Chemistry Unit 1 for CAPE® PDF

Chemistry Unit 1 for CAPE® Chemistry Unit 1 for CAPE® Roger Norris Leroy Barrett Annette Maynard-Alleyne Jennifer Murray 3 Great Clarendon Street, Oxford, OX2 6DP, United Kingdom Oxford Universi...

Chemistry Unit 1 for CAPE® Chemistry Unit 1 for CAPE® Roger Norris Leroy Barrett Annette Maynard-Alleyne Jennifer Murray 3 Great Clarendon Street, Oxford, OX2 6DP, United Kingdom Oxford University Press is a department of the University of Oxford. It furthers the University’s objective of excellence in research, scholarship, and education by publishing worldwide. Oxford is a registered trade mark of Oxford University Press in the UK and in certain other countries Text © Roger Norris, Leroy Barrett, Annette Maynard-Alleyne, Jennifer Murray 2012 Original illustrations © Oxford University Press 2015 The moral rights of the authors have been asserted First published by Nelson Thornes Ltd in 2012 This edition published by Oxford University Press in 2015 All rights reserved. No part of this publication may be reproduced, stored in a retrieval system, or transmitted, in any form or by any means, without the prior permission in writing of Oxford University Press, or as expressly permitted by law, by licence or under terms agreed with the appropriate reprographics rights organization. Enquiries concerning reproduction outside the scope of the above should be sent to the Rights Department, Oxford University Press, at the address above. You must not circulate this work in any other form and you must impose this same condition on any acquirer British Library Cataloguing in Publication Data Data available 978-1-4085-1668-3 10 9 8 7 6 5 Printed in Great Britain by CPI Group (UK) Ltd., Croydon CR04YY Acknowledgements Cover: Mark Lyndersay, Lyndersay Digital, Trinidad www.lyndersaydigital.com Illustrations: Wearset Ltd, Boldon, Tyne and Wear Page make-up: Wearset Ltd Index: Indexing Specialists (UK) Ltd Photographs Page 2 iStockphoto Although we have made every effort to trace and contact all copyright holders before publication this has not been possible in all cases. If notied, the publisher will rectify any errors or omissions at the earliest opportunity. Links to third party websites are provided by Oxford in good faith and for information only. Oxford disclaims any responsibility for the materials contained in any third party website referenced in this work. Contents Introduction v Chapter 4 Redox reactions 48 4.1 Redox reactions and oxidation Module 1 state 48 Chapter 1 Atomic structure 2 4.2 Redox equations 50 1.1 The structure of the atom 2 Chapter 5 Kinetic theory 52 1.2 Isotopes and radioactivity 4 1.3 Energy levels and emission 5.1 The gas laws 52 spectra 6 5.2 The kinetic theory of gases 54 1.4 Sub-shells and atomic orbitals 8 5.3 Using the ideal gas equation 56 1.5 Ionisation energies 10 5.4 Changing state 58 Revision questions 12 Chapter 6 Energetics 60 Chapter 2 Structure and bonding 14 6.1 Enthalpy changes 60 2.1 States of matter and forces of 6.2 Bond energies 62 attraction 14 6.3 Enthalpy changes by experiment 64 2.2 Covalent bonding – dot and cross 6.4 Hess’s law 66 diagrams 16 6.5 More enthalpy changes 68 2.3 More dot and cross diagrams 18 6.6 Lattice energy: the Born–Haber 2.4 Ionic and metallic bonding 20 cycle 70 2.5 Electronegativity and 6.7 More about lattice energy 72 intermolecular forces 22 Exam-style questions – Module 1 74 2.6 Properties of giant structures 24 2.7 Properties of simple molecular Module 2 compounds 26 2.8 Shapes of molecules (1) 28 Chapter 7 Rates of reaction 76 2.9 Shapes of molecules (2) 30 7.1 Following the course of a reaction 76 Revision questions 32 7.2 Calculating rate of reaction 78 7.3 Collision theory and rate of Chapter 3 The mole concept 34 reaction 80 3.1 Equations and moles 34 7.4 Rate equations 82 3.2 Calculations using the mole 7.5 Order of reaction from concept 36 experimental data 84 3.3 Empirical and molecular formulae 38 7.6 Reaction mechanisms 86 3.4 Using Avogadro’s law 40 Revision questions 88 3.5 Solution concentration and titrations 42 Chapter 8 Chemical equilibria 90 3.6 More about titrations 44 8.1 Characteristics of chemical Revision questions 46 equilibria 90 8.2 The equilibrium constant, K 92 c iii  Contents 8.3 Equilibrium calculations involving K 94 c Chapter 12 The chemistry of 8.4 The equilibrium constant, K 96 Groups II, IV and VII 140 p 8.5 Le Chatelier’s principle 98 12.1 Properties of Group II elements 140 8.6 Solubility product, K 100 sp 12.2 Group II carbonates, nitrates and 8.7 Solubility product and the sulphates 142 common ion effect 102 12.3 Group IV elements and their tetrachlorides 144 Chapter 9 Acid–base equilibria 104 12.4 Some properties of Group IV 9.1 Brønsted–Lowry theory of acids oxides 146 and bases 104 12.5 Relative stability of Group IV 9.2 Simple pH calculations 106 cations and oxides 148 9.3 Equilibrium calculations involving 12.6 Group VII: the halogens 150 weak acids 108 12.7 Halogens and hydrogen halides 152 9.4 Acid and base dissociation 12.8 Reactions of halide ions 154 constants 110 Revision questions 156 9.5 Changes in pH during acid–base titrations 112 Chapter 13 Transition elements 158 9.6 Acid–base indicators 114 13.1 An introduction to transition 9.7 Buffer solutions 116 elements 158 9.8 Calculations involving buffer 13.2 Physical properties of transition solutions 118 elements 160 Chapter 10 Electrode potentials 13.3 Coloured ions and oxidation and equilibrium 120 states 162 13.4 Complex ions and ligand 10.1 Introducing standard electrode exchange 164 potential 120 13.5 Complex ions and redox reactions 166 10.2 Electrode potentials and cell potentials 122 Chapter 14 Qualitative tests for 10.3 Redox reactions and cell diagrams 124 ions 168 Ø 10.4 Using E values to predict chemical change 126 14.1 Identifying cations (1) 168 10.5 Electrode potentials and 14.2 Identifying cations (2) 170 electrochemical cells 128 14.3 Identifying anions 172 Revision questions 130 Exam-style questions – Module 3 174 Exam-style questions – Module 2 132 Data sheets 176 Module 3 Answers to revision questions 178 Chapter 11 The elements of Glossary 183 Period 3 134 Index 186 11.1 Physical properties of the Period 3 elements 134 Answers to exam-style 11.2 Patterns in Period 3 elements 136 questions CD 11.3 Properties of Period 3 oxides and chlorides 138 iv Introduction This Study Guide has been developed exclusively with the Caribbean ® Examinations Council (CXC ) to be used as an additional resource by candidates, both in and out of school, following the Caribbean Advanced ® Prociency Examination (CAPE ) programme. ® It has been prepared by a team with expertise in the CAPE syllabus, teaching and examination. The contents are designed to support learning ® by providing tools to help you achieve your best in CAPE Chemistry and the features included make it easier for you to master the key concepts and requirements of the syllabus. Do remember to refer to your syllabus for full guidance on the course requirements and examination format! Inside this Study Guide is an interactive CD, which includes electronic activities to assist you in developing good examination techniques:  On Your Marks activities provide sample examination-style short answer and essay type questions, with example candidate answers and feedback from an examiner to show where answers could be improved. These activities will build your understanding, skill level and condence in answering examination questions.  Test Yourself activities are specically designed to provide experience of multiple-choice examination questions and helpful feedback will refer you to sections inside the study guide so that you can revise problem areas. This unique combination of focused syllabus content and interactive examination practice will provide you with invaluable support to help you ® reach your full potential in CAPE Chemistry. v 1 Atomic structure 1.1 The structure of the atom Learning outcomes Dalton’s atomic theory John Dalton thought of atoms as being hard spheres. He suggested that: On completion of this section, you should be able to:  all atoms of the same element are exactly alike  describe Dalton’s atomic theory  atoms cannot be broken down any further  atoms of different elements have different masses  describe the structure of the atom in terms of position,  atoms combine to form more complex structures ( compounds). relative charges and masses of the sub-atomic particles. Later changes to Dalton’s atomic theory Dalton’s atomic theory was first described in 1807. Over the past two centuries, scientists have modified the theory to fit the evidence available. 1897 Thomson discovered the electron (which he called a ‘corpuscle’). He suggested the ‘plum-pudding’ model of the atom: electrons embedded in a sea of positive charge. 1909 Rutherford suggested that most of the mass of the atom was in a tiny positively-charged nucleus in the middle of the atom. This led to the planetary model with electrons surrounding the nucleus. 1913 Bohr suggested that electrons could only orbit the nucleus at certain distances depending on their energy. John Dalton (1766–1844) 1932 Did you know? Chadwick discovered the presence of the neutron in the nucleus. Scientists now think that matter is 1926 onwards made up of sub-atomic particles Schrödinger, Born and Heisenberg described the position of an electron in called ‘leptons’ and ‘quarks’. An terms of the probability of finding it at any one point. This led to the idea electron is a type of lepton and of atomic orbitals (Section 1.4). protons and neutrons are made up of quarks. cloud of electron shells electrons (–) electrons electron (–) electron proton sea of + charge nucleus (+) nucleus Figure 1.1.1 Three different models of the atom: a The plum-pudding model; b The Bohr model; c The atomic orbital model neutron nucleus Figure 1.1.2 shows a useful model of an atom based on the Bohr model. Figure 1.1.2 The nucleus of an atom contains protons and neutrons. The electrons are outside the nucleus in fixed energy levels. 2  Chapter 1 Atomic structure Masses and charges of sub-atomic particles The actual masses and charges of protons, neutrons and electrons are –27 very small. For example, a proton has a mass of 1.67 × 10 kg and a –19 charge of +1.6 × 10 C. The relative masses and charges give an easier comparison. These are shown in the table. Proton Neutron Electron Relative mass 1 1 1 _____ 1836 Relative charge +1 0 –1 Sub-atomic particles and electric fields The effect of a strong electric field on beams of protons, neutrons and electrons is shown below. Exam tips + + + beam of beam of beam of The direction of the movement of electrons protons neutrons an electron in a magnetic eld can be found by Fleming’s left-hand deflection towards deflection towards no effect rule as shown below. Remember (1) + plate – plate The direction of the magnetic eld Figure 1.1.3 The effect of an electric field on beams of electrons, protons and neutrons is N → S (2) that the centre nger Charged particles move towards the plates with opposite charge and away points in the direction opposite to from the plates with the same charge. If we use the same voltage to the electron flow. deflect electrons and protons, the electrons are deflected to a much first greater extent. This is because electrons have a very small mass finger compared with protons. field left hand centre finger Sub-atomic particles and magnetic fields current thumb Figure 1.1.4 shows an evacuated glass tube. When heated by a low voltage motion supply, electrons are produced from the tungsten filament, F. They are accelerated towards the metal cylinder, C. They produce a glow when they collide with the fluorescent screen. When the north pole of a magnet Key points is brought near the electron beam, the glow on the fluorescent screen moves downwards. The direction of the movement can be predicted from  The atomic theory has developed Fleming’s left-hand rule. The direction of movement is at right angles to to t the experimental evidence. the direction of the magnetic field and the conventional current. Remember that the direction of electron flow is opposite that of the  Most of the mass of the atom conventional current. is in the nucleus which contains protons and neutrons. The electron flow electrons are outside the nucleus F C in shells. conventional magnetic field  Electrons, protons and neutrons current have characteristic relative N fluorescent masses and charges. screen movement  Beams of protons and electrons are deflected in opposite Figure 1.1.4 The effect of a magnetic field on a beam of electrons directions by electric and Using different apparatus, beams of protons will be deflected in the magnetic elds. Neutrons are not opposite direction to the electrons. Beams of neutrons will not be deflected. deflected since they have no charge. 3 1.2 Isotopes and radioactivity Learning outcomes Atomic number and mass number Atomic number (Z) On completion of this section, you should be able to: Atomic number ( Z) is the number of protons in the nucleus of an atom. (It also equals the number of electrons in a neutral atom and the position  dene ‘mass number’ and of the element in the Periodic Table.) ‘isotopes’  dene ‘relative atomic mass’ and Mass number (A) ‘relative isotopic mass’ using the Mass number ( A) is the number of protons plus neutrons in the nucleus 12 C scale of an atom.  know how to calculate relative So the number of neutrons in an atom is A − Z (mass number − atomic atomic masses number).  write equations involving α-, β- and γ-radiation Isotopes  state some uses of radioisotopes. Isotopes are atoms with the same atomic number but different mass numbers. We show the mass number and atomic number of an isotope like this: electron mass number → 23 Na ← element symbol 11 atomic number → proton Relative atomic and relative isotopic mass 12 hydrogen atoms The mass of a single atom is too small to weigh, so we use the C 1 proton (carbon-12) atom as a standard. We compare the masses of other atoms neutron to this standard, which has a mass of exactly 12 units. Relative atomic mass (A ) r Relative atomic mass (A ) is the weighted average mass of naturally r 12 occurring atoms of an element on a scale where an atom of C has a mass of exactly 12 units. deuterium atoms 1 proton, 1 neutron Relative isotopic mass Relative isotopic mass is the mass of a particular isotope of an element 12 on a scale where an atom of C has a mass of exactly 12 units. Calculating accurate relative atomic masses Most naturally occurring elements exist as a mixture of isotopes. So their tritium atoms mass numbers are not always exact. We have to take into account the 1 proton, 2 neutrons proportion of each isotope. This is called its relative isotopic abundance. Figure 1.2.1 The three isotopes of The relative atomic mass for neon is calculated below from the relative hydrogen isotopic masses and relative abundance. Did you know? Isotope Relative isotopic mass Relative abundance/% Isotopes with 2, 8, 20, 28, 50, 82 20 Ne 20 90.9 and 126 neutrons or protons are extremely stable. These ‘magic 21 Ne 21 0.30 numbers’ are thought to indicate 22 closed nuclear shells comparable to Ne 22 8.8 the full shells of the noble gases. (20.0 × 90.9) + (21.0 × 0.3) + (22 × 8.8) ______________________________________ A = = 20.2 r 100 4  Chapter 1 Atomic structure Radioactivity Isotopes of some elements have nuclei which break down (decay) spontaneously. These are described as radioactive isotopes. As the nuclei break down, rays or particles are given out. These are called emissions. The table shows three types of emission. Name of Type of particles/rays Stopped by emission emitted Exam tips Alpha (α) helium nuclei (positively thin sheet of paper charged particles) A common error is to suggest that α-particles are helium atoms. They Beta (β) electrons (produced by nuclear 6 mm thick are not: they are helium nuclei 2+ changes) aluminium foil He. Another error is to think that β-particles are electrons from the Gamma (γ) very high frequency thick lead sheet outside the nucleus. The β-particle electromagnetic radiation electrons arise from the nucleus not from the electron shells. Equations can be written for each of these types of decay. For example: α-decay The isotope produced has a mass number of 4 units lower and a nuclear charge of 2 units lower than the original atom: 223 219 4 Ra → Rn + He 88 86 2 β-decay The mass number stays the same but the number of protons increases by 1 1 one. This is because a neutron ( n) is changed into a proton ( p) and an 0 1 0 electron ( e): –1 1 1 0 ( n) → ( p) + ( e) 0 1 –1 For example carbon-14 decays to nitrogen-14 when it emits a β-particle: 14 14 0 C → N + e 6 7 –1 Key points γ-decay γ-rays can be emitted along with α- or β-particles or in a process called  Mass number is the total number ‘electron capture’. A proton is converted to a neutron, so the mass of protons + neutrons in an number stays the same but the atomic number decreases by one. For isotope. example:  Isotopes are atoms with the 37 0 37 Ar + e → Cl 18 –1 17 same number of protons but different numbers of neutrons. Uses of radioisotopes  Relative atomic and isotopic  Tracers for searching for faults in pipelines and for studying the masses are compared using the 131 12 working of certain organs in the body, e.g. I is used to study thyroid C scale. function.  The relative atomic mass of an  In medicine, for radiotherapy in the treatment of cancers. element is the weighted mean of 14  Dating the age of objects, e.g. using C to date objects which were the atoms of the isotopes they once living. contain. 241  Smoke detectors often use Am.  When an unstable isotope decays 235  Generating power, e.g. U is used in many nuclear reactors as a rays or particles are emitted. source of energy.  α-, β-, and γ-radiation have characteristic properties. 5 1.3 Energy levels and emission spectra Learning outcomes Energy quanta and energy levels The electrons in atoms have certain fixed values of energy. The smallest On completion of this section, you fixed amount of energy required for a change is called a quantum of should be able to: energy. We say that the energy is discrete or quantised. The electrons are  explain how data from emission arranged in energy levels which are quantised. The atom is most stable spectra give evidence of discrete when the electrons are in the lowest energy levels possible for each of energy levels for electrons them. The electrons are then in their ground state. When an electron absorbs a quantum of radiation it can move up to a higher energy level.  explain the lines of the Lyman The electron is then said to be in an excited state. When an excited and Balmer series in emission electron falls to a lower energy level a quantum of radiation is given out. spectra The electron moves to a lower energy level.  understand how the difference in energy between two electrons’ energy levels is related to the energy frequency of radiation. electrons in electrons in ground state excited state Figure 1.3.1 An electron can jump from one energy level to another by absorbing or emitting a fixed amount (quantum) of energy The difference in energy between two energy levels is related to the frequency of radiation by the relationship: –1 energy (J) → ΔE = hν ← frequency of radiation (s ) ↑ –34 –1 Planck’s constant (6.63 × 10 J s ) In any atom there are several possible energy levels. These are called principal quantum levels or principal energy levels. They are also called electron shells. They are given numbers n = 1, n = 2, n = 3, etc. going further away from the nucleus. Atomic emission spectra When electrical or thermal energy is passed through a gaseous sample of an element, the radiation is emitted only at certain wavelengths or frequencies. An emission spectr um differs from a normal visible light spectrum in that:  It is made up of separate lines.  The lines converge (get closer to each other) as their frequency increases. lines get closer together Figure 1.3.2 Part of a simplified line emission spectrum of hydrogen 6  Chapter 1 Atomic structure Interpreting the hydrogen emission spectrum Each line in the hydrogen emission spectrum is a result of electrons moving from a higher to a lower energy level. Among the several series of lines seen are the:  Lyman series (seen in the ultraviolet region), where previously excited electrons fall back to the n = 1 energy level.  Balmer series (seen in the visible region), where previously excited electrons fall back to the n = 2 energy level. n = 5 n = 4 n = 3 n = 2 n = 1 frequency emission spectrum Lyman series Balmer series Figure 1.3.3 A line emission spectrum is the result of electrons falling from higher energy levels to lower energy levels The energy associated with each line is found by using the relationship Did you know? ΔE = hν. You can see that the greater the frequency (the smaller the wavelength), the more energy is released. The point where the lines There are other sets of emission eventually come together is called the convergence limit. This represents spectrum lines in the infrared an electron falling from the highest possible energy level. If the electron region. These are called the Paschen, has more energy than this, it becomes free from the pull of the nucleus of Brackett and Pfund series. the atom. Then the atom is converted to an ion. The Bohr model of an atom n = 5 n = 4 The energy levels get closer together towards the outside of the atom. n = 3 This is mirrored by the spectral lines getting closer together. In the Bohr n = 2 model of the hydrogen atom shown opposite, a quantum of energy moves n = 1 the hydrogen electron in the n = 1 level (ground state) to the n = 2 level. A further quantum of energy will excite the electron to the n = 3 level. The more energy an electron has, the higher the energy level it will be in. When the electron loses energy, it will fall down to the lower energy levels, emitting radiation of characteristic frequencies. Key points electron shells (principle quantum shells, n)  Electrons occupy specic energy levels (quantum levels) in the atom. Figure 1.3.4 The Bohr model of the atom.  When electrons gain specic quanta of energy they move from lower to The principal energy levels get closer higher energy levels. They become ‘excited’. together towards the outside of the atom.  When excited, electrons lose energy, they fall back to lower energy levels emitting radiation of characteristic frequency. This is the origin of the line emission spectrum.  The energy difference between any two energy levels is given by ΔE = hν 7 1.4 Sub-shells and atomic orbitals Learning outcomes Shells and sub-shells The principal quantum shells apart from the first, are split into On completion of this section, you sub-shells (sub-levels). These are named s, p, d. Figure 1.4.1 shows the should be able to: sub-shells for the first four quantum shells in terms of increasing energy.  know that the principal quantum Notice that after the element calcium the order of the sub-shells, in shell for n = 2 and above can be terms of energy, does not always conform to the pattern of s followed by p split into sub-shells (subsidiary followed by d. For example in scandium the 4s sub-shell is at a lower quantum shells) energy level than the 3d.  know the relative energies of the For a neutral atom, the number of electrons is equal to the number of shells and sub-shells protons. The maximum number of electrons each shell and sub-shell can hold is shown in the table.  know how to determine the electronic conguration of atoms Principal Total number of Maximum number of electrons in and ions quantum electrons in principal sub-shell  understand the term ‘atomic shell quantum shell orbital’ s p d  describe the shapes of s and p orbitals. n = 1 2 2 – – n = 2 8 2 6 – 4d n = 3 18 2 6 10 5s n = 4 4p ygrene 3d 4s Atomic orbitals n = 3 3p gnisaercni Each sub-shell contains one or more atomic orbitals. An orbital is a 3s region of space where there is a high probability of finding an electron. Each orbital can hold a maximum of two electrons. So the number of 2p n = 2 orbitals in each sub-shell is: s = 1, p = 3 and d = 5. Orbitals differ from 2s each other in shape (see Figure 1.4.2). n = 1 1s z z z z prinicipal quantum sub-shell y y y y shell x x x x Figure 1.4.1 The principal quantum shells, apart from n = 1, are split into sub-shells s orbital p orbital p orbital p orbital x y z Figure 1.4.2 The shapes of s and p orbitals Note that:  A 2s orbital is larger than the 1s orbital and a 3p is larger than the 2p.  The three different p orbitals are mutually at right angles to each other. Here are two examples of how to write the type of orbital and the number of electrons present in them: 2 electrons in orbital 5 electrons ↓ ↓ 2 5 First quantum shell → 1s Third quantum shell → 3p s sub-shell ↑ p sub-shell ↑ 8  Chapter 1 Atomic structure The electronic configuration of atoms Exam tips The electronic configuration shows the number and type of the This diagram may help you electrons in a particular atom. The procedure is: remember the order of lling the 1 Make sure you know the order of the orbitals, 1s 2s 2p 3s etc., as well sub-shells. as the number of electrons in the atom. 1s 2 Fill up the sub-shells to the maximum amount starting with those of lowest energy, e.g. an s sub-shell can have a maximum of 2 electrons 2s 2p and a p sub-shell a maximum of 6. 3s 3p 3d 3 Stop when you have the correct number of electrons. 4s 4p 4d Example 1 5s 5p Electronic configuration of sodium (Z = 11): 6s 2 2 6 1 1s 2s 2p 3s Example 2 Did you know? Electronic configuration of scandium (Z = 21): 2 2 6 2 6 2 1 2 2 6 2 6 1 2 1s 2s 2p 3s 3p 4s 3d (or 1s 2s 2p 3s 3p 3d 4s ) The electronic conguration can also be written to show the separate Note that the 4s sub-shell is filled before the 3d (see Figure 1.4.1). orbitals p , p , p. (see Figure 1.4.2). x y z For example, the electronic The electronic configuration of ions conguration of sodium (Example 1) can be written: When ions are formed one or more electrons are removed from or added 2 2 2 2 2 1 to the atom. In general, the electrons are lost or gained from the outer 1s 2s 2p 2p p 3s x y z sub-shell. Example 1 2 2 6 2 2+ 2 2 6 Magnesium atom 1s 2s 2p 3s ; Magnesium ion (Mg ) 1s 2s 2p Example 2 2 2 3 3– 2 2 6 Nitrogen atom 1s 2s 2p ; Nitride ion (N ) 1s 2s 2p Note that the electronic configuration of some transition element atoms and ions are less straightforward (see Section 13.1). Key points  The principal quantum shells (apart from the rst) are divided into sub- shells named s, p, d, etc.  Each sub-shell can hold a maximum number of electrons.  An orbital is a region of space where there is a good chance of nding an electron.  s orbitals are spherical in shape and p have an hour-glass shape.  Each orbital can hold a maximum of 2 electrons.  The maximum number of electrons in each sub-shell is s = 2, p = 6, d = 10.  Electronic congurations of atoms and ions are found by adding electrons to the energy sub-levels placed in order of increasing energy. 9 1.5 Ionisation energies Learning outcomes Ionisation energy The first ionisation energy , ΔH , is the energy needed to remove one i1 On completion of this section, you electron from each atom in one mole of atoms of an element in its should be able to: gaseous state to form one mole of gaseous ions. For example:  know the factors influencing + –1 Cl(g) → Cl (g) + e ΔH = +1260 kJ mol i1 ionisation energy The value of the ionisation energy depends on:  explain how ionisation energy data give evidence for sub-shells  Size of nuclear charge: In any one period, the higher the nuclear  know how electron conguration charge (greater number of protons), the greater the attractive force between the nucleus and outer electrons and the more energy needed can be determined from to remove these electrons. So ΔH tends to increase with increase in successive ionisation energies. i1 nuclear charge.  Distance of the outer electrons from the nucleus: The further the outer electrons are from the nucleus, the less attractive force there is between them and the nucleus and the lower the value of ΔH i1  Shielding: Electrons in full inner shells reduce the attractive force between the nucleus and the