Tutorial LU1-7 Sem 1

PRK 1016 CHEMISTRY I UNIMAS 2013 NOT FOR SALE Page 1 of 38

Tutorial 1

LU 1: Chemistry and Chemical Nomenclature

Physical and chemical changes

1. Decide whether each of the following processes in physical change or a chemical

change, and explain briefly.

(a) Droplets of water forms on the outside of a glass of ice tea.

(b) A match ignites to form ash and a mixture of gases.

(c) Perspiration evaporates when you relax after jogging.

(d) An iron nail forms rust slowly in air.

Element, compound and mixture

2. Classify each of the following as an element, a compound, or a mixture.

(a) Gold (b) Air (c) Water

(d) Carbon dioxide (e) Seawater (f) Oxygen

Chemical formula

3. Write the formulae for the following compounds:

Expressions of numbers

4. Express the following numbers in scientific notation:

(a) 0.000000027

(b) 47 764

(c) 0.00096

5. Express the following numbers as decimals:

(a) 1.52 102

(b) 7.78 108

(c) 2.357 105

6. Express the answers to the following calculations in scientific notation:

a) 145.75 + (2.3 101)

b) 79 500 (2.5 102)

Compound Formula Compound Formula

Copper(II) chloride Lithium acetate

Iron(II) sulphate Manganese(IV) nitrate

Sodium sulphate Potassium permanganate

Ammonium nitrate Iron(III) oxide

Sodium nitrite Potassium sulphite

Potassium sulphate Potassium hydrogen carbonate

Magnesium phosphate Sodium iodate


7. A child has a body temperature of 38.7oC.

(a) If normal body temperature is 98.6oF, does the child have a fever?

(b) What is the childs body temperature in Kelvin?

Significant numbers

8. State the number of significant figure in each of the following numbers.

(a) 2.0560

(b) 0.000372

(c) 300.6

(d) 70030

9. What is the number of significant figures in each of the following measurements?

(a) 2900 g

(b) 40.2 g mL1

(c) 0.00000030 cm

(d) 4.6 1019 atoms

10. Perform the following calculations and round the answers to the correct number of

significant figures.

(a) 3.476 + 0.002 (b) 81.4 g 0.112 g

(c) 81.4 104.2 (d) 4.003 18.3

(e) 0.321 + 0.0041 1.2 (f) 6.810 0.0230 3.21

Percentages

11. A student experimentally determines the specific heat of water to be 4.29 J/g /oC. He

then looks up the specific heat of water on a reference table and finds that is

4.18 J/g/oC. What is his absolute error and percent error?

12. The human body is 0.0040% iron. How many milligrams of iron does a 64 kg person

contain?

Dimensional analysis

13. The density of ammonia gas under certain conditions is 0.625 g/L. Calculate the

density in g/cm3.

14. The speed of light is 3.0 108 m s1. What would be the speed of light in km hr1?

15. A company wants 800 square feet of carpet, but the carpet store sells only by the

square meter. How many square meters does the company need to buy?

[1 m = 39.37 inches]


Sub-atomic particles and relative atomic mass

16. Complete the table for some of the elements occurring naturally.

Element

Symbol

of

Element

Proton

number

Nucleon

number

Number

of

neutrons

Number

of

electrons

Symbolic

representation

Beryllium Be 4 5

Oxygen O 16 8

Potassium K 19 39 20 19

Copper Cu 29 64 27

17. Strontium consists of four isotopes with masses of 84 (abundance 0.50%), 86

(abundance of 9.9%), 87 (abundance of 7.0%), and 88 (abundance of 82.6%).

Calculate the average atomic mass of strontium.

18. Naturally occurring hydrogen on earth has an atomic mass of 1.0079 amu. Suppose

you were on another planet and found the atomic mass of hydrogen to be

1.2000 amu.

(a) How would you explain this?

(b) What is the percentage abundance of 1H on the other planet?(the isotopic mass of 1H is 1.0079 amu and 2H is 2.1041 amu)

Mass spectrometer and mass spectrum

19. List and describe briefly the functions of the five major components of a mass

spectrometer.

20. Figure below shows the mass spectrum of water.

(a) Identify the ions which cause the peaks at 1, 16, 17 and 18.

(b) Which of the species is the most stable? Why?

(c) Which of the species are called molecular fragments?

(d) Which species is called molecular ion?


Consolidated problem

21. The relative atomic mass of A is 79.99.

(a) Why the relative atomic mass of A not a whole number?

(b) A contains two isotopes with mass numbers 79 and 81. Calculate the

percentage abundance of the isotope 79A and 81A.

(c) If A exists as a monoatomic element, state the ion and its m/e detected in the

mass spectrometer.

(d) Sketch the mass spectrum for A in (c).

(e) If A exists as a diatomic element, state the molecular ion and its m/e detected

in the mass spectrometer.

(f) Sketch the mass spectrum for A in (e).

22. Figure below shows the mass spectrum of water.

(a) Name the equipment used to analyse water.

(b) What is meant by a base peak?

(c) In which part of the equipment will ions be formed?

(d) Write a chemical equation to show formation of molecular ion of water.

(e) In which part of the equipment where fragmentation will happen?

(f) From the figure above, indicate clearly what chemicals detected on m/e with 17

and 16.

(g) If hydrogen has isotopes of 1H and 2H, and oxygen has 16O and 18O, what are

the molecular ions of water can be detected. Make a table to show the

molecular ion and m/e.

Group Discussion & Presentation

23. Your group has been assigned to run a brief presentation on the determination of

the molecular mass of one of the following elements/compounds by using a mass

spectrometer:

Group 1 bromine, Br2

Group 2 propane, CH3CH2CH3

Group 3 methanol, CH3OH

Group 4 carbon dioxide, CO2

Your presentation may include:


- What happened to the sample in the mass spectrometer, starting from the

vapourisation chamber

- The expected mass spectrum (with important peaks only)

- How to recognise the base peak

- Interpretation of the mass spectrum obtained


Tutorial 2

LU2: Stoichiometry

Mole concept and Avogadro constant

1. Calculate the number of particles / atoms in:

[assume NA = 6.02 1023 ]

(a) 0.35 mole of Al

(b) 2.7 moles of Cl

(c) 0.048 mole of CO2

(d) 14.8 g of copper

(e) 7.5 g of ammonia, NH3

2. Calculate the number of moles for the following substances:

[assume NA = 6.02 1023 ]

(a) 6.02 1020 iron atoms

(b) 8.15 1023 carbon monoxide molecules

(c) 6.02 1022 sodium ions

(d) 2.34 g of bromine gas, Br2

(e) 17.5 g of copper(II) nitrate, Cu(NO3)2

3. Determine the mass for each of the following substances:

[assume NA = 6.02 1023 ]

(a) 0.091 mole of ascorbic acid, C6H8O6

(b) 0.125 mole of magnesium hydroxide, Mg(OH)2

(c) 1.5 1022 zinc atoms

(d) 2 1023 molecules of ethanol, C2H5OH

(e) 5.72 1020 calcium ions

4. When aqueous sodium hydroxide solution reacts with dilute sulphuric acid solution,

sodium sulphate and water are produced.

(a) Write a balanced chemical equation for the reaction

(b) State the ratio for number of moles between sodium hydroxide, sulphuric acid,

sodium sulphate and water based on the equation in (a)

(c) Calculate the number of moles of sodium hydroxide needed to produce 0.5 mole of

sodium sulphate.

Molar volume of gases

5. Calculate the volume occupied by the following gases at room temperature .

[Assume 1 mole of gas occupies a volume of 24.0 L at room temperature]

(a) 0.75 mole of nitrogen gas

(b) 3.0 mole of chloromethane gas, CH3Cl(g)

(c) 2.4 g of ammonia gas

(d) 0.056 g of carbon monoxide, CO

(e) 1.7 1023 molecules of ethane gas

(f) 9 1021 molecules of chlorine gas


6. Calculate the number of moles for each of the following gases at STP.

[Assume 1 mole of gas occupies a volume of 22.4 L STP]

(a) 1802 mL of methane gas, CH4

(b) 0.71 m3 of carbon dioxide

7. Calculate the mass for each of the following gases at STP.

[Relative atomic mass: H, 1.0; C, 12.0; O, 16.0; S, 32.0]


(a) 9.3 L of sulphur dioxide gas

(b) 700 mL of methane gas, CH4

8. Calculate the number of molecules present at STP for each of the following

substances:[Relative atomic mass: C, 12.0; N, 14.0; O, 16.0; S, 32.0]


(a) 0.37 L of nitrogen gas

(b) 480 mL of carbon monoxide

9. In an experiment, a student notices that carbon dioxide, CO2, occupy the same

volume as 0.036 kg of butane gas, C4H10, under the same conditions. Calculate the

mass of carbon dioxide.

10. Hydrogen gas, H2, burns in oxygen gas, O2, to form water.

(a) Write a balanced chemical equation for this reaction.

(b) What is the molar ratio for the reactants involved?

(c) If 1 mole of hydrogen gas is used, what will be the volume of hydrogen gas,

oxygen gas and water you expect at 25C?

(d) What is the assumption you made in the calculation of volume in 10(c)?

(e) What is the mass of the oxygen gas needed to react with 20 L of hydrogen gas at

room temperature and what is the volume of water produced?

11. Ethanol, C2H5OH(l), burns in oxygen to form carbon dioxide and water.

(a) Write a balanced chemical equation for this reaction.

(b) State the molar ratio for each substance involved in the reaction.

(c) Calculate the mass of ethanol burnt if 2.40 L of carbon dioxide is produced at

room temperature.

Molarity

12. What is the molarity of a solution prepared by dissolving 0.10 mole of NaOH in

0.50 L of solution?

13. Calculate the molarity of a solution prepared by dissolving

[Relative atomic mass: Na, 23.0; O, 16.0; H, 1.0; Ca, 40.0; Cl, 35.5; K, 39.0]

(a) 4.00 g of NaOH in 100.0 mL of solution

(b) 16.0 g of CaCl2 in 250.0 mL of solution

(c) 14.0 g of KOH in 75.0 mL of solution


14. How many moles of solute are there in each of the following solutions:

(a) 1.0 L of 0.10 molar NaOH

(b) 250 mL of 1.0 molar NaOH

(c) 20.00 mL of 0.20 molar H2SO4

15. How many grams of solute are there in each of the following solutions:

[Relative atomic mass: Na, 23.0; O, 16.0; H, 1.0; C, 12.0]

(a) 250 mL of 1.0 M NaOH

(b) 20.00 mL of 0.10 M CH3COOH

Dilution

16. 25.00 mL of 0.10 M NaOH was diluted to become 250 mL solution. Determine the

concentration of the diluted solution.

17. 0.22 M hydrochloric acid is diluted to become 250.0 mL of 0.10 M. How much of the

original acid must be used for dilution?

18. To what volume must 25.0 mL of 18.0 M H2SO4 be diluted to produce 1.50 M H2SO4?

19. Given a solution of 2.50 M KOH, how do you prepare 1.00 M KOH solution from it?

(state all the steps involved in the dilution process)

20. (Challenging) Concentrated hydrochloric acid has a purity of 36.0 % by mass. The

density of the concentrated acid is 1.19 g/mL. Determine the molar concentration of

this acid.

Molality

21. The molality of a sodium chloride solution is 0.150 m. How many moles of NaCl must

be dissolved in 500 g of water to prepare a solution of this molality?

22. If you dissolve 4.00 g NaOH in 250 g water, what is the molality of the solution

produced?

23. Calculate the mass of methanol, CH3OH needed to prepare a 0.250 m solution using

2000 g of water.

24. (Challenging) Commercial hydrobromic acid can be purchased as 40.0 % by mass of

HBr. The density of this acid is 1.38 g mL1 . Calculate the molality of this

commercial acid.

25. (Challenging) A solution of ethanol(C2H5OH) has a concentration of 4.613 mol L1 at

20C. The density of the solution is 0.9667 g mL1. Calculate the molality of the

solution.


Other units of concentration

26. How much ethanol(C2H5OH) must be used to prepare 150 mL 10% by volume?

27. A solution of acetic acid was prepared by mixing glacial acetic acid with water. The

solution was 25% of acetic acid by volume. How much glacial acetic acid must be

used to prepare 120 mL of the desired solution?

28. A limestone sample weighing 1.267 g contains 0.3684 g iron. Determine the % of iron

by mass in the sample.

29. (Challenging) A solution of ethanol(C2H5OH) has a concentration of 4.613 mol L1 at

20C. The density of the solution is 0.9667 g mL1. Calculate the percent of ethanol

by mass.

30. A sample contains 2.91% by mass of iron. Express the result in

(a) ppm

(b) ppb

31. A 0.456 g sample of an ore is found to contain 0.280 mg chromium.

(a) Express the result in % by mass

(b) Express the result in ppm

(c) Express the result in ppb

32. A vessel contains a mixture of 2 moles of oxygen and 4 moles of helium. What is the

mole fraction of oxygen?

33. Consider a solution containing 0.50 mole of KNO3 and 0.30 mole of NaCl in 100 g of

water. Calculate the mole fraction of each compound in the solution.

34. (Challenging) A solution of ethanol (C2H5OH) has a concentration of 4.613 mol L1 at

20C. The density is 0.9677 g mL1. Calculate the mole fraction of ethanol in the

solution.

Empirical formulas & molecular formulas

35. An organic acid, with molecular mass 88, contains 54.55% C, 9.09% H and 36.36% O.

Determine its:

(i) empirical formula, and

(ii) molecular formula.

36. 6.92 g of X reacts with 0.584 g of carbon to form a compound. Given that the relative

atomic mass of C is 12.0 and the molecular formula of the compound is CX4.

Calculate the relative atomic mass of X.


37. Mannose is a sugar containing only C, H and O, and its relative molecular mass is

180. If 2.36 g of a sample of mannose contains 0.944 g of carbon and 0.158 g of

hydrogen, what is its molecular formula?

Writing & balancing chemical equations

38. Balance the following chemical equations:

(a) CaCO3(s) + HCl(aq) CaCl2(aq) + CO2(g) + H2O(l)

(b) KClO3(s) KCl(s) + O2(g)

(c) Fe2(SO4)3(aq) + NaOH(aq) Fe(OH)3(s) + Na2SO4(aq)

(d) C2H5CH2OH(l) + O2(g) CO2(g) + H2O(l)

39. Write the ionic equations for the following reactions:

(a) CuCl2(aq) + Pb(NO3)2(aq) Cu(NO3)2(aq) + PbCl2(s)

(b) ZnSO4(aq) + BaCl2(aq) ZnCl2(aq) + BaSO4(s)

(c) CaCl2(aq) + Na2SO4(aq) CaSO4(s) + NaCl(aq)

Interpreting a chemical equation

40. Based on the following equation:

ClO2 + H2O HClO3 + HCl

(a) Balance the above equation.

(b) How many grams of HClO3 can be obtained from 24.4 g of ClO2?

(c) How many grams of H2O are needed to produce 11.44 g of HCl?

(d) Find the mass of HClO3 produced when 6.25 g of ClO2 reacts with an excess

amount of H2O.

41. (a) Write a balanced equation for the reaction between aluminium and sulphuric

acid.

(b) If 20.0 g of aluminium is placed in a solution containing an excess amount of

sulphuric acid,

(i) how many mole of H2 will be formed?

(ii) what volume of H2 will be formed at STP?

(iii) how many gram aluminium sulphate will be produced?

Limiting reactant and yield

42. (a) Write a balanced equation for the reaction between aluminium and sulphuric

acid.

(b) If 20.0 g of aluminium is placed in a solution containing 115.0 g of sulphuric

acid,

(i) which one is the limiting reactant?

(ii) how many mole of H2 will be formed?

(iii) how many grams of aluminium sulfate will be produced?

(iv) calculate the mass of the excess reactant remains after the reaction.


43. Potassium nitrate can be prepared by the reacton between KCl and HNO3. The

equation is shown below.

KCl + HNO3 KNO3 + Cl2 + NOCl + H2O

(a) Balance the above equation.

(b) 100 g of KCl is mixed with 100 g of HNO3. Determine which reactant is the

limiting reactant.

(c) What is the mass of KNO3 formed?

(d) Calculate the mass of the excess reactant left after the reaction is completed.

44. When 4.8 g of a sample of impure anhydrous sodium carbonate reacted with excess

of sulfuric acid, 1.96 g of carbon dioxide was produced. Calculate the percent purity

of the sample.

45. (a) Write a balanced chemical equation for the reaction between iron and

hydrochloric acid.

(b) If 0.40 mol of iron powder was mixed with 0.75 mol hydrochloric acid,

(i) which one is the limited reactant?

(ii) how many moles of the excess reactant will remain after the reaction is

completed?

(iii) what is the volume of hydrogen gas produced at STP?

46. A bleaching agent, sodium hypochlorite, NaClO can be prepared as shown by the

equation below:

2NaOH + Cl2 NaCl + NaClO + H2O

If 75.0 g NaOH was mixed with 50g Cl2,

(a) which one is the limited reactant?

(b) what is the theoretical yield of NaClO that can be obtained from the mixture?

(c) 43.2 g of NaClO was obtained from the mixture. What is the percentage yield of

NaClO in this preparation?

Atom economy

47. Hydrogen gas is made on a large scale by reacting natural gas (methane) with steam.

(a) Write the balanced equation for the reaction.

(b) Calculate the atom economy of this reaction.

48. Based on the following equation:

2SO2 + O2 2SO3

(a) calculate the maximum theoretical mass of sulfur trioxide that can be made by

reacting 96 g of sulfur dioxide with an excess of O2.

(b) calculate the atom economy of this reaction.


Oxidation numbers & balancing redox equations

49. (a) What is oxidation number?

(b) Determine the oxidation number of the underlined elements in the following

species.

(i) Na2CO3 (v) H2SO4

(ii) Ca(OH)2 (vi) NaHSO4

(iii) MnO2 (vii) IO3

(iv) MnO4 (viii) Cr2O72

50. Balance the following equations:

(a) Cr2O72 + H2S Cr3+ + S (in acidic solution)

(b) MnO4 + C2O42

MnO2 + CO32 (in alkaline solution)

(c) Zn + MnO4 Zn(OH)2 + MnO2 (in alkaline solution)

(d) Fe2+ + MnO4- Fe3+ + Mn2+ (in acidic solution)

Consolidated problem

51. A bleaching agent, sodium hypochlorite, NaClO can be prepared as shown by the

equation below:

2NaOH(aq) + Cl2(aq) NaCl(aq) + NaClO(aq) + H2O(l)

If 400 mL of 4 M NaOH was mixed with 55 g Cl2,

(a) which one is the limiting reactant? (show calculations)

(b) which one will be the excess reagent?

(c) what is the theoretical yield of NaClO that can be obtained from the mixture?

(d) 42.2 g of NaClO was obtained from the reaction mixture. What is the percentage

yield of NaClO in this preparation?

(e) What is the mass of excess reagent left after the chemical reaction?

(f) Determine the oxidation number of Cl in NaCl and NaClO


Tutorial 3

LU3: Thermochemistry

System and surrounding

1. Explain the following situations in terms of heat transfer, system and surrounding:

(a) Heating of water using a kitchen gas

(b) You cool down after swimming

(c) You used water to put out a fire

(d) Heat packs used by sportspeople to relax their muscles

Enthalpy change of reaction

2. Consider the following reaction:

2CH3OH(l) + 3O2(g) 4H2O(l) + 2CO2(g) H = 1452.8 kJ

What is the value of H if:

(a) The equation is multiplied by 2.

(b) The direction is reversed so that the products become the reactants and vice

versa.

Specific heat capacity and q

3. A 50.0 g of sample of a dilute acid solution is added to 50.0 g of a base solution in a

coffee cup calorimeter. Temperature of the solution increases from 18.20C to

21.30C. Calculate q for the neutralisation reaction. [The specific heat of water is

4.20 J/g.K]

4. The reaction of 0.440 g magnesium with 400 g of hydrochloric acid solution causes

the temperature of the solution to increase by 5.04C. Assume that the specific heat

and density of the solution is the same as the water and Mg is the limiting reactant.

Calculate the heat produced in the reaction.

5. When a 43.0 g of sample of metal at 100C is added to 38.0 g water at 23.72C. The

final temperature of both metal and the water is 29.33C. What is the specific heat of

the metal? [The specific heat of water is 4.20 J/g.K]

Enthalpy change of reaction

6. When 4.14 g of potassium carbonate was added to 50.0 mL of 2.0 M of hydrochloric

acid the temperature of the solution increases by 4.6C.

(a) Write an equation for the reaction.

(b) Determine q in the reaction mixture. [The specific heat capacity and density of

the solution is 4.20 J/g.K and 1.0 g mL1 respectively]

(c) Hence, calculate the enthalpy change for the reaction per mole of potassium

carbonate.


Enthalpy change of formation and enthalpy change of combustion

7. Which are the thermochemical equations for enthalpy change of formations and

explain why?

(a) N2(g) + 2

3H2(g) NH3(g)

(b) 2H2(g) + O2(g) 2H2O(l)

(c) Fe2O3(s) + 3SO3(g) Fe2(SO4)3(s)

(d) Fe(s) + N2(g) + 3O2(g) Fe(NO3)2(s)

8. Write the correct thermochemical equations for the formation of the following

substances:

(a) NCl3 (b) O3 (c) BaCO3 (d) NaNO3

9. Write the thermochemical equations for the combustion of the following substances:

(a) Hydrogen

(b) Propane, C3H8(g)

(c) Ethanol, CH3CH2OH(l)

Use of enthalpy change of formation and enthalpy change of combustion

10. One step in the production of nitric acid, a powerful acid used in the production of

fertilizers is the combustion of ammonia.

4NH3(g) + 5O2(g) 4NO(g) + 6H2O(g)

By using standard enthalpy of formation given, find the standard enthalpy change of

reaction. [Hf for NH3, NO and H2O are 46.11 kJ mol1, +90.25 kJ mol1 and

241.82 kJ mol1 respectively]

11. When 1.15 g of ethanol was burned under a container of water, it was found that 200

mL of water was heated from 28.0C to 60.4C. The process was found to be 80%

efficient. [The specific heat of water is 4.20 J/g.K].

(a) Write a balanced equation for the combustion of ethanol

(b) Calculate the enthalpy change of combustion per mole of ethanol

(c) The process was 80% efficient, means that the experimental value of Hc is less

than theoretical value. Suggest why?

12. Using the following standard enthalpy of formation, calculate the standard enthalpy

change of combustion of methanol.

Compound Hf (kJ mol1)

CH3OH 238.7

H2O 285.9

CO2 93.5


Hesss Law and Born-Haber Cycle

13. Two gaseous pollutants that form in auto exhaust are CO and NO. A chemist is

studying ways to convert them to less harmful gases through the following reaction:

CO(g) + NO(g) CO2(g) + N2(g) H = ?

Given the following equation, calculate the unknown H?

CO(g) + O2(g) CO2(g) H = 283.0 kJ

N2(g) + O2(g) 2NO(g) H = 180.6 kJ

14. Calculate the H for the reaction of H2O(l) H2(g) + O2 (g) by using data below:

H2O(l) H2O(g) H = 44 kJ

H2O(g) H2(g) + O2(g) H = 242 kJ

15. Calculate the lattice energy of lithium chloride from the following data:

H/ kJmol-1

First ionisation energy of lithium +522

Enthalpy change of atomisation of lithium +160

Enthalpy change of atomisation of chlorine +121

Electron affinity of Cl2 atoms 365

Enthalpy change of formation of lithium chloride 410

16. Construct a Born-Haber cycle energy diagram and use it to determine the 1st

electron affinity of chlorine. Given to you the following data:

H/ kJ mol-1 H atomisation of Cl +122

H atomisation of Mg +148

H 1st ionisation energy of Mg +738

H 2nd ionisation energy of Mg +1451

H lattice energy of MgCl2 2526

H formation of MgCl2 641

Bond energy and enthalpy change of reaction

17. Calculate the standard enthalpy of reaction:

C2H4(g) + Br2(g) CH2BrCH2Br(l)

[Mean standard bond enthalpies are in kJ mol1: CC, 348; CBr, 276; BrBr, 193;

HBr, 366; C=C, 614]


Entropy

18. How does the entropy of a system change for each of the following processes?

(a) A solid melts

(b) A liquid freezes

(c) A solid sublimes

(d) A vapor condenses to a liquid

(e) A gas dissolves in another gas

(f) A pure solute dissolves in water

(g) Formation of solid magnesium oxide

(h) Reaction between ethanoic acid and ammonium carbonate

Consolidated problems

19. When 0.020 mol potassium carbonate is added to 30.0 mL of 2.0 mol L1 HCl, the

temperature of the solution increased by 5.18C. The reaction occurred is

K2CO3(s) + 2HCl(aq) 2KCl(aq) + CO2(g) + H2O(l)

(a) Determine which one is the limiting reactant.

(b) Determine the heat produced in the reaction.

(c) Determine the enthalpy change for the chemical reaction.

(Specific heat of solution = 4.5 J g1 C1 ; density of solution = 1.25 g mL1)

20. When 50.0 mL of 1 M hydrochloric acid is added into a beaker containing 50.0 mL of

1 M potassium hydroxide, the temperature increases from 25.0C to 31.9C.

(a) Write the chemical equation for the reaction.

(b) Write the ionic equation for the reaction.

(c) Calculate the heat released from the neutralisation. [specific heat capacity for

solution is 4.20 J g1 C1 and density of solution is 1.00 g mL1 ]

(d) Calculate the number of mol HCl used in the reaction.

(e) Calculate the enthalpy change of neutralisation as in equation (a).


Tutorial 4

LU4: Structure of Atoms

Electronic Energy Levels

1. Using a Bohrs model of a hydrogen atom, draw a diagram to show what happens to

its electron when the atom is heated?

2. Using a Bohrs model of a hydrogen atom, draw a diagram to show what happens to

its excited electron when the atom is cooled?

3. If a line spectra is made up by electronic transitions from higher levels to n = 1,

arrange the transitions below in the order of increasing wavelength ( ).

(a) n2 n1 (b) n3 n1 (c) n4 n1 (d) n5 n1 (e) n7 n1

4. How many electronic transitions can possibly occur for a hydrogen atom as

indicated by the following energy levels (shells)? (try drawing the lines in the

diagram below)

n=5

n=4

n=3

n=2

n=1

5. Draw a diagram to indicate the electronic transitions of the first three line spectra

formed in the Lyman series.

6. The diagram below shows a line spectra in the visible region.

i) Indicate clearly the first three lines.

ii) Indicate also the electronic transitions that can give rise to each line.

7. When an electron is excited because of energy absorption, it moves up to a higher

energy shell. What will happen to the atom if the electron is moved to an infinite

number of shells away from the nucleus of the atom?

8. What is the wavelength of light with a frequency of 3.75 1014 s1?

9. What is the wavelength of light with the energy of 2.25 1018 J?

E

n

e

r

g

y


10. A line spectrum in the Lyman Series has a wavelength of 103 nm. Calculate the

energy given out by the electron that corresponds to the wavelenght. (use c = f and

E = hf) [Speed of light, c = 3.00 108 m s1, Plancks constant, h = 6.63 1034 Js]

11. Calculate the wavelength of the first line in a Lyman Series.

(Rydbergs constant, RH = 1.097 107 m1, c = 3.000 108 m s1)

12. Calculate the wavelength of the first line in a Balmer Series.

(Rydbergs constant, RH = 1.097 107 m1, c = 3.000 108 m s1)

13. For a hydrogen atom, if the electron is promoted from n = 1 to n = , ion H+ will be

formed.

(a) Use the Rydbergs equation to calculate the wavelength of the light required to

cause the movement of the electron.

(b) Calculate the energy required by an electron in this situation.

(c) Calculate the energy required by 1 mole of hydrogen atoms for similar

transition.

(d) Write an equation for the ionisation a hydrogen atom.

(e) What is the ionisation energy of 1 mole of hydrogen atoms?

14. Light of wavelength 162.7 nm is required to ionise magnesium atoms to magnesium

ions, Mg+.

(a) Calculate the energy required to change one magnesium atom to an ion, Mg+.

(b) If 1 mole of magnesium atoms are used, calculate the energy used.

[1 nm = 1 109 m, Speed of light, c = 3.00 108 m s1, Plancks constant,

h = 6.63 1034 Js, Avogadros number = 6.02 1023 mol1]

15. With reference to line spectrum in a Lyman Series and a Balmers Series, show

their differences by using a table.

Atomic Orbitals

16. Write the possible values for l and m for each of the principal quantum numbers

below:

(a) n = 1

(b) n = 2

(c) n = 4

(d) n = 5

17. For each of the principal quantum numbers listed in question 16 above, state the

number of allowed orbitals.


18. Write down the sublevel name, magnetic quantum numbers, and total

number of orbitals allowed in each sublevel with the following quantum numbers:

(a) n = 2, l = 0

(b) n = 2, l = 1

(c) n = 3, l = 2

(d) n = 4, l = 2

(e) n = 4, l = 3

19. Determine the total number of allowed orbitals in each of the following atoms:

(a) 1H

(b) 9Be

(c) 16O

(d) 18O

(e) 27Al

(f) 40Ar

Filling of Electrons in Many Electrons Atoms & the Periodic Table

20. For each of the atoms listed in question 19 above, write the:

(a) Electronic configuration

(b) Condensed electronic configuration

21. The transition elements in period 4 of the periodic table are the first 10 elements to

fill up a d orbital:

(a) Write down the electronic configuration for 23V, 24Cr and 25Mn.

(b) Explain why the configuration for 24Cr does not follow the electron filling up

pattern followed by 23V and 25Mn.

(c) Give another example of a transition element in period 4 that does not follow

the electron filling up pattern of its nearby elements. Write down the electronic

configuration for this element and explain why the configuration is as such.

22. Write down the electronic configurations (full or condensed) for the elements below:

(a) 33As

(b) 37Rb

(c) 39Y

(d) 53I

23. Write down the electronic configurations for the ions below:

(a) 1H+

(b) 4Be2+

(c) 8O2

(d) 13Al3+

(e) 17Cl

(f) 26Fe2+

(g) 29Cu+

(h) 23V5+


24. Write down any two ions in question 8 above that are isoelectronic? Explain your

answer.

Consolidated Problems

25. For a hydrogen atom, electronic transitions can occur if it is heated.

(a) If the electron is promoted from n = 1 to n = 4 and then cooled, show all the

possible electronic transitions in the Lyman series and its relation to the line

spectrum formed.

(b) If the electron is promoted from n = 1 to n = , ion H+ will be formed.

(i) Use the Rydbergs equation to calculate the wavelength of the light required to cause the movement of the electron.

(ii) Calculate the energy required by an electron in this situation.

(iii) Calculate the energy required by 1 mole of hydrogen atoms for similar

transition.

(iv) Write an equation for the ionisation a hydrogen atom.

(v) What is the ionisation energy of 1 mole of hydrogen atoms?

26. Light of wavelength 162.7 nm is required to ionise magnesium atoms to magnesium

ions, Mg+.

(a) In which region is light of 162.7 nm?

(b) Write an equation for the ionisation.

(c) Calculate the energy required to change one magnesium atom to an ion, Mg+.

(d) If 1 mole of magnesium atoms are used, calculate the energy used.

[1 nm = 1 109 m, Speed of light, c = 3.00 108 m s1, Plancks constant,

h = 6.63 1034 Js, Avogadros number = 6.02 1023 mol1]

(e) What do you expect the ionisation energy to be if Mg2+ were to be formed from

Mg+?

27. An element W (not the real chemical symbol) has 3 protons.

(a) Write the electronic configuration for atom W.

(b) Determine the number of valence electrons in atom W.

(c) Determine the Group & Period in which atom W belongs to in the periodic

table.

(d) Describe how atom W can become an ion.

(e) Write down the electronic configuration for the ion formed in (d).

28. An element X (not the real chemical symbol) has 17 protons.

(a) Write the electronic configuration for atom X.

(b) Determine the number of valence electrons in atom X.

(c) Determine the Group & Period in which atom X belongs to in the periodic table.

(d) Describe how atom X can become an ion.


(f) Which noble gas is isoelectronic to the ion formed in (e) ?


29. An element Y (not the real chemical symbol) has 30 protons.

(a) Write the electronic configuration for atom Y.

(b) Determine the number of valence electrons in atom Y.

(c) Determine the Group & Period in which atom Y belongs to in the periodic table.

(d) Describe how atom Y can become an ion.


(f) Which noble gas is isoelectronic to the ion formed in (e) ?

(g) If the ions for elements X from question 28 and Y were to form an ionic

compound, write the chemical formula for the compound.


Tutorial 5

LU5 : Periodic Table of Elements

Information from electronic configurations of atoms / ions

1. Given the electronic configuration of some atoms as in table below:

Atom Electronic configuration

A 1s2 2s2 2p1

B 1s2 2s2 2p6 3s2 3p4

C 1s2 2s2 2p6 3s2 3p6 4s2

D 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p5

E 1s2 2s2 2p4

Arrange the atoms in order of increasing sizes.

2. Arrange the following chemical species in order of increasing sizes:

Na+ , Na, Cl, F, Ca

3. Given the electronic configuration of some atoms as in table below:

Atom Electronic configuration

U 1s2 2s1

V 1s2 2s2 2p6 3s2 3p4

W 1s2 2s2 2p6 3s2 3p6 4s2

X 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p5

Y 1s2 2s2 2p5

Arrange the atoms in order of increasing electronegativities.

4. What are the 2 factors that affect the first ionisation energy of an element?

5. Sulphur has a smaller atomic size as compared to phosphorus. Why sulphur has

smaller first ionisation energy?


Information from successive ionisation energies from an element

6. Successive ionisation energies can be used to determine the number of valence

electrons in an atom. Study the data in the table below:

Number of electrons removed 1 2 3 4 5

Ionisation energy / kJ mol1 800 2426 3659 25020 32820

(a) Determine the number of valence electrons in the atom.

(b) Determine the Group of the element.

7. Successive ionisation energies can be used to determine the number of valence

electrons in an atom. Study the data in the table below:

No. electron 1 2 3 4 5 6 7

Ionisation

energy/ kJ mol1 1310 3390 5320 7450 11000 13300 71000

(a) Determine the number of valence electrons in the atom

(b) Determine the Group of the element

8. The graph shows a plot of log (ionisation energy) against the number of the electron

removed for an element Z.

(a) What is the number of valence electron in element Z?

(b) Element Z is expected to be placed in which period and group of the Periodic

Table?

(c) Give the name of element Z.

9. Write an equation, including state symbols, to show the process that occurs when the

third ionisation energy of chlorine is measured.


Trends across a Period and down a Group

10. For the elements in Period 3, state the general trend in terms of:

(a) Atomic sizes

(b) First ionisation energies

(c) Electronegativities

(d) Melting points

11. For the elements in Period 3

(a) Write the formula of the oxides.

(b) When the oxides are dissolved in water separately, will it be acidic, basic, neutral

or amphoteric?

(c) Write the formulae of the chlorides.

12. For the elements in Group 2, what is the trend in terms of

(a) Atomic sizes


(c) Densities

(d) Melting points

(e) Electronegativities

(f) Thermal stabilities of the carbonates

(g) Solubilities of the sulphates

(h) Solubilities of the hydroxides

(i) Reactivities with oxygen

(j) Reactivities with chlorine

(k) Reactivities with water

13. Diagram shows part of the Periodic Table. Positions of six elements are shown by

using letters U, N, I, M, A, and S. These are not the symbols of the elements, but you

need to use them to answer the following questions.

(a) Which is the most electronegative element?

(b) Which is a monatomic element?

(c) Which is the most stable element?

(d) Which element is a semi-conductor?

(e) Which element forms strong metallic bond between them?

(f) Which element has the highest first ionisation energy?

(g) Which element has the biggest atomic size?

(h) Which element forms amphoteric oxide?

(i) Which element will reacts explosively with water?

(j) Which element has exceptionally high melting point?

U N

I M

A

S


Reactions of Group 2 elements

14. Write balanced equations (indicate the states also) for the following reactions:

(a) Reaction between Ca and water

(b) Mg and chlorine

(c) Ba and oxygen

(d) Na and water

(e) Heating of Mg(OH)2

(f) Heating of MgCO3

(g) Heating of Mg(NO3)2

15. A solution of 50.0 mL contains Barium ions (Ba2+) react with excess sulphuric acid to

produce 0.244 g BaSO4 (Relative molecular mass of 233).

(a) Write an ionic equation for the reaction.

(b) Calculate the concentration of the Ba2+ ion in the original solution.

16. 25.00 mL of M(OH)2 requires 0.10 mol L1 of 10.50 mL HCl to have complete

reaction. [Relative atomic mass: M = 137, O = 16.0, H = 1.0]

(a) Write a balanced equation for the reaction

(b) Calculate the concentration of M(OH)2 solution

(c) What is the mass of M(OH)2 that must be dissolved in 1 L of solution ?

17. Y is a mixture of two compounds of an element from Group 2. It gives a brick red

colour in flame test. When heated strongly, a reddish brown gas can be observed. Y

dissolves in water to give a colourless solution. Explain all the reactions and identify

Y.

18. When magnesium is burnt in air, a white powder is left. When some water is added

to the white powder and the solution tested with red litmus, a blue colour is seen.

Explain all observations and write chemical equations for the reactions involved.

Chemistry of Group 14 elements

19. For Group 14 elements, from C Si Ge Sn Pb, state how the general

variation changes down the Group.

(a) Atomic sizes

(b) Melting points

(c) First ionisation energies

(d) Metallic character

(e) +4 oxidation states

20. Indicate which of the compounds is more stable in each pair below:

(a) CO and CO2

(b) PbO and PbO2

(c) SnCl2 and SnCl4

(d) PbCl2 and PbCl4

21. Carbon cannot form complex ions while Si, Ge and Pb can. Explain the reason why

does this happen.


22. Two allotropes of carbon are shown below.

(a) Explain what is meant by allotropes of carbon.

(b) Name the allotropes of carbon in I and II.

(c) Explain why allotrope I is a good conductor of electricity but not allotrope II.

(d) Explain the hardness of allotrope I and II.

(e) What are the commercial uses of allotrope I and II?

(f) Name another 2 elements that can exhibit allotropy.

Chemistry of Group 17 elements

23. For the elements in Group 17, describe the trend in terms of:

(a) Atomic radius


(c) Electronegativities

(d) Melting points

(e) Solubility

24. Discuss the reactivity of halogens in term of displacement reactions between halogen

and halide ions.

25. Write the balance equations (indicate the states) for the following reactions:

(a) Phosphorus and excess chlorine

(b) Chlorine and iron(II) chloride solution

(c) Bromine bubbled through dilute sodium hydroxide

(d) Chlorine bubbled through concentrated sodium hydroxide at 70C

(e) NaBr(s) with concentrated sulphuric acid

(f) KI(s) with concentrated sulphuric acid

(g) NaCl(s) with concentrated sulphuric acid

(h) Aqueous solution of calcium chloride with silver nitrate solution

(i) Sodium bromide solution with silver nitrate solution

26. Which of the following reactions are classified as disproportionation reaction?

(a) I2(s) + KI(aq) KI3(aq)

(b) NH4NO2(s) N2(g) + 2H2O(g)

(c) 2ClO2(g) + H2O(l) HClO3(aq) + HClO2(aq)

(d) Cl2(g) + 2NaOH(aq) NaCl(aq) + NaClO(aq) + H2O(l)

(e) 3MnO42(aq) + 2H2O(l) 2MnO4

(aq) + MnO2(s) + 4OH(aq)


27. A solid compound releases a gas, P, when heated with concentrated sulphuric acid. P

is highly soluble in water to form a solution. When solution P is added to silver

nitrate solution, a white precipitate is formed. The white precipitate dissolves when

excess ammonia solution is added.

(a) Explain the chain of reactions.

(b) Write reactions for all the observations

(c) Name the gas P

Transitional elements

28. One of the characteristics of transition elements is the ability to show different

oxidation states in their compounds. Discuss this statement with reference to iron,

and explain the relative stability of the two main oxidation states of iron based on

their electronic structures.

29. Explain how the element chromium forms the complex ion, [Cr(H2O)6]3+, with water

molecules.

30. What are the factors that can determine the colour of a complex ion? Why complexes

of scandium and zinc are colourless?

31. Study the ionisation energies of the elements Sc to Zn as shown in graph below:

(a) Write the equation for the first, second and third ionisation energy for Mn.

(b) Why the 2nd ionisation energy Cr is higher than that of Mn?

(c) Why the 2nd ionisation energy of Cu is higher than that of Ni?

(d) Why the 3rd ionisation energy of Mn is higher than that of Fe?


Tutorial 6

LU6: Chemical bonding

Physical properties and its relation with bonding

1. Study the information provided in the table below:-

Substance Melting

Point/oC

Boiling

Point/oC

Solubility

in Water

Electrical

Conductivity

of Aqueous

Solution

Electrical

Conductivity

of Solid

Electrical

Conductivity

of Molten

Substance

T 43 23 Soluble Good Poor Poor

U 600 1190 Poor Poor Good Good

N 1700 2500 Insoluble Poor Poor Poor

E 780 1367 Good Good Poor Good

H 33 80 Insoluble Poor Poor Poor

(a) Which substances are likely to be solid at room temperature and pressure?

(b) Which substance is likely to be a liquid at room temperature

(c) Give a substance that is likely to be an ionic compound. Explain your choice.

(d) Is there a metallic element given in the table? Which one? Explain.

Lewis structure of a molecule

2. Draw a Lewis structure for each of the following molecules.

(a) CH4

(b) PCl3

(c) AlCl4

(d) NH4+

Accuracy of a Lewis Structure

3. Draw 2 possible Lewis structures of NOCl in accordance with octet rule. Then, select

the more likely structure of NOCl by means of formal charge.

Information from a Lewis structure of a molecule

4. Complete the table below:

Molecule/

Ion

Lewis structure

of the

Molecule /ion

No. of

electron pairs in

the central atom

No. of

bonded pairs

No. of lone

pairs in the

central atom

BeCl2

BF3

SO2

CH4

NH3


H2S

PCl5

SF4

ClF3

I3

SF6

BrF5

XeF4

Polarity of a covalent bond and dipole moment of a molecule

5. Indicate clearly the polarity in each of the following covalent bonds:-

(a) HI (b) HF (c) NH (d) BF

(e) NO (f) SO (g) SCl (h) HS

6. Study each molecule below, and decide whether there is a resultant dipole moment.

Then decide whether the molecular is polar or not?

(a) HBr (b) H2O (c) BCl3 (d) CH4

(e) CO2 (f) HCHO (g) CH3Cl (h) NF3

Valence Shell Electron Pairs Repulsion Theory to determine shape, bond angle

and dipole moment of a molecule


Molecule

No. of

electron pairs in

the central atom

Orientation

Of electron

Pairs

Shape of

molecule

Any resultant

dipole moment

Is the

molecule

polar?

BeCl2

BF3

SO2

SiCl4

NH3

H2S

PCl5

SF4

ClF3


8. Arrange the bond angles for the following compounds in ascending order:

(a) HCN (b) CO2 (c) H2O

9. Arrange the bond angles for the following compounds in ascending order:

(a) CCl4 (b) H2O (c) BCl3 (d) CH3

Co-ordinate bonding (Dative bonding)

10. When ammonia (NH3) is mixed with boron trifluoride (BF3), a compound H3N-BF3 is

formed.

(a) Show the co-ordinate bonding in the final product.

(b) What do you expect the orientation of the electron pairs in B?

11. Aluminium chloride (AlCl3 ) is a covalent compound. When ammonia (NH3) is mixed

with aluminium chloride, a compound H3NAlCl3 is formed.

(a) Show the co-ordinate bonding in the final product.

(b) What do you expect the orientation of the electron pairs in Al?

Valence bond theory

12. By using the Valence Bond Theory, draw diagrams to indicate the bonds formed.

(a) H2 (b) HCl (c) O2 (d) N2

13. Count the number of and bonds in each of following compounds:

(a) CH3CH3 (b) H2O (c) BCl3 (d) CH2=CH2

(e) CO2 (f) HCHO (g) CH3Cl (h) HCN

Hybridisation theory


Molecule /

ion

No. of

bonds

No. of lone

pairs in the

central atoms

Type of

hybridisation of

the central atom

Shape of

molecule

CH4

NH3

H2O

BeH2

BF3

CO2

CH3Cl

SO2

SO3


15. State the type of hybridisation in each of the carbon atoms below:

(a) H3CCH2CH=CH2

(b) H2C=CHCH=CH2

O

(c) HCCCH2COH

16. Count the total number of and bonds in each of the molecules below:

(a) H3CCH2CH=CH2

(b) H2C=CHCH=CH2

(c) HCCCH2COOH

Inter-molecular attractive forces

17. Show how hydrogen bonding exists in the following molecules:

(a) HF (b) H2O (c) NH3 (d) CH3OH

18. Relative molecular masses of H2O and H2S are18 and 34 respectively. The lighter

H2O is a liquid, whereas the heavier H2S is a gas. Explain.

19. Van der Waals forces are forces of attraction between covalent molecules. Write

short notes on the following:

(a) Permanent dipoles

(b) Induced dipoles

(c) Instantaneous dipoles (London forces or dispersion forces)

20. Arrange the strength of inter chemical species attractive forces in order of ascending

order:

(a) Ionic bonding

(b) Hydrogen bonding

(c) Van der waals forces

21. Which has the highest boiling point? Explain

H2S, HI, H2O

22. Which has the highest melting points? Explain

CO, H2S, CH3OH, CH3CH2OH

23. Discuss solubility with suitable example in situations like

(a) An ionic substance in water

(b) A polar substance in water

(c) A non polar substance in a non polar solvent

24. Explain in terms of bonding or intermolecular forces for the following observations:

(a) Ethanol, C2H5OH, mixes with water in all proportion, but ethane, C2H6, does not

mix with water.

(b) Relative molecular mass of ethanoic acid in benzene is found to be 120.



25. Potassium, Silicon, Trichloromethane, Helium, Sodium chloride

From the above list, chose a substance that shows the following characteristic:

(a) Monoatomic with Van der Waals forces.

(b) Macromolecule consists of atoms held together by covalent bonds.

(c) Polyatomic molecule with high boiling point.

(d) Solid that conducts electricity in molten state and in aqueous solution.

(e) Substance containing delocalised electrons.

26. Explain in terms of bonding and structure for the following observations:

(a) Magnesium chloride is a solid with high melting point.

(b) Tetrachlomethane does not form precipitate with aqueous silver nitrate.

(c) Sodium chloride dissolves in water but does not dissolve in benzene.

(d) Ice has lower density than water.

27. Using a suitable example to explain each of the following:

(a) An ionic bond

(b) A covalent bond

(c) A co-ordinate bond

(d) Metallic bonding

(e) A hydrogen bonding

28. Ammonia (NH3) is a covalent compound.

(a) How many valence electrons are there in a molecule of NH3?

(b) Show the Lewis structure of NH3

(c) How many lone pair of electrons are there in the central atom of NH3?

(d) How many electron pairs are there in the central atom of NH3?

(e) What type of orientation for the electron pairs in (d)?

(f) What is the shape of a NH3 molecule?

(g) What is the bond angle in NH3?

(h) Show the resultant dipole moment in NH3.

(i) Phosphorus (P) and Nitrogen(N) are both Group 15. What do you expect the

shape and dipole moment in PH3 ?

(j) RMM of NH3 is smaller compared to PH3, but NH3 has a higher boiling point than

PH3. Explain.


Tutorial 7

LU7: Organic chemistry

Differentiate organic compounds from inorganic compounds

1. Which of the followings are considered as organic compounds? For each organic

compound, name the homologous series it belongs and identify its functional group.

(a) CH3CH3 (b) CH3CHClCH3 (c) CH2=CH2

(d) CH3CHO (e) C6H5CH2Br (f) CaCO3

(g) HCHO (h) H2SO4 (i) CH3CH2COOH

(j) C10H8 (k) CuO (l) C6H5COOH

(m) CH3CH(CH3)COCH2CH3 (n) HCOOH (o) NH3

(p) CH3CH=CH(CH2)12CH3 (q) NaBr

Relation between empirical formula and molecular formula

2. What is the empirical formula of each compound below:

No. Name of

compound

Actual formula of

compound Empirical formula

(a) Butane CH3CH2CH2CH3

(b) But-1-ene CH3CH2CH=CH2

(c) Hex-2-ene CH3CH2CH2CH=CHCH3

(d) Methyl propanoate CH3CH2COOCH3

(e) Butanoic acid CH3CH2CH2COOH

3. A hydrocarbon, Q, has an empirical formula of CH2. It has a molar mass of 56.

(a) Determine the molecular formula of Q.

(b) Draw a displayed formula of Q.

(c) Write the structural formula of Q in (b).

(d) Show the skeletal formula of Q in (c).

Ways to represent structural formula

4. Draw any two possible displayed formulae for each of the following compounds.

(a) C4H10 (b) C3H8O (c) C3H7Br (d) C4H8O

5. Draw all the displayed formula for C4H8.

6. Show the skeletal formulae for the compounds below.

(a) CH3CH2CH2CH3 (b) CH3CH=CHCH3

(c) CH3CH2COCH2CH3 (d) CH3CH=CHCH2OH


7. Write a structural formula for each of the following and then draw the skeletal

formula of it.

(a) C4H10 (b) C3H8O (c) C3H7Br

(d) C4H8 (e) C3H9N

8. Given the molecular formula of C5H10, show at least 3 displayed formulae, and then

show the skeletal formula for each. (you may go beyond the open chains)

9. Draw the 3-D structural formulae for the following compounds.

(a) CH3F (b) CH2Cl2 (c) CH3OH (d) CF4

Nomenclature of organic compounds

10. Name the following according to IUPAC systems.

(a) CH3CH2CH2CH3 (b) CH3C(CH3)2CH2CH2CH3

(c) CH3(CH2)3CH(CH3)CH3 (d) CH3CH(CH2CH3)2

11. Give the structural formulae for the following compounds.

(a) 2,5-dimethylhexane (b) 2,3-dimethyloctane

(c) 2,2,3,3-tetramethylpentane (d) 4-ethyl-3,4-dimethylheptane


(a) (b)

(c) (d)

13. Give the IUPAC names for the following compounds.

(a) CH3CH2CHBrCH3 (b) CH3C(CH3)2CHClCHClCH3

(c) CH3(CH2)3CCl(CH3)CH2Br (d) CH2ICH(CH2CH3)2

(e) (f)


(a) CH3CH2CH=CH2 (b) CH3C(CH3)2CHCHCH3

(c) CH3(CH2)3C(CH3)CH2 (d) CH2C(CH2CH3)2

15. Give the structural formulae for the following compounds.

(a) 2,5-dimethylhex-1-ene (b) 2,3-dimethyloct-3-ene

(c) 2,3-dimethylpent-1-ene (d) 3,4-dimethylhept-1,3-diene

(e) 3,3-dimethylbut-1-ene

CH3

CH3

CH3

Cl

CH

Br

CH3


16. Draw displayed formulae for the following compounds.

(a) 2,3-dimethylpent-2-ene

(b) 1,2-dimethylcyclopentene


(a)

(b)

(c)

CH3H3C

(d)

(e)

Cl

(f)

Isomerism in organic compounds

18. Using the molecular formula of C4H10O, explain the meaning of chain isomers,

positional isomers and functional isomers.

19. Which of the following alkenes will exhibit cis-trans forms?

(a) CH3CH2CH2CH=CH2 (b) CH3CH2CH=CHCH3

(c) CH3CH2C(CH3)=CH2 (d) CH3CH(CH3)CH=CH2

(e) CH3CH=C(CH3)2 (f) CH3CH2CH2Cl

(g) CH2=CHCH2Cl (h) CH3CH=CCl2

(i) CH3CCl=CClCOOH (j) CH3CHClCH2CH3

20. Identify the chiral carbon atom(s) in each of the following compounds.

(a) CH(NH2)(CH3)CO2H (b) CH3CHBrCH2CH3

(c) C6H5CH2CH3 (d) C6H5CH(OH)CN

21. Study each compound below and identify the chiral carbon atom (if any) and state

whether the compound is optically active or not.

(a)

(b)

(c)

CH3

Br

CH3

CH3CH2CH2 CH CH2OH

CH=CH2

CH3CH CH

Cl Cl

Cl

CH CH3

Cl


22. Which of the following compounds can show cis-trans isomers?

(a) CH2=CHCl (b) CH3CH=CH2

(c) CH3CH=CHCH3 (d) CH3(CH2)7CH=CH(CH2)12CH3

(e) C6H5CH=CH2 (f) CH3CHBrCHBrCH3

(g) HOOCCH=CHCOOH

Chemical reactions of alkanes and alkenes

23. Show clearly the free radical reaction mechanism for the chlorination of methane.

24. Show clearly the free radical reaction mechanism for the bromination of ethane.

25. Write an equation for each of the following chemical reactions.

(a) Combustion of ethene

(b) Ethene and bromine.

(c) Ethene and hydrogen

(d) Ethene and hydrogen bromine.

(e) Ethene and cold dilute KMnO4 solution

26. Write the addition reaction mechanism between ethene and bromine.

27. Write the addition reaction mechanism between cyclohexene and bromine.

28. There were two bottles of liquid, cyclohexane and cyclohexene. Unfortunately, the

labels were missing. How would you confirm the identity of the liquids? Give 2 tests.

29. What would be the structure of the product if HI is reacted with each of the

following?

(a) CH3CH2CH=CH2 (b) CH2=C(CH3)CH2CH3

(c)

(d) CH3

(e) CH3CH=C(CH3)CH3

30. Write the structure of the polymerised product for each of the following monomers.

(a) CH2=CHCN

(b) CH2=C(CH3)COOCH3

(c) Propene

(d) C6H5CH=CH2

(e) Buta-1,3-diene


Nomenclature of arenes and their derivatives

31. Name the followings compounds systematically.

(a) CH3

(b)

(c) CH3H3C

(d)

(e)

CH3

C2H5

(f)

(g)

C2H5

NO2

(h)

32. Draw the displayed formulae for the followings.

(a) p-Dinitrobenzene

(b) 2,3-Dinitro-4-chlorotoluene

(c) 2-Amino-5-bromo-3-nitrobenzoic acid

(d) 2,4,6-Trinitrophenol

Reactions of arenes

33. Draw the electron cloud to show the delocalised electrons in the benzene ring.

34. The reaction between benzene and bromine gives bromobenzene as the product with

FeBr3 as the catalyst. Write the reaction mechanism.

35. What is the structure of the product when CH3Cl reacts with benzene in the

presence of AlCl3? Write the reaction mechanism.

36. What is the structure of the product when concentrated nitric acid reacts with

benzene in the presence of concentrated sulphuric acid? Write the reaction

mechanism.

37. What is the structure of the product when CH3COCl reacts with benzene in the

presence of AlCl3?

38. What is the structure of the product when toluene boiled for a prolonged period with

acidified KMnO4?

39. What is the structure of the product when m-xylene boiled for a prolonged period

with acidified KMnO4?

C2H5

CH3

NO2

C2H5

C2H5

NO2

NO2


40.

+ YAlCl3

C(CH3)3

What reagent could be Y?


41. Propane is a saturated hydrocarbon that reacts with chlorine according to chemical

equation shown below:

CH3CH2CH3(g) + Cl2(g) CH3CH2CH2Cl(g) + HCl(g)

The reaction is known as a free radical substitution reaction that can occur in the

presence of sunlight.

(a) What is meant by a saturated organic compound?

(b) What is the purpose of sunlight in the reaction?

(c) The free radical substitution reaction is known to be carried out in steps. Write

clearly the steps involved.

(d) Regarding the product CH3CH2CH2Cl,

(i) What is the IUPAC name?

(ii) Draw the displayed formula.

(iii) Draw the skeletal formula.

(iv) Give the possible isomers and name them systematically.

(e) From the view point of chemical properties, give one similarity and one difference

between propane and propene.

42. Benzene is an unsaturated hydrocarbon. It has typical reactions quite different from

an alkene.

(a) What is the typical reaction of benzene?

X

(b) C6H6 + CH3CH2Cl Y + HCl

(i) State what is X and Y.

(ii) Give the IUPAC name of Y

(iii) Give 2 isomers of Y and name them systematically

(iv) If Y is vigorously oxidised, what will be the product?

(v) Name the reagents to carry out the reaction in (iv).

(vi) Count the number of bonds in Y

(vii) State the type of hybridisation in the carbon atoms in Y.

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Tutorial LU1-7 Sem 1

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