Cambridge international AS and a level chemistry coursebook by roger norris 1

Roger Norris, Lawrie Ryan and David Acaster Cambridge International AS and A Level Chemistry Coursebook c a mb r id g e u n i ve r s i t y p re s s Cambridge, New York, Melbourne, Madrid, Cape Town, Singapore, São Paulo, Delhi, Mexico City Cambridge University Press The Edinburgh Building, Cambridge CB2 8RU, UK www.cambridge.org Information on this title: www.cambridge.org/9780521126618 © Cambridge University Press 2011 This publication is in copyright Subject to statutory exception and to the provisions of relevant collective licensing agreements, no reproduction of any part may take place without the written permission of Cambridge University Press First published 2011 5th printing 2012 Printed in Dubai by Oriental Press A catalogue record for this publication is available from the British Library ISBN 978-0-521-12661-8 Paperback with CD-ROM for Windows and Mac Cambridge University Press has no responsibility for the persistence or accuracy of URLs for external or third-party internet websites referred to in this publication, and does not guarantee that any content on such websites is, or will remain, accurate or appropriate noti c e to t e ach e r s The photocopy masters in this publication may be photocopied or distributed electronically free of charge for classroom use within the school or institute which purchases the publication Worksheets and copies of them remain in the copyright of Cambridge University Press and such copies may not be distributed or used in any way outside the purchasing institution Contents Introduction Moles and equations 1.1 1.2 1.3 1.4 1.5 1.6 1.7 1.8 Introduction Masses of atoms and molecules Accurate relative atomic masses Amount of substance Mole calculations Chemical formulae and chemical equations Solutions and concentration Calculations involving gas volumes Test yourself questions Atomic structure 2.1 2.2 2.3 2.4 Elements and atoms Inside the atom Numbers of nucleons How many protons, neutrons and electrons? Test yourself questions Electrons in atoms 3.1 3.2 3.3 3.4 3.5 Simple electronic structure Evidence for electronic structure Sub-shells and atomic orbitals Electronic configurations Patterns in ionisation energies in the Periodic Table Test yourself questions Chemical bonding 4.1 4.2 4.3 4.4 4.5 4.6 4.7 Introduction: types of chemical bonding Ionic bonding Covalent bonding Shapes of molecules Metallic bonding Intermolecular forces Bonding and physical properties Test yourself questions States of matter 5.1 5.2 5.3 5.4 5.5 5.6 States of matter The gaseous state The liquid state The solid state Ceramics Conserving materials Test yourself questions Enthalpy changes 6.1 6.2 6.3 6.4 6.5 6.6 Introduction: energy changes What are enthalpy changes? Standard enthalpy changes Measuring enthalpy changes Hess’s law Bond energies and enthalpy changes Test yourself questions Redox reactions and electrolysis 7.1 7.2 7.3 7.4 What is a redox reaction? Redox and electron transfer Oxidation numbers Electrolysis Test yourself questions Equilibrium 8.1 Reversible reactions and equilibrium 8.2 Changing the position of equilibrium 8.3 Equilibrium expressions and the equilibrium constant, Kc 8.4 Equilibria in gas reactions: the equilibrium constant, Kp 8.5 Equilibria and the chemical industry 8.6 Acid–base equilibria Test yourself questions Rates of reaction 9.1 Introduction to reaction kinetics 9.2 The effect of concentration on rate of reaction 9.3 The effect of temperature on rate of reaction 9.4 Catalysis Test yourself questions 10 Periodicity 10.1 Introduction – structure of the Periodic Table 10.2 Periodicity of physical properties 10.3 Periodicity of chemical properties 10.4 Oxides of Period elements 10.5 Chlorides of Period elements Test yourself questions Contents iii 11 Groups II and VII 11.1 Physical properties of Group II elements 11.2 Reactions of Group II elements 11.3 Thermal decomposition of Group II carbonates and nitrates 11.4 Some uses of Group II compounds 11.5 Physical properties of Group VII elements 11.6 Reactions of Group VII elements 11.7 Reactions of the halide ions 11.8 Disproportionation 11.9 Uses of the halogens and their compounds Test yourself questions 12 Nitrogen and sulfur 12.1 12.2 12.3 12.4 Nitrogen gas Ammonia and ammonium compounds Sulfur and its oxides Sulfuric acid Test yourself questions 13 Introduction to organic chemistry 13.1 13.2 13.3 13.4 13.5 13.6 13.7 13.8 Introduction Representing organic molecules Functional groups Naming organic compounds Bonding in organic molecules Structural isomerism Stereoisomerism Organic reactions – mechanisms Test yourself questions and answers 13.9 Types of organic reactions Test yourself questions 14 Hydrocarbons 14.1 14.2 14.3 14.4 14.5 Introduction – the alkanes Sources of the alkanes Reactions of alkanes The alkenes Addition reactions of the alkenes Test yourself questions 15 Halogenoalkanes 15.1 Introduction 15.2 Nucleophilic substitution reactions 15.3 Mechanism of nucleophilic substitution in halogenoalkanes 15.4 Elimination reactions 15.5 Uses of halogenoalkanes Test yourself questions iv Contents 16 Alcohols and esters 16.1 Introduction – the alcohols 16.2 Reactions of the alcohols Test yourself questions 17 Carbonyl compounds 17.1 17.2 17.3 17.4 17.5 Introduction – aldehydes and ketones Preparation of aldehydes and ketones Reduction of aldehydes and ketones Nucleophilic addition with HCN Testing for aldehydes and ketones Test yourself questions 18 Lattice energy 18.1 Introducing lattice energy 18.2 Enthalpy change of atomisation and electron affinity 18.3 Born–Haber cycles 18.4 Factors affecting the value of lattice energy 18.5 Ion polarisation 18.6 Enthalpy changes in solution 19 Electrode potentials 19.1 Redox reactions revisited 19.2 Electrode potentials 19.3 Measuring standard electrode potentials ntials tials 19.4 Using E values 19.5 Cells and batteries 19.6 More about electrolysis 19.7 Quantitative electrolysis 20 Ionic equilibria 20.1 Introduction 20.2 pH calculations 20.3 Weak acids – using the acid dissociation constant, Ka 20.4 Indicators and acid–base titrations 20.5 Buffer solutions 20.6 Equilibrium and solubility 21 Reaction kinetics 21.1 Introduction 21.2 Rate of reaction 21.3 Rate equations 21.4 21.5 21.6 21.7 21.8 Which order of reaction? Calculations involving the rate constant, k Deducing order of reaction from raw data Kinetics and reaction mechanisms Catalysis 22 Group IV 22.1 22.2 22.3 22.4 22.5 Introduction Variation in properties The tetrachlorides The oxides Relative stability of the +2 and +4 oxidation states 22.6 Ceramics from silicon(IV) oxide 28 The chemistry of life 28.1 Introduction 28.2 Reintroducing amino acids and proteins 28.3 The structure of proteins 28.4 Enzymes 28.5 Factors affecting enzyme activity 28.6 Nucleic acids 28.7 Protein synthesis 28.8 Genetic mutations 28.9 Energy transfers in biochemical reactions 28.10 Metals in biological systems 29 Applications of analytical chemistry 23 Transition elements 23.1 23.2 23.3 23.4 What is a transition element? Physical properties of the transition elements Redox reactions Ligands and complex formation 29.1 29.2 29.3 29.4 Electrophoresis Nuclear magnetic resonance (NMR) Chromatography Mass spectrometry 30 Design and materials 24 Benzene and its compounds 24.1 24.2 24.3 24.4 Introduction to benzene Reactions of arenes Phenol Reactions of phenol 25 Carboxylic acids and acyl compounds 25.1 The acidity of carboxylic acids 25.2 Acyl chlorides 25.3 Reactions to form tri-iodomethane 26 Organic nitrogen compounds 26.1 26.2 26.3 26.4 Amines Amides Amino acids Peptides and proteins 27 Polymerisation 27.1 27.2 27.3 27.4 Types of polymerisation Polyamides Polyesters Polymer deductions 30.1 30.2 30.3 30.4 30.5 Designing new medicinal drugs Designing polymers Nanotechnology Fighting pollution ‘Green chemistry’ Appendix 1: The Periodic Table Appendix 2: Standard electrode potentials Answers to check-up questions Answers to end-of-chapter questions Answers to Test yourself questions Advice on the practical exam Revision skills Glossary Index Acknowledgements Contents v Introduction Cambridge CIE AS and A Level Chemistry This new Cambridge AS/A Level Chemistry course has been specifically written to provide a complete and precise coverage for the Cambridge International Examinations syllabus 9701 The language has been kept simple, with bullet points where appropriate, in order to improve the accessibility to all students Principal Examiners have been involved in all aspects of this book to ensure that the content gives the best possible match to both the syllabus and to the type of questions asked in the examination The book is arranged in two sections Chapters 1–17 correspond to the AS section of the course (for examination in Papers 1, and 31/32) Chapters 18–30 correspond to the A level section of the course (for examinations in papers and 5) Within each of these sections the material is arranged in the same sequence as the syllabus For example in the AS section, Chapter 1 deals with atoms, molecules and stoichiometry and Chapter deals with atomic structure The A level section starts with lattice energy (Chapter 18: syllabus section 5) then progresses to redox potentials (Chapter 19: syllabus section 6) Nearly all the written material is new, although some of the diagrams have been based on material from the endorsed Chemistry for OCR books and (Acaster and Ryan, 2008) There are separate chapters about nitrogen and sulfur (Chapter 12) and the elements and compounds of Group IV (Chapter 22), which tie in with the specific syllabus sections Electrolysis appears in Chapter and quantitative electrolysis in Chapter 19 The chapter on reaction kinetics (Chapter 21) includes material about catalysis whilst the organic chemistry section has been rewritten to accommodate the iodoform reaction and to follow the syllabus more closely The last three chapters have been developed to focus on the applications of chemistry (Paper 4B) These chapters contain a wealth of material and questions which will help you gain confidence to maximise your potential in the examination Important definitions are placed in boxes to highlight key concepts Several features of the book are designed to make learning as effective and interesting as possible • Objectives for the chapter appear at the beginning of each chapter These relate directly to the statements in the syllabus, so you know what you should be able to when you have completed the chapter vi Introduction • Important definitions are placed in boxes to highlight key concepts • Check-up questions appear in boxes after most short sections of text to allow you to test yourself They often address misunderstandings that commonly appear in examination answers The detailed answers can be found at the back of the book • Fact files appear in boxes at various parts of the text These are to stimulate interest or to provide extension material They are not needed for the examination • Worked examples, in a variety of forms, are provided in chapters involving mathematical content • Experimental chemistry is dealt with by showing detailed instructions for key experiments, e.g calculation of relative molecular mass, titrations, thermochemistry and rates of reaction Examples are also given of how to process the results of these experiments • A summary at the end of each chapter provides you with the key points of the chapter as well as key definitions • End-of-chapter questions appear after the summary in each chapter Many of these are new questions and so supplement those to be found on the Cambridge Students’ and Teachers’ websites The answers to these questions, along with exam-style mark schemes, can be found at the back of the book • Examiner tips are given with the answers to the endof-chapter questions in the supplementary materials (see below) • A full glossary of definitions is provided at the back of the book Supplementary materials In this e-book version of the Cambridge International AS and A Level Chemistry Coursebook, the CD-ROM content is included as ‘supplementary materials’ These materials include the following: • test-yourself questions (multiple choice) for Chapters 1–17 These are new questions and will help you with Paper They can be found at the end of their respective chapters • study skills guidance to help you direct your learning so that it is productive, provided at the back of the book • advice on the practical examination to help you achieve the best result, also provided at the back of the book Moles and equations Learning outcomes Candidates should be able to: define the terms relative atomic, isotopic, molecular and 12 formula masses based on the C scale analyse mass spectra in terms of isotopic abundances (no knowledge of the working of the mass spectrometer is required) calculate the relative atomic mass of an element given the relative abundances of its isotopes or its mass spectrum define the term mole in terms of the Avogadro constant define the terms empirical and molecular formulae calculate empirical and molecular formulae using combustion data or composition by mass 1.1 Introduction For thousands of years, people have heated rocks and distilled plant juices to extract materials Over the past two centuries, chemists have learnt more and more about how write and/or construct balanced equations perform calculations, including use of the mole concept involving – reacting masses (from formulae and equations) – volumes of gases (e.g in the burning of hydrocarbons) – volumes and concentrations of solutions perform calculations taking into account the number of significant figures given or asked for in the question deduce stoichiometric relationships from calculations involving reacting masses, volumes of gases and volumes and concentrations of solutions to get materials from rocks, from the air and the sea and from plants They have also found out the right conditions to allow these materials to react together to make new substances, such as dyes, plastics and medicines When we make a new substance it is important to mix the reactants in the correct proportions to ensure that none is wasted In order to this we need to know about the relative masses of atoms and molecules and how these are used in chemical calculations 1.2 Masses of atoms and molecules Relative atomic mass, Ar Atoms of different elements have different masses When we perform chemical calculations, we need to know how heavy one atom is compared with another The mass of a single atom is so small that it is impossible to weigh it directly To overcome this problem, we have to weigh a lot of atoms We then compare this mass with the mass of the same number of ‘standard’ atoms Scientists have chosen to use the isotope carbon-12 as the standard This has been given a mass of exactly 12 units The mass of other atoms is found by comparing their mass with the mass of carbon-12 atoms This is called the relative atomic mass, Ar Figure 1.1 A titration is a method used to find the amount of a particular substance in a solution The relative atomic mass is the weighted average mass of naturally occurring atoms of an element on a scale where an atom of carbon-12 has a mass of exactly 12 units Moles and equations From this it follows that Relative formula mass Ar [element Y ] average mass of one atom of element Y × 12 = mass of one atom of carbon-12 For compounds containing ions we use the term relative formula mass This is calculated in the same way as for relative molecular mass It is also given the same symbol, Mr For example, for magnesium hydroxide: We use the average mass of the atom of a particular element because most elements are mixtures of isotopes For example, the exact Ar of hydrogen is 1.0079 This is very close to and most Periodic Tables give the Ar of hydrogen as 1.0 However, some elements in the Periodic Table have values that are not whole numbers For example, the Ar for chlorine is 35.5 This is because chlorine has two isotopes In a sample of chlorine, chlorine-35 makes up about three-quarters of the chlorine atoms and chlorine-37 makes up about a quarter Relative isotopic mass Isotopes are atoms which have the same number of protons but different numbers of neutrons (see page 28) We represent the nucleon number (the total number of neutrons plus protons in an atom) by a number written at the top left-hand corner of the atom’s symbol, e.g 20 Ne, or by a number written after the atom’s name or symbol, e.g neon-20 or Ne-20 We use the term relative isotopic mass for the mass of a particular isotope of an element on a scale where an atom of carbon-12 has a mass of exactly 12 units For example, the relative isotopic mass of carbon-13 is 13.00 If we know both the natural abundance of every isotope of an element and their isotopic masses, we can calculate the relative atomic mass of the element very accurately To find the necessary data we use an instrument called a mass spectrometer Relative molecular mass, Mr The relative molecular mass of a compound (Mr) is the relative mass of one molecule of the compound on a scale where the carbon-12 isotope has a mass of exactly 12 units We find the relative molecular mass by adding up the relative atomic masses of all the atoms present in the molecule For example, for methane: formula atoms present add Ar values Mr of methane CH4 × C; × H (1 × Ar[C]) + (4 × Ar[H]) = (1 × 12.0) + (4 × 1.0) = 16.0 Moles and equations formula ions present add Ar values Mr of magnesium hydroxide Mg(OH)2 × Mg2+; × (OH−) (1 × Ar[Mg]) + (2 × (Ar[O] + Ar[H])) = (1 × 24.3) + (2 × (16.0 + 1.0)) = 58.3 Check-up Use the Periodic Table on page 497 to calculate the relative formula masses of the following: a calcium chloride, CaCl2 b copper(II) sulfate, CuSO4 c ammonium sulfate, (NH4)2SO4 d magnesium nitrate-6-water, Mg(NO3)2.6H2O Hint: for part d you need to calculate the mass of water separately and then add it to the Mr of Mg(NO3)2 1.3 Accurate relative atomic masses Mass spectrometry A mass spectrometer (Figure 1.2) can be used to measure the mass of each isotope present in an element It also compares how much of each isotope is present – the relative abundance A simplified diagram of a mass spectrometer is shown in Figure 1.3 You will not be expected to know the details of how a mass spectrometer works, but it is useful to understand how the results are obtained The atoms of the element in the vaporised sample are converted into ions The stream of ions is brought to a detector after being deflected (bent) by a strong magnetic field As the magnetic field is increased, the ions of heavier and heavier isotopes are brought to the detector The van der Waals’ forces between individual atoms are very small However, the total van der Waals’ forces between very long non-polar molecules such as poly(ethene) molecules (see page 224) can be much larger That is why poly(ethene) is a solid at room temperature Fact file Two types of poly(ethene) are low-density poly(ethene), LDPE, and high-density poly(ethene), HDPE Both have crystalline and non-crystalline regions in them (Figure 4.36) crystalline region non-crystalline region b The table lists the formulae and boiling points of some alkanes Explain the trend in these boiling points Alkane methane ethane propane butane Structural formula CH4 CH3CH3 CH3CH2CH3 CH3CH2CH2CH3 Boiling point / °C −164 −88 −42 Permanent dipole–dipole forces In some molecules, the dipole is permanent Molecules with a permanent dipole are called polar molecules A fine jet of polar molecules will be attracted towards an electrically charged plastic rod or comb (The rod can be charged by rubbing it with a woollen cloth.) Figure 4.37 shows the result of this experiment Figure 4.36 Crystalline and non-crystalline regions in poly(ethene) HDPE has more crystalline regions where the molecules are closer together than LDPE The total van der Waals’ forces are greater in HDPE, so HDPE is the stronger of the two Check-up 10 a The boiling points of the halogens are: fluorine −188 °C chlorine −35 °C bromine +59 °C iodine +184 °C i Describe the trend in these boiling points going down Group VII ii Explain the trend in these boiling points continued Figure 4.37 The deflection of water by an electrically charged nylon comb Chemical bonding 65 The molecules are always attracted to the charged rod, whether it is positively or negatively charged This is because the molecules have both negatively and positively charged ends The forces between two molecules having permanent dipoles are called permanent dipole–dipole forces The attractive force between the δ+ charge on one molecule and the δ− charge on a neighbouring molecule causes a weak attractive force between the molecules (Figure 4.38) Check-up 11 Bromine, Br2, and iodine monochloride, ICl, have the same number of electrons But the boiling point of iodine monochloride is nearly 40 °C higher than the boiling point of bromine Explain this difference Hydrogen bonding weak permanent dipole – dipole force CH3 δ– O δ+ C CH3 δ– O δ+ C CH3 CH3 Figure 4.38 Dipole–dipole forces in propanone For small molecules with the same number of electrons, permanent dipole–dipole forces are often stronger than van der Waals’ forces For example, propanone (CH3COCH3, Mr = 58) has a higher boiling point than butane (CH3CH2CH2CH3, Mr = 58) (Figure 4.39) This means that more energy is needed to break the intermolecular forces between propanone molecules than between butane molecules CH3 CH2 CH3 CH3 CH2 δ+ C H δ– O CH3 butane, boiling point °C propanone, boiling point 56 °C Figure 4.39 The difference in the boiling points of propanone and butane can be explained by the different types of intermolecular force between the molecules The permanent dipole–dipole forces between propanone molecules are strong enough to make this substance a liquid at room temperature There are only van der Waals’ forces between butane molecules These forces are comparatively weak, so butane is a gas at room temperature 66 Chemical bonding Hydrogen bonding is the strongest type of intermolecular force For hydrogen bonding to occur between two molecules we need: • one molecule having a hydrogen atom covalently bonded to F, O or N (the three most electronegative atoms) • a second molecule having a F, O or N atom with an available lone pair of electrons When a hydrogen atom is covalently bonded to a very electronegative atom, the bond is very highly polarised The δ+ charge on the hydrogen atom is high enough for a bond to be formed with a lone pair of electrons on the F, O or N atom of a neighbouring molecule (Figure 4.40) The force of attraction is about one-tenth of the strength of a normal covalent bond For maximum bond strength, the angle between the covalent bond to the hydrogen atom and the hydrogen bond is usually 180° H H N H H N H Figure 4.40 Hydrogen bonding between two ammonia molecules A hydrogen bond is represented by a line of dots The average number of hydrogen bonds formed per molecule depends on: • the number of hydrogen atoms attached to F, O or N in the molecule • the number of lone pairs present on the F, O or N Water has two hydrogen atoms and two lone pairs per molecule (Figure 4.41) So water is extensively hydrogen bonded with other water molecules It has an average of two hydrogen bonds per molecule H δ+ H H δ+ H O δ– H O δ– δ+ H O δ– Figure 4.41 Water can form, on average, two hydrogen bonds per molecule Ammonia is less extensively hydrogen bonded than water (see Figure 4.40) It can form, on average, only one hydrogen bond per molecule Although each ammonia molecule has three hydrogen atoms attached to the nitrogen atom, it has only one lone pair of electrons which can be involved in hydrogen bond formation Check-up 12 Draw diagrams to show hydrogen bonding between the following molecules: a ethanol, C2H5OH, and water b ammonia and water c two hydrogen fluoride molecules How does hydrogen bonding affect boiling point? Some compounds may have higher boiling points than expected This can be due to hydrogen bonding Figure 4.42 shows a graph of the boiling points of the hydrogen halides, HF, HCl, HBr and HI, plotted against the position of the halogen in the Periodic Table Boiling point / °C +50 –50 –100 HF HCl HBr Hl Figure 4.42 The boiling points of the hydrogen halides The rise in boiling point from HCl to HI is due to the increasing number of electrons in the halogen atoms as we go down the group This leads to increased van der Waals’ forces as the molecules get bigger If hydrogen fluoride only had van der Waals’ forces between its molecules, we would expect its boiling point to be about −90 °C However, the boiling point of hydrogen fluoride is 20 °C, which is much higher This is because of the stronger intermolecular forces of hydrogen bonding between the HF molecules Check-up 13 The table lists the boiling points of some Group V hydrides Hydride ammonia, NH3 phosphine, PH3 arsine, AsH3 stibine, SbH3 Boiling point / °C −33 −88 −55 −17 a Explain the trend in the boiling points from phosphine to stibine b Explain why the boiling point of ammonia does not follow this trend The peculiar properties of water Enthalpy change of vaporisation and boiling point Water has a much higher enthalpy change of vaporisation and boiling point than expected This is due to its extensive hydrogen bonding Figure 4.43 shows the enthalpy changes of vaporisation of water and other Group VI hydrides The rise in enthalpy change of vaporisation from H2S to H2Te is due to the increasing number of electrons in the Group VI atoms as we go down the group This leads to increased van der Waals’ forces as the molecules get bigger If water only had van der Waals’ forces between its molecules, we would expect its enthalpy change to be about 17 kJ mol−1 But the enthalpy change of vaporisation of water is much higher This is because water is extensively hydrogen bonded The boiling point of water is also much higher than predicted by the trend in boiling points for the other Group VI hydrides This Chemical bonding 67 50 Enthalpy change of vaporisation / kJ mol–1 H 2O 40 30 H2Te H2S 20 H2Se 10 20 40 60 80 100 Number of electrons 120 140 Figure 4.43 Enthalpy changes of vaporisation for Group VI hydrides plotted against number of electrons present Figure 4.44 Ice floats on water also indicates that much more energy is required to break the bonds between water molecules compared with other hydrides of Group VI elements molecules This produces a rigid lattice in which each oxygen atom is surrounded by a tetrahedron of hydrogen atoms This ‘more open’ arrangement allows the water molecules to be slightly further apart than in the liquid (Figure 4.45) So the density of ice is less than that of liquid water Surface tension and viscosity Water has a high surface tension and high viscosity Hydrogen bonding reduces the ability of water molecules to slide over each other, so the viscosity of water is high The hydrogen bonds in water also exert a significant downward force at the surface of the liquid This causes the surface tension of water to be higher than for most liquids Fact file You can float a needle on water! Place a small piece of tissue paper on the surface of some water in a bowl The water surface must be still Place a needle on the tissue The paper will sink slowly, leaving the needle to float The hydrogen bonding in the water gives it a high surface tension If you then add some detergent, the needle will sink Ice is less dense than water Most solids are denser than their liquids This is because the molecules are more closely packed in the solid state But this is not true of water In ice, there is a three-dimensional hydrogen-bonded network of water 68 Chemical bonding Figure 4.45 A model of ice Oxygen atoms are red, hydrogen atoms are white, hydrogen bonds are lilac This hydrogen-bonded arrangement makes ice less dense than water 4.7 Bonding and physical properties The type of bonding between atoms, ions or molecules influences the physical properties of a substance Physical state at room temperature and pressure Ionic compounds Ionic compounds are solids at room temperature and pressure This is because: • there are strong electrostatic forces (ionic bonds) holding the positive and negative ions together • the ions are regularly arranged in a lattice (see Chapter 5), with the oppositely charged ions close to each other Ionic compounds have high melting points, high boiling points and high enthalpy changes of vaporisation It takes a lot of energy to overcome the strong electrostatic attractive forces Metals Metals, apart from mercury, are solids Most metals have high melting points, high boiling points and high enthalpy changes of vaporisation This is because it takes a lot of energy to overcome the strong attractive forces between the positive ions and the ‘sea’ of delocalised electrons Covalent compounds Covalently bonded substances with a simple molecular structure, for example water and ammonia, are usually liquids or gases This is because the forces between the molecules are weak It does not take much energy to overcome these intermolecular forces, so these substances have low melting points, low boiling points and low enthalpy changes of vaporisation compared with ionic compounds Some substances that have covalently bonded molecules may be solids at room temperature, for example iodine and poly(ethene) These are usually molecules where the van der Waals’ forces are considerable However, the melting points of these substances are still fairly low compared with ionic compounds or most metals Solubility attracted to the ions on the surface of the ionic solid These attractions are called ion–dipole attractions (see page 263) These attractions replace the electrostatic forces between the ions and the ions go into solution Metals Metals not dissolve in water However, some metals, for example sodium and calcium, react with water Covalent compounds Covalently bonded substances with a simple molecular structure fall into two groups • Those that are insoluble in water Most covalently bonded molecules are non-polar Water molecules are not attracted to them so they are insoluble An example is iodine • Those that are soluble in water Small molecules that can form hydrogen bonds with water are generally soluble An example is ethanol, C2H5OH Some covalently bonded substances react with water rather than dissolving in it For example, hydrogen chloride reacts with water to form hydrogen ions and chloride ions, and the ions are soluble Silicon chloride reacts with water to form hydrogen ions, chloride ions and silicon dioxide This reaction is called a hydrolysis reaction Electrical conductivity Ionic compounds Ionic compounds not conduct electricity when in the solid state This is because the ions are fixed in the lattice and are not free to move When molten, an ionic compound conducts electricity because the ions are free to move Metals Metals conduct electricity both when solid and when molten This is because the delocalised electrons are free to move Covalent compounds Covalently bonded substances with a simple molecular structure not conduct electricity This is because they have neither ions nor electrons which are free to move Ionic compounds Most ionic compounds are soluble in water This is because water molecules are polar and they are Chemical bonding 69 Check-up 14 Explain the following differences in terms of the type of bonding present a Aluminium oxide has a melting point of 2980 °C but aluminium chloride changes to a vapour at 178 °C b Magnesium chloride conducts electricity when molten but not when solid c Iron conducts electricity when solid but the ionic solid iron(II) chloride does not conduct when solid d Sodium sulfate dissolves in water but sulfur does not e Propanol, CH3CH2CH2OH, is soluble in water but propane, CH3CH2CH3, is not f A solution of hydrogen chloride in water conducts electricity continued Summary Ions are formed when atoms gain or lose electrons Ionic (electrovalent) bonding involves an attractive force between positively and negatively charged ions A covalent bond is formed when atoms share a pair of electrons When atoms form covalent or ionic bonds each atom or ion has a full outer electron shell of electrons (Some covalent compounds may be electron deficient or have an ‘expanded octet’.) Dot-and-cross diagrams can be drawn to show the arrangement of electrons in ionic and covalent compounds In dative covalent bonding one atom provides both electrons in the formation of the covalent bond The shapes and bond angles in molecules can be predicted using the idea that lone pairs of electrons repel other lone pairs more than bond pair electrons, and that bond pair to bond pair repulsion is least σ bonds (sigma bonds) are formed by end-on overlap of atomic orbitals whereas π bonds (pi bonds) are formed by sideways overlap of p-type atomic orbitals Three types of relatively weak intermolecular forces are hydrogen bonding, permanent dipole–dipole forces and van der Waals’ forces Electronegativity differences can be used to predict the type of weak intermolecular forces between molecules Hydrogen bonding occurs between molecules that have a hydrogen atom covalently bonded to an atom of a very electronegative element (fluorine, oxygen or nitrogen) The reactivities of covalent bonds can be explained in terms of bond energy, bond length and bond polarity Intermolecular forces are based on either permanent dipoles, as in CHCl3(l), or temporary induced dipoles (van der Waals’ forces), as in Br2(l) Metallic bonding can be explained in terms of a lattice of positive ions surrounded by mobile electrons The physical properties of substances may predicted from the type of bonding present Substances with ionic bonding have high melting and boiling points, whereas simple molecules with covalent bonding have low melting points The presence of hydrogen bonding in a molecule influences its melting point and boiling point 70 Chemical bonding End-of-chapter questions The table shows the atomic number and boiling points of some noble gases Gas Atomic number Boiling point / °C helium −253 neon 10 −246 argon 18 −186 krypton 36 −152 xenon 54 −107 [2] a Use ideas about forces between atoms to explain this trend in boiling points b Xenon forms a number of covalently bonded compounds with fluorine i What you understand by the term covalent bond? ii Draw a dot-and-cross diagram for xenon tetrafluoride, XeF4 iii Suggest a shape for XeF4 Explain why you chose this shape c The structure of xenon trioxide is shown below [1] [1] [3] Xe O O O i By referring to electron pairs, explain why xenon trioxide has this shape [2] ii Draw the structure of xenon trioxide to show the partial charges on the atoms and the direction of the dipole in the molecule [2] Total = 11 Aluminium chloride, AlCl3, and ammonia, NH3, are both covalent molecules a i Draw a diagram of an ammonia molecule, showing its shape Show any lone pairs of electrons ii State the bond angle N H H [3] [1] in the ammonia molecule b Explain why ammonia is a polar molecule c An ammonia molecule and an aluminium chloride molecule can join together by forming a co-ordinate bond i Explain how a co-ordinate bond is formed ii Draw a dot-and-cross diagram to show the bonding in the compound formed between ammonia and aluminium chloride, H3NAlCl3 (Use a • for a nitrogen electron, a ∘ for an aluminium electron and an × for the hydrogen and chlorine electrons.) d Aluminium chloride molecules join together to form a compound with the formula Al2Cl6 Draw a displayed formula (showing all atoms and bonds) to show the bonding in one Al2Cl6 molecule Show the dative covalent bonds by arrows [2] [1] [3] [2] Total = 12 Chemical bonding 71 Electronegativity values can be used to predict the polarity of bonds a Explain the term electronegativity [2] b The electronegativity values for some atoms are given below: H = 2.1, C = 2.5, F = 4.0, Cl = 3.0, I = 2.5 Use these values to predict the polarity of each of the following bonds by copying the bonded atoms shown below and adding δ+ or δ− above each atom i H I ii F I [2] iii C Cl c The shape of iodine trichloride, ICl3, is shown below Cl Cl I Cl [2] [2] [1] i Describe the shape of this molecule ii Explain why the ICl3 molecule has this shape iii Suggest a value for the Cl I Cl bond angle d The boiling points of the hydrogen halides are shown in the table Hydrogen halide Boiling point / °C HF +20 HCl −85 HBr −67 HI −35 i Explain the trend in boiling points from HCl to HI ii Explain why the boiling point of HF is so much higher than the boiling point of HCl e Tetrachloromethane, CCl4, is a non-polar molecule i Draw a diagram to show the shape of this molecule ii Explain why this molecule is non-polar [2] [1] Total = 17 The diagram below shows part of a giant metallic structure – +e – e e– + + e– e– + + e– e– + –+ e + a b c d 72 [2] [3] – +e – e e– + + e– e– + + e– e– + –+ e + + + + + e– e– e– e– Use this diagram to explain the main features of metallic bonding Explain why metals are good conductors of electricity Explain why, in general, metals have high melting points Suggest why potassium is a better conductor of electricity than lithium Chemical bonding [3] [2] [2] [4] Total = 11 Methane, CH4, is a gas at room temperature a Explain why methane is a gas at room temperature b Draw a diagram to show the shape of a molecule of methane On your diagram show a value for the C H H [2] [3] bond angle c Perfumes often contain molecules that have simple molecular structures Explain why [2] d When a negatively charged rod is held next to a stream of propanone, CH3COCH3, the stream of propanone is attracted to the rod Draw the full structure of a molecule of propanone and use your diagram to explain why the stream of propanone is attracted to the rod [3] Total = 10 Sodium iodide and magnesium oxide are ionic compounds Iodine and oxygen are covalent molecules a Draw dot-and-cross diagrams for: i magnesium oxide ii oxygen [2] b How sodium iodide and iodine differ in their solubility in water? Explain your answer [3] c Explain why molten sodium iodide conducts electricity but molten iodine does not [2] d The boiling point of sodium iodide is 1304 °C The boiling point of iodine is 184 °C Explain this difference [5] Total = 12 Hydrogen sulfide, H2S, is a covalent compound a Draw a dot-and-cross diagram for hydrogen sulfide b Draw a diagram of a hydrogen sulfide molecule to show its shape Show on your diagram: i the value of the S H H [2] bond angle ii the partial charges on each atom as δ+ or δ− iii an arrow showing the exact direction of the dipole in the molecule as a whole c Oxygen, O, sulfur, S, and selenium, Se, are in the same group in the Periodic Table i Explain why hydrogen selenide, H2Se, has a higher boiling point than hydrogen sulfide, H2S ii Explain why the boiling point of water is so much higher than the boiling point of hydrogen sulfide [4] [2] [5] Total = 13 The table shows the type of bonding in a number of elements and compounds Element or compound Fe, Na NaCl, MgCl2 CO2, Br2 Type of bonding metallic ionic covalent within the molecules a Draw a labelled diagram to show metallic bonding b Explain why magnesium chloride has a high melting point but bromine has a low melting point [2] [5] Chemical bonding 73 c Explain why solid sodium conducts electricity but solid sodium chloride does not conduct electricity [2] d i Draw a dot-and-cross diagram for carbon dioxide [1] ii Describe the shape of the carbon dioxide molecule [1] iii Explain why a carbon dioxide molecule has this shape [2] e Bromine is a liquid at room temperature Weak van der Waals’ forces hold the bromine molecules together Describe how van der Waals’ forces arise [5] Total = 18 Water is extensively hydrogen bonded This gives it anomalous (peculiar) properties a Explain why ice is less dense than liquid water b State two other anomalous properties of water c Propanone has the structure shown below [3] [2] CH3 C O CH3 When propanone dissolves in water, it forms a hydrogen bond with water i What features must water and propanone molecules posses in order to form a hydrogen bond? ii Draw a diagram to show a propanone molecule and a water molecule forming a hydrogen bond d Propanone has a double bond One of the bonds is a σ bond (sigma bond) The other is a π bond (pi bond) i Explain the difference between a σ bond and a π bond in terms of how they are formed ii Copy the diagram, then complete it to show the shapes of the electron clouds in the σ bond and the π bond between the carbon atoms in ethene Label your diagram H [3] [3] H C H [2] [2] C H Total = 15 74 Chemical bonding Test yourself Chapter Which one of the following is the most likely Cl C Cl bond angle in tetrachloromethane, CCl4? A 90° B 107.5° C 109° D 120° Which one of these statements about boron trifluoride, BF3, is correct? A The boron atom in boron trifluoride has a lone pair of electrons B Boron trifluoride has a tetrahedral structure C The F B F bond angle in boron trifluoride is approximately 107° D Boron trifluoride can form a co-ordinate bond with ammonia Which one of the following pairs of substances not form hydrogen bonds with each other? A CH3Cl and CH3COCH3 B CH3OH and CH3CH2OH C NH3 and CH3OH D H2O and CH3CH2OH AS and A Level Chemistry © Cambridge University Press Test yourself: Chapter 4 Which one of these statements about the bonding between the carbon atoms in ethene is correct? A The σ bond is formed by the sideways overlap of modified p orbitals B The π bond is formed by the sideways overlap of modified p orbitals C The π bond is formed by the end-on (linear) overlap of modified p orbitals D The π bond contains two pairs of electrons In which one of the following substances would you expect the total van der Waals’ forces between the molecules to be the greatest? A CH3CH2CH(CH3)CH3 B CH3CH2CH2CH2CH3 C CH3Br D H3C CH3 C H2C CH2 Which one of the following statements about metallic bonding is false? A The strength of metallic bonding increases with increasing size of the metal ions B The strength of metallic bonding increases with increasing number of delocalised electrons in the metallic structure C In metallic bonds, the ions are held together by electrostatic attractions between the ions and the delocalised electrons D The strength of metallic bonding increases with the increasing positive charge on the metal ions Test yourself: Chapter AS and A Level Chemistry © Cambridge University Press Which one of the following substances is very soluble in water? A Iodine, I2 B Carbon monoxide, CO C Sodium nitrate, NaNO3 D Propane, CH3CH2CH3 Which one of the following substances has the highest boiling point? A CH3CH2CH2CH3 B CH3CH2Cl C CH3COCH3 D CH3CH2CH2OH Which one of the following describes the shape of a PCl3 molecule? A Linear B Octahedral C Planar D Pyramidal AS and A Level Chemistry © Cambridge University Press Test yourself: Chapter 4 10 Which one of these statements about covalent bonding is false? A The bond energy of a C  C bond is numerically greater than the bond energy of a C C bond B The bond energy of H I is numerically greater than the bond energy of H Br C Once a co-ordinate bond is formed, it cannot be distinguished from a covalent bond formed in the normal way D Sulfur has more than eight electrons in its outer principal energy level when it forms covalent bonds with more than two fluorine atoms Test yourself: Chapter AS and A Level Chemistry © Cambridge University Press States of matter Learning outcomes Candidates should be able to: state the basic assumptions of the kinetic theory as applied to an ideal gas explain qualitatively in terms of intermolecular forces and molecular size – the conditions necessary for a gas to approach ideal behaviour – the limitations of ideality at very high pressures and very low temperatures state and use the general gas equation pV = nRT in calculations, including the determination of Mr describe, using a kinetic-molecular model, the liquid state, melting, vaporisation and vapour pressure describe in simple terms the lattice structure of a crystalline solid which is – ionic, as in sodium chloride, magnesium oxide – simple molecular, as in iodine – giant molecular, as in graphite, diamond, silicon(IV) oxide 5.1 States of matter In the last chapter we looked at the types of forces which keep the particles in solids and liquids together and make it possible to liquefy gases In this chapter, we shall also consider how the closeness and motion of the particles influences the properties of these three states of matter (Figure 5.1) – hydrogen bonded, as in ice – metallic, as in copper (the concept of the ‘unit cell’ is not required) explain the strength, high melting point and insulating properties of ceramics in terms of their giant molecular structure relate the uses of ceramics based on magnesium oxide, aluminium oxide and silicon(IV) oxide to their properties, e.g furnace linings, electrical insulators, glass, crockery describe and interpret the uses of the metals aluminium (including its alloys) and copper (including brass) in terms of their physical properties understand that materials are a finite resource and the importance of recycling processes suggest from quoted physical data the type of structure and bonding present in a substance Gases have no fixed shape or volume Gas particles: • are far apart, therefore gases can be compressed • are randomly arranged • can move freely from place to place, in all directions Liquids take the shape of the container they occupy Liquid particles: • are close together, so liquids have a fixed volume and can only be compressed slightly Figure 5.1 The three states of water are ice, water and steam The ‘steam’ we see from the kettle is condensed droplets of water The real gaseous water is in the area between this condensation and the spout of the kettle We can’t see it because it is colourless States of matter 75 ... the amount of a particular substance in a solution The relative atomic mass is the weighted average mass of naturally occurring atoms of an element on a scale where an atom of carbon-12 has a mass... Roger Norris, Lawrie Ryan and David Acaster Cambridge International AS and A Level Chemistry Coursebook c a mb r id g e u n i ve r s i t y p re s s Cambridge, New York, Melbourne, Madrid, Cape... bromide, magnesium sulfate and ammonium nitrate Acids and alkalis also contain ions For example H+(aq) and Cl−(aq) ions are present in hydrochloric acid and Na+(aq) and OH−(aq) ions are present in sodium