LC Revision 2014/15

1. Atomic Theory: Chapters 2,3,4,5,6,7,8,9,112. Organic Chemistry: 21,22,233. Volumetric: 13,15 (19)4. Acids, Bases, pH & Indicators: Chapter 12, 185. Chemical Equilibrium: Chap 176. Oxidation & Reduction: Chap 147. The Gas Laws: Chap 108. Rates of Reaction: Chap 169. Water: Chap 1910. Electrochemistry: Chap 2011. Thermochemistry: Chap 2112. Q.4 & Options

Section 1 Atomic Theory

Atomic Theory covers Chapters 2,3,4,5,6,7,8,9 & 11. This involves a lot of work but the topics are short and easy to master with a little practice.

1 The Periodic Table2 Radioactivity3 Electronic Structure of atoms4 Bonding & Electronegativity5 Stoichiometry6 Shapes of molecules 7 Mole

Chapter 2 Atomic Theory

A brief history of the atom also with special attention to Dalton, Crookes, Thompson, Millikan, Rutherford and Chadwick

Q.1 What is meant by diffusion? Give e.g. Q.2 What is formed when hydrochloric acid and ammonia react?Q.3 Who was the first man to put forward an atomic theory? Q.4 What material did Rutherford hit with alpha particles in his famous experiment that led to the discovery of the nucleus? Q.5 What is the mass of a neutron? Q.6 What was Dalton’s theory?Q.7 Who studied the discharge of electricity through gases?Q.8 What are cathode rays?Q.9 Draw a cathode ray tube?Q.10 What name is given to the negative/positive end of a battery?Q.11 Who discovered the electron?Q.12 What does e/m mean?Q.13 Who devised an expt. to measure the charge on the electron? Q.14 What did Thomson’s model of the atom look like?Q.15 Describe Rutherford’s expt. Give his findings.Q.16 Who discovered the proton?Q.17 Who discovered the neutron?Q.18 Describe the expt.that led to the discovery of the neutron?Q.19 What is the mass of the proton/the electron?Q.20 What did Thomson discover about cathode rays?Q.21 What contribution did Stoney make in the discovery of the electron?Q.22 Compare the mass, charge and location of the proton, neutron and electron.

Atomic Structure - History of Development

Democritus Atoms = indivisible – 4 elements earth, wind, fire and water

Dalton John Atomic Theory - element made of tiny particles – all particles of element identical – compounds combinations of these particles

Crookes Maltese Cross Experiment – Cathode Rays – radiometer

Thompson Deflected cathode rays with charged plates – must be negativeCalled them electrons [after George Stoney of NUIG]

Magnet also deflects raysWorked out Charge/Mass ratioPlum Pudding model

Millikan Mass of Electron Oil Drop Experiment – charged oil drops with X-rays – measured charge to stop them falling – used this with charge/mass ratio to work out mass

Rutherford Gold Foil Experiment - Discovered nucleus – very small, dense, positively charged – deflected alpha particles

Moseley Each element has a characteristic positive charge – Atomic NumberElement substance whose atoms all have same atomic number

Bohr Niels Electrons orbit nucleus in shells – energy quantised – only certain values allowedMaximum number of electrons in shell = 2n2

Rutherford Called positive particles ProtonsSuspected other particles to cement nucleus together

Chadwick Showed Neutrons by bombarding Be with alpha particles

Sub-atomic ParticlesParticle Mass Charge Location

Proton 1 +1 NucleusNeutron 1 neutral NucleusElectron 1/1838 -1 Orbiting nucleus

Atomic Number = Number of Protons This determines what the element is.

Mass Number = Number of Protons plus the number of neutrons

Relative Atomic Mass is the average mass of the element as it occurs in nature [when the isotopes and abundance are taken into account) and compared to C12 isotope.

Number of Protons = Atomic Number

Number of Electrons = number of protons = atomic number [in atoms]In ions add an electron for each minus charge and subtract an electron for each positive charge

Number of Neutrons = Mass Number – Atomic Number

IsotopesAtoms of an element which have the same atomic number but different mass numbers due to different numbers of neutrons

Not all isotopes are radioactive but 14C is

Chapter 3: Arrangement of Electrons in an atom

Electronic Structure of AtomsElectronic configurations of atoms and ions must be known here along with the shapes of s and p orbitals.Atomic emission and atomic absorption spectra can be asked so be able to distinguish clearly between the two. Know the ‘flame’ test experiment and the colours produced by the metals. e.g. potassium – lilac.

Learn Bohr diagrams for first 20 elements

k = 2, l = 8, m = 8, n = >2

Atomic spectra

3 types of spectra

Continuous – from sun

Line – Emission – from heated or shocked elements – Absorption – frequencies removed as light passes through gaseous element

Emission Spectrum – how formed

- Electron normally in lowest available energy level – Ground State – E1- Electron excited by heat or electricity [high voltage]- Jumps to higher energy level – Excited State – E2

- Unstable- Drops back to lower level- Energy released as photon or Electro Magnetic Radiation- Frequency proportional to drop - (E2-E1) = hf h is Planck’s Constant and f is frequency of radiation- lines formed are evidence of energy levels

Electron Configuration

n is the main energy level or Quantum Number

Number of electrons that a main energy level [shell] can hold is 2n2 [n = main energy level number]later found that main energy levels were divided into sublevels [up to 4 of these s, p, d and f.sublevels have different energies

n = 1 1sn = 2 2s, 2pn = 3 3s, 3p, 3dn = 4 4s, 4p, 4d, 4fn = 5 5s, 5p, 5d, 5f

Heisenberg’s Uncertainty Principle

The more accurately we determine the position of an electron the less accurately we can determine its velocity.

Electrons have Wave-particulate duality Have properties of both waves and particles at same time

Sublevels further divided into orbitals of equal energy

Each s sublevel contains 1 s orbital – spherical in shape

Each p sublevel contains 3 p orbitals - px, py, pz – each dumbbell shaped and mutually at right angles

Each d sublevel contains 5 d orbitals

Each f sublevel contains 7 f orbitals

Orbital Region in space around the nucleus of an atom in which electrons are most likely to be found

Each orbital can hold up to 2 electrons

- Each s sublevel can hold up to 2 electrons

- Each p sublevel can hold up to 6 electrons

- Each d sublevel can hold up to 10 electrons

- f sublevel

Electron configuration [pattern]

Aufbau PrincipleElectrons occupy the lowest available energy level

Order of filling

1s2s, 2p3s, 3p, 3d4s, 4p, 4d, 4f5s, 5p, 5d, 5f

Number of electrons [in an atom] = atomic number

Learn patterns for first 36 elements

e.g. Calcium Atomic number 20 therefore 20 electrons

1s2, 2s2, 2p6, 3s2, 3p6,4s2

Exceptions

Chromium [4s1 3d5] and Copper [4s1 3d10]

Half full and full d sublevel is more stable

Electron configuration of ions

Each electron lost = 1+ chargeEach electron gained = 1- charge

Put pattern in square brackets with charge outside

e.g. Calcium ion Ca2+ Lost 2 electron thus 2+

[1s2, 2s2, 2p6, 3s2, 3p6]2+

Only patterns of ions of first 20 elements required

Arrangement of electrons in orbitals of equal energy

Hund’s Rule of Maximum MultiplicityWhen two or more orbitals of equal energy are available, the electrons occupy them singly and then in pairs.

Final point is electrons spin in opposite direction

Pauli Exclusion PrincipleNo more than two electrons may occupy an orbital and to do so they must have opposite spin

Each electron can be specified by 4 numbers

- Principal Quantum Number- Sublevel - Orbital- Spin

1. Who was the first scientist to give information about the arrangement of electrons in an atom?

2. What is a continuous spectrum/line spectrum?3. Name two ways in which spectra can be seen.

Na Mg Al Si P S Cl Ar

Na

Mg

Al

Si

PS

Cl

Ar

c

Mg higher than general trend due to full s sublevel and P is higher than general trend due to a half- filled p sub-level

Na Mg Al Si P S Cl Ar

Na

Mg

Al

Si

PS

Cl

Ar

c

Mg higher than general trend due to full s sublevel and P is higher than general trend due to a half- filled p sub-level

1st Ionisation energies

4. Describe an expt. to investigate the flame colours of different salts.5. Name two parts of a spectrometer.6. What colour is emitted by lithium, potassium, barium, strontium, copper and

sodium?7. What salt commonly causes contamination?8. How do you reduce contamination?9. What is an energy level?10. What is a quantum of energy?11. What is Heisenberg’s Uncertainty Principle?12. Who worked out the likely probability of finding a particular electron in an

atom? (mathematically)13. What is the shape of an s/p orbital?14. How many type of p orbital exist? Draw, give letters to label each.15. What is the ground state of an atom?16. How do you work out how much energy is emitted when an electron falls

back down to the ground state, from its excited state.17. Give detailed description (using diagrams) of how elements are able to

produce their own particular line spectra. 18. What is the Balmer series?19. What is the definition of an orbital?20. What did De Broglie say about moving particles?21. What is an absorption spectrum? Give some of its uses (two)22. What does LASER stand for?23. Give another example of a piece of equipment that makes use of electron

transitions.24.What does ‘h’ stand for in the equation relating energy of light and its frequency.25. Name the instrument used to study emission line spectra? 26. Give another word to explain ‘quantised’, with reference to energy of an electron?27. E = hf. What does ‘f’ stand for?

28. What electron transitions give rise to lines in the visible spectrum? 29. What is the maximum number of electrons that can occupy energy Level 3? 30 Which of the following orbitals is spherical: ‘s’, ‘p’, ‘d’ or ‘f’31. Give the name of two ‘heavy’ metals that might be found by Atomic Absorption Spectrometer in water analysis? ****************************

Chapter 4: Periodic Table

The Periodic Table – Know the contributions made by Dobereiner (triads), Newlands (octaves), Mendeleev and Mosely.. Know the difference between Atomic number and Relative Atomic Mass, group and period, metals and non-metals and gases. Be familiar with the following groups – alkali metals, alkaline earth metals, halogens and noble gases.Calculation of the relative atomic mass from the percentage isotopes should be practised.

1. Give explanation and example of how Dobereiner and Newlands grouped elements.

2. Why would Newland’s classification not work today? Give the main reason.

3. In what order did Mendeleev arrange elements?4. Who changed this order?5. Give three differences between Mendeleev’s table and the modern day one?6. What is meant by ‘eka silicon’?7. Who discovered sodium and potassium?8. What is the Periodic Law?9. How many (a) electrons (b) neutrons are in 2311Na+

10. What is Relative Atomic Mass?11. What is a Mass Spectrometer used for? (Give 2)12. Who built the first Mass Spectrometer?13. Give the Principles behind the operation of the Mass Spectrometer.14. Name the parts and be able to draw the instrument.15. What is an isotope? Give an e.g.16. Calculate the relative atomic mass of an element give its mass and %

abundance.17. Write the electronic configuration of an element..18. Write the s,p configuration of an element..19. What is the order in which the sub levels are filled?20. What does isoelectronic mean?21. What is Pauli’s Exclusion Principle?22. What is Hund’s Rule?23. What is the Aufbau Principle?24. How many electrons can be accommodated in the p orbitals/d orbitals?25. Why is the arrangement of electrons in potassium 2,8,8,1 and not 2,8,9,? 26.Write the s,p, configuration for Copper and Chromium? What is different about

these?

Periodic Table - Atomic Theory Summary

Boyle Robert Earl of Meath – Elements cant be broken down into anything simplerBoyles Law – volume inversely proportional to pressure at constant temperature

Lavoisier Anton

Listed known elements Law of conservation of Mass – matter neither created nor destroyed in a chemical reaction simply rearranged.

Davy Humphrey

Discovered new elements Na, K Ca etc using Electrolysis

Dalton John Atomic Theory - element made of tiny particles – all particles of element identical – compounds combinations of these particles

Dobereiner Triads – groups of 3 elements – middle one intermediate

Newlands Law of Octaves – properties repeat every eighth elements when arranged in order of mass – no noble gases

Mendeleev Dmitri

Periodic Table – Arranged by property space rather than weight – left gaps – predicted properties of elements to go in gaps – Germanium

Moseley Henri

Atomic number – number of protons

Modern Table v. Mendeleev’s

1. Arranged by atomic number2. More elements in modern – Noble Gases3. No gaps – Mendeleev left gaps to make elements fit into proper

column – In a few cases reversed the order of elements so they fitted into groups with similar properties

4. Transition elements in a separate block

Symbols *Symbols of first 36 elementsGroups Vertical columns, Gp I – alkali metals; Gp II – alkaline earth metals ,

Gp VII – halogensGp 0 – Noble gases – same number of electrons in out shell – similar properties

Periods Rows – outer shell filling

Blocks s, p, d [transition elements] , f

DEFINITIONS Chapters 2,3,4.

AtomSmallest particle of matter that can exist by itself

MatterAnything that occupies space

ElementSubstance made up of one type of atom – can’t be broken into anything simpler by chemical means

MoleculeSmallest particle of substance that shows properties of that substanceGroup of atoms chemically joined

IsotopesForms of element with different mass number due to different numbers of neutrons

The Mole 1 mole = 1 mole

= RMM (relative molecular mass) in grams = Avogadro’s number (6*1023)= 22.4L of any gas at STP [ 273 K and 760mm of Hg ]

Avogadro’s Law 1 mole of any gas at STP occupies 22.4 L

Atomic NumberNumber of protons in an atom. Determines what the element is.

Mass number Number of Protons + neutrons in an atom

Relative Atomic Mass1. average of the mass numbers of the isotopes of the element.2. as they occur naturally3. taking their abundances into account

4. expressed on a scale on which atoms of the carbon 12 isotope have a mass of exactly 12 units.

Relative Molecular Mass1. The sum of the relative atomic masses of all the atoms in a molecule of the

compound.2. The mass of one molecule of that compound compared with one twelfth of

the mass of the carbon 12 isotope.3. Mass of one mole of a compound = Relative Molecular Mass in grams.

Energy LevelThe fixed energy value that an electron in an atom may have.

Atomic OrbitalThe region in space within which there is a high probability of finding an

electron.

Hund’s Rule of Maximum MultiplicityWhen 2 or more orbitals of equal energy are available, the electrons occupy them singularly before filling them in pairs.

Aufbau PrincipleElectrons occupy the lowest available energy level.

Pauli Exclusion PrincipleNo more than 2 electrons may occupy an orbital and they must have opposite spin.

Heisenberg’s Uncertainty PrincipleThe more accurately we know the position of a particle the less accurately we know its velocity.

Bonding & ElectronegativityKnow the concept of electronegativity well i.e. the attraction of an atom for electrons. Link electronegativity with different types of bonding on the course. Be able to discuss the trends as you go across a period and down a group in the periodic table. Intramolecular bonding includes: Pure covalent, Polar covalent and IonicIntermolecular forces that occur between molecules are: Dipole-dipole, Hydrogen bonding and Van der Waals.

Shapes of Molecules – You need to be able to explain the shapes of molecules containing 2,3,4 and 5 atoms using the VSEPR theory. Know the repulsive forces between bonding pairs and lone pairs of electrons and how the presence of a lone pair affects its shape.

2 atoms in a molecule – linear (180o)3 atoms in a molecule – linear (no lone pairs – 180o) or V – shaped (has lone pairs – 104.5o)4 atoms in a molecule – triangular planar (no lone pairs – 120o) or pyramidal (has lone pairs – 107o)5 atoms in a molecule – tetrahedral (no lone pairs – 109o)

Chapter 5: Bonding & Shapes of molecules

1. What is meant by the octet rule?2. Who were the 2 American chemists that proposed this rule to predict how

elements combine together?3. Give 2 exceptions to the octet rule.4. What is an ion?5. What is an ionic bond?6. What is the formula for calcium fluoride? Lithium sulphide?7. Give the formulae for the following complex ions: carbonate, ammonium,

sulphate.8. Why do transition elements have variable valency?9. What does –ate mean at the end of a formula?10. What is the definition of a transition metal?11. Why are scandium and zinc not considered true transition elements?12. What is a molecule?13. What is a ‘lone pair’?14. What element is the standard by which valency is measured?15. What is meant by the term valency?16. What is a sigma/pi bond?17. List 3 properties of a covalent compound.18. What is recrystallisation?19. Give two benefits of knowing a compound’s melting point.20. What does VSEPR mean?21. Define electronegativity.22. Who set up the electronegativity scale?23. If electronegativity values are between 0 and 0.4 what does this tell you

about the type of bonding?24. Name 2 polar substances.25. What is the shape of the ammonia molecule?26. Name a non-metal that exists as single atoms.27. Which 2 of the first 36 elements would form the compound with the

greatest ionic character?28. What type of bond exists between the molecules of hydrogen chloride.

List the 3 types of intermolecular forces

Chapter 6: Anion Tests

1. What is the Law of Conservation of Mass?2. Be able to balance simple equation.3. What is an anion?4. What is the name given to a material that settles out of solution?5. Give reagents/chemicals needed to test for all 7 anions. (Usually 2)6. Give equations where necessary.

Chapter 7: Trends

Ionisation Energy

First Ionisation EnergyThe energy required to completely remove the most loosely bound electron from a neutral gaseous atom in its Ground State.[ X – e- = X+ ]

TrendsAcross Table – goes up – bigger nuclear charge – same distance from nucleus

Down table – goes down – bigger nuclear charge but further from nucleus so more shielding

Jumps at various places – Group V elements due to half filled sub-level being more stable and Group 0 elements Filled sublevel being very stable

Second Ionisation EnergyThe energy required to completely remove the second most loosely bound electron from the ion X+ – e- = X2+

Na Mg Al Si P S Cl Ar

Na

Mg

Al

Si

PS

Cl

Ar

c

Mg higher than general trend due to full s sublevel and P is higher than general trend due to a half- filled p sub-level

Electronegativity

The relative attraction that an atom in a molecule has for the shared pair of electrons in a covalent bond.

Invented by Linus Pauling

Difference in electronegativity and type of bonding

= 0 – 0.4 Pure Covalent bond > 0.4 < 1.7 Polar Covalent

> 1.7 Ionic

TrendsAcross – increases – bigger nuclear charge

Down – decreases – further from nucleus – more important than increased nuclear charge

Group O – none because do not form bonds

Atomic Radius

Half the distance between 2 adjacent atoms of the same element joined by a single covalent bond

TrendsAcross table – decreases – greater nuclear charge – same shielding effect

Down table – increases – new shell added – so more shielding – so increased nuclear charge not as significant.

Noble gases don’t bond so NO atomic radius

Chapter 7: Trends in the Periodic Table

1. Why is there a decrease in (i) the atomic radius, (ii) the first ionisation energy, in going from the element of atomic number 12 to element number 13?

2. Explain why the value of the first ionisation energy of oxygen is lower than that of fluorine and is also lower than that of nitrogen.

3. Electronegativity values __________gown the groups and ________across the periods.

4. The most reactive halogen is _____________5. Which one of the halogens is a liquid at room temperature?6. Sodium reacts with water to form _______ and _________

Measure of the attractiveness of an atom for electrons in a shared pair

7. Write the equation for the second ionisation energy of element X.8. Define atomic radius.

Radioactivity Chapter 8

Radioactivity – There are three types of radioactivity: alpha, beta and gamma. Be able to write brief notes on each type of radiation using the following guidelines.

1 What each type consists of2 Charge3 Symbol4 Effect on atomic number5 Effect on Mass number6 Penetrability

Radioactivity discovered by Henri Becquerel

Curies discovered two new elements

Polonium 84 [after Poland] and Radium 88

Radiation types and properties

Type Composition Penetrating Power

Source Use

Alpha particle

2 P + 2 N Few cm of air Americium 241

Smoke alarms

Beta particle

Electrons 5 mm Al Carbon 14 Carbon dating

Gamma rays

EMR Few cm lead Cobalt 60 Killing tumours

Alpha Particle Emission Equation 22688 Ra ® 22286 Rn + 42 He Beta Particle Emission Equation 146 C ® 147 N + 0e- Half-life

Uses of Radioactivity Carbon Dating

Medical – killing tumours, sterilising equipment

This is the time taken for half of the nuclei in any given sample to decay

Agricultural – killing insects Food Irradiation – increases shelf life

Industrial - finding leaks

Research – radioactive IAA trace movement

All around usRadon in granite – causes cancer if inhaled for a long period

1. Who discovered radioactivity?2. Who discovered polonium and radium?3. What is radioactivity?4. What are the three types of radiation?5. What instrument is used to detect radiation?6. What units is radioactivity measured in?7. What is half-life?8. What is an alpha particle? Beta particle?9. How can they be stopped?10. Give a source of each of the three particles.11. What is a radioisotope?12. Complete nuclear equations to show the action of alpha and beta particles.

Chapter 9: The Mole

1. Give a definition of a mole.2. What is the Relative Molecular Mass of a compound?3. What is Avogadro’s number?4. Calculations

- What is the mass of one mole of sodium?... 23g- What is the mass of a molecule of oxygen? 32g- What is the Relative molecular mass of water?.. 18g- How many moles are in 10g of sodium?..10/23 = .43moles- How many atoms are in 10g of sodium?..

10/23 = .43 x 6 x 1023

= 2.5 x 1023

5. What mass of chromium has the same number of atoms as 8g of calcium?8/40 = 0.2moles 1 mole Cr = 52g 0.2moles Cr = 10.4

6. How many atoms are present in 0.12g carbon. 0.12/12 = 0.01moles 1 mole carbon = 6 x 1023

0.01 x 6 x 1023

7. How many moles are in 2.6 x 1012 atoms of sodium? 6 x 1023 = 1 mole1 atom of sodium = 1/6 x 1023

2.6 x 1012 = 2.6 x 1012

________ 6 x 1023

= 0.43 x 1012-23

= 4.3 x 10-12

Relative Molecular Mass of a substance is the mass of one molecule of that substance compared with one twelfth of the mass of the carbon-12 isotope

1. Find Mr of CuSO4

Answer: 159.5

2. Find Mr of Mg (NO3) 2

Answer: 148g

A Mole is a measure of the amount of a substance

One mole of a substance is the amount of that substance that contains 6 x 10 23

particles of that substanceThis number 6 x 10 23 is known as the Avogadro Constant

A mole of a substance is either its relative atomic mass or its relative molecular mass (which ever is suitable) expressed in grammes.

One mole of carbon has a mass of 12gOne mole of oxygen molecules has a mass of 32gOne mole of CuSO4 has a mass of 159.5g

No. of moles = m/M = mass in grammes/ mass of one mole

3. How many moles are in 16.6g of CO2Answer: 0.377 moles

4. How many moles are in 10g of NaOH?

Answer: 0.25moles5. What is the mass of 0.05 moles of CuSO4Answer: 0.05 x 159.5 = 7.9moles

6. If 0.2 mole of a compound has a mass of 9.2g, what is the relative molecular mass of the compound?

Answer 9.2 /0.2 = 46g

No. of Particles

One mole of any substance contains 6 x 10 23

7. How many (a) molecules (b) atoms are there in 42.4g of Na2CO3Chapter 11: Stoichiometry

– the mole forms the basis of most calculations. Many of the calculations involve converting one quantity to another so is important to practice this.I mole = molar massI mole = 22.4 litres at S.T.P.I mole = 6 x 1023

1. What is meant by empirical formula?2. Calculate the empirical formula of the following compound: sulphur 50%, oxygen

50%.3. Which of the following compounds has the highest % of carbon: carbon dioxide,

calcium carbonate, ethanol?4. Oxygen is prepared by decomposing hydrogen peroxide according to the

equation: 2H2O2 → 2H2O + O2 Calculate the volume of oxygen formed at s.t.p. by the decomposition of a solution containing 41g of hydrogen peroxide.

Exam Questions on Atomic Theory

You might try to get through 2014, 2013, 2012 2011 2010, 2009 and 2008.

1) 2014 Q.4 (a), (b), (c), (d), (e), (f), (g) Q.5 Q.10 (b), Q.11 (a)2) 2013 Q.4 (a), (b), (c), (d), (e), (g), Q.5, Q.10 (a) (c), Q.11 (a)3) 2012 Q.4 (a), (b), (c), (d), (f) Q.5 Q.11 (a)4) 2011 Q.4 (a), (b), (c), (d), (e) Q.5 Q.10 (c) Q.11 (b)5) 2010 Q.4 (a),(b),(c),(d),(e),(g) Q.5, Q.10 (a) Q.11 (b)6) 2009 Q.3,Q.4 (a),(b),(c),(d) Q.5,Q.10 (c),Q.11(b)7) 2008 Q.4 (a),(b),(c),(d), (g) Q.5, Q.10 (c)8) 2007 Q.4 (a),(c),(k)-B Q.5, Q.11 (a)9) 2006 Q.4 (a),(b),(c), (f) Q.5, Q.10 (a)10) 2005 Q.4 (a),(b),(c),(d),(e),(h) Q.5 Q.10 (b) Q.11 (b)11) 2004 Q.4 (a),(b),(c),(h) Q.5 Q.10(b) Q.11(a)12) 2003 Q.4 (a),(b),(d),(f),(g),(i) Q.5 Q.11(b)

13) 2002 Q.4 (a),(b),(c),(h) Q.5 Q.10 (b) Q.11 (b)