the Periodic Table combines with another element Y from Group VI. Which of the statements is least likely to be true?

A. The compound will be ionic. B. X+2 and Y–2 ions will be formed.

C. The compound will have the formula XY. D. The compound will readily dissolve in water.

4. Given the following electronegativities: H = 2.2 N = 3.0 O = 3.5 F = 4.0

Which bond would be the most polar?

A. N − F in NF2

B. O − H in H2O

C. N − O in NO2

D. N − H in NH3

5. A container at standard atmospheric pressure contains 60 cm3 of a gas. What pressure is required to reduce its volume to 20 cm3?

A. ⅓_atmosphere

B. ½ atmosphere C. 2 atmospheres D. 3 atmospheres

6. Gas pressure is caused by the

A. weight of the gas molecules in the container. B. repulsion between gas molecules in the container. C. kinetic energy of the gas molecules in the container.

D. collision of the gas molecules with the walls of the container.

7. The substance that will form the strongest hydrogen bond in the liquid state is

A. NH3

B. H2O

C. H2S

D. CH4

8. Which of the following statements about transition metals is correct?

A. The solutions of transition metals are rarely coloured.

B. Complex ions of transition metals always have a coordination number 6.

C. Transition metals exhibit more than one oxidation state when they form compounds. D. There is a large difference between the first ionisation energies of iron and copper.

Turn Over

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9. The Group IV elements of the Periodic Table namely carbon, silicon, germanium, tin and lead

A. have stable oxidation states of +3 and +4. B. show the same properties as transition metals.

C. show increasing stability of the lower oxidation state.

D. become increasingly electronegative as the atomic number increases.

10. The standard heat of formation of carbonic acid is the heat change associated with the reaction

A. H2(g) + C(s) +

2 3

O2(g) H2CO3(l)

B. 2H(g) + CO2(g) + O(g) H2CO3(l)

C. 2H(g) + C(s) + 3O(g) H2CO3(l)

D. H2O(l) + CO2(g) H2CO3(l)

11. The thermochemical equation for the combustion of carbon is given as:

C(s) + O2(g) CO2(g); ∆Hºc = −394 kJmol−1

The mass of carbon that needs to be completely burned to produce 1 000 kJ of heat energy is

A. 2.5 g B. 12.0 g C. 30.5 g D. 111.7 g

12. The compound that will show geometrical isomerism is

A. CH2CBr2

B. CBr2CBr2

C. CH2CHBr

D. CHBrCHBr

13. The reaction between 2-chloropropane and hydroxide ion is an example of

A. an addition reaction. B. a condensation reaction.

C. a nucleophilic substitution reaction. D. an electrophilic substitution reaction.

14. The solution of the compound which is most basic is

A. NH3

B. (CH3)3N

C. (CH3)2NH

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15. Which of the reactions stated below will not result in the formation of ethanol?

A. oxidation of ethanal B. fermentation of glucose

C. the reaction of ethene with steam

D. the reaction of ethyl ethanoate with water

16. Which functional groups are present in alanine in the structure given below?

A. an amine and a methyl B. an amine and a carbonyl C. an amine and a carboxyl D. an amine and a hydroxyl

17. Which one of the following is the weakest acid?

Acid Ka pKa

A. sulphurous acid 1.2 x 10−2 1.92 B. phosphoric acid 7.1 x 10−3 2.15 C. citric acid 7.1 x 10−4 3.15 D. carbonic acid 4.5 x 10−7 6.35

18. When 1.0 cm3 of a weak acid solution is added to 100 cm3 of a buffer solution,

A. there will be almost no change in the pH of the solution. B. the volume of the resulting mixture will be 100 cm3. C. the pH of the solution will decrease noticeably. D. the pH of the solution will increase noticeably.

19. The oxidation number of sulphur in S2O32− is

A. +4

B. +2

C. −2 D. −4

20. Which one of the statements given below is not true for an electrochemical cell?

A. It cannot be recharged.

B. It converts chemical energy into electrical energy.

C. It is made up of two half cells connected by a salt bridge.

D. Electrons flow from the negative electrode to the positive electrode.

Turn Over

© MINISTRY OF EDUCATION, FSFCE 2013: CHEMISTRY. C C OH

CH3

H

N

O

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SECTION B

[30 marks] The ten questions in this section are all compulsory. Each question is worth 3 marks.

QUESTION 1

(a) Arrange the atoms of elements Mg, N, Cs and F in order of increasing

electronegativity. (1 mark)

(b) The table below shows the atomic radius and ionic radius for the element chlorine, Cl.

Atomic radius of Cl atom 99 picometres Ionic radius of Cl− 181 picometres

Explain why the chloride ion, Cl− has a larger radius than the chlorine atom, Cl. (1 mark)

(c) “Polar molecules such as HCl have a permanent dipole moment.”

Using HCl as an example, explain the statement above. (1 mark)

QUESTION 2

An athelete prepared a glucose drink before training by dissolving 25 g of glucose in 250 g of water.

(a) Determine the mole fraction of glucose (C6H12O6) in the drink. (2 marks)

(b) Calculate the molality of the glucose drink. (1 mark)

QUESTION 3

Transition metals are capable of forming complex ions like [Cu(NH3)4]2+.

(a) Define the term complex ion. (1 mark)

(b) Name the complex ion given above and state the coordination number of the

ion. (1 mark)

(c) Explain why ammonia, NH3 can effectively bond to Cu2+ in the formation of the

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QUESTION 4

Sodium oxide (Na2O) and phosphorus pentaoxide (P2O5) are examples of oxides of

elements in Period 3 of the Periodic Table.

(a) State with a reason the type of bonding present in Na2O. (1 mark)

(b) Write an equation to explain why the resulting solution of Na2O with water

is a basic solution. (1 mark)

(c) Complete the equation given below by writing the product for the reaction of P2O5 with water, in the box in the Answer Book.

P2O5(s) + H2O(l) (1 mark)

QUESTION 5

(a) Using the bond energy data given in the table, calculate the enthalpy change for the reaction:

C3H6(g) + Cl2(g) C3H5Cl(g) + HCl(g)

Bond energies (kJmol-1)

C – C 348

C – H 413

Cl – Cl 242

C – Cl 339

H – Cl 431

(2 marks)

(b) Is the reaction exothermic or endothermic? Explain your answer. (1 mark)

Turn Over

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SECTION B (continued)

QUESTION 6

(a) Write the structural formula of 2-methylpropan-l-ol in the box drawn in the

Answer Book. (1 mark)

(b) The structural formula of methyl propanoate is

Which of the two compounds (methyl propanoate or 2-methylpropan-l-ol) would you expect to have a greater solubility in water?

Explain your answer. (2 marks)

QUESTION 7

(a) In the box provided in the Answer Book, write the structural formula of the major organic product formed in the reaction shown below:

(1 mark)

(b) To which organic group of compounds does the major product belong? (1 mark)

(c) With a reason, classify the major product as either primary (1º), secondary (2º)

or tertiary (3º). (1 mark)

QUESTION 8

Carbon disulphide is formed according to the equation:

C(s) + 2S(s) CS2(l)

Using the data given below, calculate the heat of formation of carbon disulphide, CS2(l).

1. C(s) + O2(g) CO2(g); ∆Hºc = −393 kJmol−1

2. S(s) + O2(g) SO2(g); ∆Hºc = −297 kJmol−1

3. CS2(l) + 3O2(g) CO2(g) + 2SO2(g); ∆Hºc = −1007 kJmol−1

(3 marks)

H H

C C Cl

H

H H

+ NH3

alcohol

O

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QUESTION 9

A 1.00 L of a 0.100 molL-1 solution of sodium nitrate is added to 1.00 L of 0.100 molL-1 solution of calcium nitrate. Calculate the concentration of the nitrate ions in the final

solution. (3 marks)

QUESTION 10

Given the following standard electrode potentials:

Ag+/Ag + 0.80 V Cu2+/Cu + 0.34 V Zn2+/Zn − 0.76 V Al3+/Al − 1.66 V Mg2+/Mg − 2.32 V

(a) Identify the strongest oxidizing agent. (1 mark)

(b) Calculate the maximum emf which could be obtained by combining a pair of the

above couples. (2 marks)

Turn Over

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SECTION C

[50 marks]

The five questions in this section are all compulsory. Each question is worth 10 marks.

QUESTION 1

(a) Naturally-occurring boron is composed of 19.8% of 10B and 80.2% of 11B. The relative atomic mass of 10B is 10.0129 and that of 11B is 11.0093. Calculate

the average atomic mass of boron. (2 marks)

(b) Quantum numbers are used to describe the orbitals in which the electrons can be found.

(i) Name the quantum number that describes the main energy levels. (1 mark)

(ii) Draw and describe the shape (electron cloud) of a ‘p’ orbital. (1 mark)

(iii) What would be the maximum number of electrons that could be held by

a ‘d’ orbital? (1 mark)

(c) Draw the Lewis structure for carbonate ion, CO₃2− and predict its shape. (2 marks)

(d) Transition metals such as iron, nickel and cobalt are paramagnetic. Draw the

orbital diagram for nickel and explain why it is paramagnetic. (2 marks)

(e) Silicon tetrachloride, SiCl4 is a covalent halide which hydrolyses readily in water

as shown by the following equation:

SiCl4 + 2H2O SiO2 + 4HCl

Diagrammatic representation:

Using the diagrammatic representation above, explain how hydrolysis is brought

about. (1 mark)

H

O H Cl

Si Cl Cl

Cl

H O H

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QUESTION 2

(a) Some oxygen gas was collected over water at 20ºC and 101.3 kPa and was found to occupy 30 mL.

Calculate:

(i) the partial pressure of the oxygen gas. [SVP of water at 20ºC = 2 kPa] (1 mark)

(ii) the volume that the gas would occupy at STP. (2 marks)

(iii) the mass of oxygen gas collected. [R = 8.314 JK−1mol−1] (1 mark)

(b) Graphs I and II show the relative concentration of species in solutions of ethanoic (acetic) acid and magnesium chloride.

(i) Identify the species B and C present in the solution, in Graph I. (1 mark)

(ii) Identify the graph which shows the species present in ethanoic acid. (1 mark)

(iii) Which solution will be a better electrolyte? Give a reason for your

answer. (1 mark)

(c) Balance the following redox equation in an acidic solution:

) g ( 2 2 ) aq ( 2 ) aq ( 4 2 ) aq (

4 C O Mn CO

MnO− + − + + (3 marks)

QUESTION 3

(a) Write the systematic name of the following compound of benzene.

(1 mark)

Turn Over

OH

Relative concentra-

tion

A B C D E

Graph II Graph I

Relative concentra-

tion

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SECTION C (continued)

(b) Analysis of an organic compound indicated an empirical formula of C3H6O2

and a relative molecular mass of 74.

Write the structural formula of this organic compound and name one of its

isomers. (2 marks)

(c) The following graph is the titration curve for an acid-base reaction.

(i) Which type of acid-base titration is represented by the above curve? (1 mark)

(ii) What is the approximate acid dissociation constant, pKa, for the acid? (1 mark)

(iii) What would be the pH at the end point of the titration? (1 mark)

(iv) Explain which indicator from the following list would be the most suitable for the above titration.

Indicator pKa

methyl orange 3.7

methyl red 5.1

bromothymol blue 7.0 phenolphthalein 9.6

(1 mark)

0 2 4 6 8 10 12 14 16 18 20

0 2 4 6 8 10 12 14

♦

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(d) When water reacts with propene, the products formed are propan-2-ol and propan-1-ol.

(i) Using the structural formula, write an equation to show the formation

of propan-2-ol. (1 mark)

(ii) Use your knowledge of Markovnikov’s Rule to explain the formation of

propan-2-ol. (1 mark)

(iii) For the reaction shown below, explain why Markovnikov’s Rule does not apply.

(1 mark)

QUESTION 4

(a) In an experiment, a 75 mL of 0.002 molL−1 solution of potassium chloride, KCl, was mixed with 25 mL of 0.005 molL−1 solution of silver nitrate, AgNO3.

[Ksp (AgCl) = 1.7 x 10−10 mol2 L−2]

(i) Write the ionic equation to show the formation of the product, silver

chloride. (1 mark)

(ii) Calculate the concentration of the silver ions [Ag+] and chloride ions

[Cl−]. (2 marks)

(iii) Determine whether a precipitate of silver chloride will form or not.

(Show full working.) (2 marks)

(b) The standard enthalpy (heat) of combustion of carbon is given as follows:

C(s) + O2(g) CO2(g); ∆Hºc = − 394 kJmol−1

and the standard enthalpy (heat) of formation of carbon dioxide is given as follows:

C(s) + O2(g) CO2(g); ∆Hºf = − 394 kJmol−1

Explain why the two equations above have the same ∆Hº values. (1 mark)

Turn Over

C C H

H H

H OH H C C H

H OH H H

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SECTION C (continued)

(c) Given: Eº (Fe3+ / Fe2+) = + 0.77 V Eº (Br2 / Br−) = + 1.09 V

Determine whether the following reaction will be spontaneous under standard conditions:

Br2(aq) + 2 Fe2_(aq)+ 2 Fe3_(aq)+ + 2 Br__(aq) (2 marks)

(d) Esters undergo alkaline hydrolysis in a reaction known as saponification.

(i) Complete the equation given below by writing the products in the boxes in the Answer Book.

(1 mark)

(ii) Explain how the saponification process takes place. (1 mark)

QUESTION 5

(a) Consider the following reaction sequence.

(i) Name the compounds C and D. (2 marks)

(ii) Identify the reagents X and Y. (2 marks)

(iii) Write an equation to show the conversion of Compound B back to

Compound A in the laboratory. State the conditions in the equation. (2 marks) CH3CH2 C

O

OCH3

+ NaOH +

Compound D

Compound A CH2 = CH2

Compound C Br2

H2

Ni

H2O

H3PO4/330ºC

Reagent Y

Reagent X CH3COCl

CH2Br – CH2Br

CH3COOH

CO2 + H2O

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(b) Zinc reacts with sulphur according to the equation:

Zn(s) + S(s) ZnS(s)

30.0 g of zinc was mixed and heated with 36.0 g of sulphur until the reaction was

complete. Which reactant is the limiting reagent? (2 marks)

(c) Ethanol and dimethyl ether have the same molecular formula, C2H6O. The boiling

point of ethanol is 78.4ºC while that of dimethly ether is –23.7ºC.

Account for the difference in the boiling point of the two liquids. (2 marks)

THE END

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