Preview Cambridge International AS and A Level Chemistry Coursebook, 1st Edition by Roger Norris, Lawrie Ryan and David Acaster (2011)

Preview Cambridge International AS and A Level Chemistry Coursebook, 1st Edition by Roger Norris, Lawrie Ryan and David Acaster (2011) Preview Cambridge International AS and A Level Chemistry Coursebook, 1st Edition by Roger Norris, Lawrie Ryan and David Acaster (2011) Preview Cambridge International AS and A Level Chemistry Coursebook, 1st Edition by Roger Norris, Lawrie Ryan and David Acaster (2011) Preview Cambridge International AS and A Level Chemistry Coursebook, 1st Edition by Roger Norris, Lawrie Ryan and David Acaster (2011)

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 What is the concentration of chloride ions in a solution containing 0.02 mol of calcium chloride, CaCl2, in 200 cm3 of solution? A 0.01 mol dm−3 B 0.02 mol dm−3 C 0.1 mol dm−3 D 0.2 mol dm−3 What information is given by a molecular formula? A The ratio of each type of atom in a molecule B The number of each type of atom in a molecule C The arrangement of the atoms in a molecule D The number of atoms in a molecule The equation for a reaction is: Ba(NO3)2(aq) + Na2SO4(aq) → 2NaNO3(aq) + BaSO4(s) What is the ionic equation for the reaction? A Ba2+(aq) + Na2SO4(aq) → 2Na+(aq) + BaSO4(s) B Ba(NO3)2(aq) + 2Na+(aq) → 2NaNO3(aq) + Ba2+(s) C Ba2+(aq) + SO42−(aq) → Ba2+(s) + SO42−(aq) D Ba2+(aq) + SO42−(aq) → BaSO4(s) Test yourself: Chapter AS and A Level Chemistry © Cambridge University Press An oxide of copper contains 0.635 g of copper and 0.080 g of oxygen by mass (Ar values: Cu = 63.5, O = 16.0) What is the empirical formula of this oxide of copper? A Cu2O B CuO C CuO2 D Cu2O3 100 cm3 of an aqueous solution of sodium hydroxide of concentration 0.10 mol dm−3 reacts with 200 cm3 of aqueous phosphoric acid The equation for the reaction is: 2NaOH + H3PO4 → Na2HPO4 + 2H2O What is the concentration of the phosphoric acid? A 0.025 mol dm−3 B 0.050 mol dm−3 C 0.10 mol dm−3 D 0.040 mol dm−3 Which one of the following reactions is accompanied by the largest percentage increase in volume? A 4NH3(g) + 5O2(g) → 4NO(g) + 6H2O(g) B 2NH3(g) → N2(g) + 3H2(g) C CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) D S(s) + O2(g) → SO2(g) AS and A Level Chemistry © Cambridge University Press Test yourself: Chapter 10 Hydrochloric acid reacts with barium hydroxide: 2HCl + Ba(OH)2 → BaCl2 + 2H2O What is the volume of hydrochloric acid of concentration 0.050 mol dm−3 which exactly neutralises 25.0 cm3 of a 0.10 mol dm−3 solution of barium hydroxide? A 25 cm3 B 50 cm3 C 100 cm3 D 200 cm3 Test yourself: Chapter AS and A Level Chemistry © Cambridge University Press Atomic structure Learning outcomes Candidates should be able to: identify and describe protons, neutrons and electrons in terms of their relative charges and relative masses deduce the behaviour of beams of protons, neutrons and electrons in electric fields describe the distribution of mass and charges within an atom deduce the numbers of protons, neutrons and electrons present in both atoms and ions given proton and nucleon numbers (and charge) describe the contribution of protons and neutrons to atomic nuclei in terms of proton number and nucleon number distinguish between isotopes on the basis of different numbers of neutrons present 2.1 Elements and atoms Every substance in our world is made up from chemical elements These chemical elements cannot be broken down further into simpler substances by chemical means A few elements, such as nitrogen and gold, are found on their own in nature, not combined with other elements Most elements, however, are found in combination with other elements as compounds Every element has its own chemical symbol The symbols are often derived from Latin or Greek words Some examples are shown in Table 2.1 Chemical elements contain only one type of atom An atom is the smallest part of an element that can take part in a chemical change Atoms are very small The diameter of a hydrogen atom is approximately 10−10 m, so the mass of an atom must also be very small A single hydrogen atom weighs only 1.67 × 10−27 kg Element carbon lithium iron potassium Symbol C Li (from Greek ‘lithos’) Fe (from Latin ‘ferrum’) K (from Arabic ‘al-qualyah’ or from the Latin ‘kalium’) Table 2.1 Some examples of chemical symbols Figure 2.1 Our Sun is made largely of the elements hydrogen and helium This is a composite image made using X-ray and solar optical telecopes 2.2 Inside the atom The structure of an atom Every atom has nearly all of its mass concentrated in a tiny region in the centre of the atom called the nucleus The nucleus is made up of particles called nucleons There are two types of nucleon: protons and neutrons Atoms of different elements have different numbers of protons Outside the nucleus, particles called electrons move around in regions of space called orbitals (see page 38) Atomic structure 25 electron Fact file Nanotechnology is the design and making of objects that may have a thickness of only a few thousand atoms or less Groups of atoms can be moved around on special surfaces In this way scientists hope to develop tiny machines that may help deliver medical drugs to exactly where they are needed in the body Chemists often find it convenient to use a model of the atom in which electrons move around the nucleus in electron shells Each shell is a certain distance from the nucleus at its own particular energy level (see page 37) In a neutral atom, the number of electrons is equal to the number of protons A simple model of a carbon atom is shown in Figure 2.3 Atoms are tiny, but the nucleus of an atom is far tinier still If the diameter of an atom were the size of a football stadium, the nucleus would only be the size of a pea This means that most of the atom is empty space! Electrons are even smaller than protons and neutrons nucleus electron shells (energy levels) neutron proton Figure 2.3 A model of a carbon atom This model is not very accurate but it is useful for understanding what happens to the electrons during chemical reactions – Experiments with sub-atomic particles We can deduce the electric charge of sub-atomic particles by showing how beams of electrons, protons and neutrons behave in electromagnetic fields If we fire a beam of electrons past electrically charged plates, the electrons are deflected (bent) away from the negative plate and towards the positive plate (Figure 2.4) This shows us that the electrons are negatively charged A cathode-ray tube (Figure 2.5) can be used to produce beams of electrons At one end of the tube is a metal wire (cathode) which is heated to a high temperature when a electron beam + Figure 2.4 The beam of electrons is deflected away from a negatively charged plate and towards a positively charged plate – cathode + fluorescent screen with scale cathode rays charged plates (anode) magnets causing electromagnetic field Figure 2.2 Ernest Rutherford (left) and Hans Geiger (right) using their alpha-particle apparatus Interpretation of the results led to Rutherford proposing the nuclear model for atoms 26 Atomic structure beam deflected downwards Figure 2.5 The electron beam in a cathode-ray tube is deflected (bent) by an electromagnetic field The direction of the deflection shows us that the electron is negatively charged low voltage is applied to it At the other end of the tube is a fluorescent screen which glows when electrons hit it The electrons are given off from the heated wire and are attracted towards two metal plates which are positively charged As they pass through the metal plates the electrons form a beam When the electron beam hits the screen a spot of light is produced When an electromagnetic field is applied across this beam the electrons are deflected (bent) The fact that the electrons are so easily attracted to the positively charged anode and that they are easily deflected by an electromagnetic field shows us that: • electrons have a negative charge • electrons have a very small mass In recent years, experiments have been carried out with beams of electrons, protons and neutrons The results of these experiments show that: • a proton beam is deflected away from a positively charged plate; since like charges repel, the protons must have a positive charge (Figure 2.7) • an electron beam is deflected towards a positively charged plate; since unlike charges attract, the electrons must have a negative charge • a beam of neutrons is not deflected; this is because they are uncharged – beam of protons protons detected on walls of apparatus + + + Figure 2.7 A beam of protons is deflected away from a positively charged area This shows us that protons have a positive charge In these experiments, huge voltages have to be used to show the deflection of the proton beam This contrasts with the very low voltages needed to show the deflection of an electron beam These experiments show us that protons are much heavier than electrons If we used the same voltage to deflect electrons and protons, the beam of electrons would have a far greater deflection than the beam of protons This is because a proton is about 2000 times as heavy as an electron Check-up A beam of electrons is passing close to a highly negatively charged plate When the electrons pass close to the plate, they are deflected (bent) away from the plate a What deflection would you expect, if any, when the experiment is repeated with beams of i protons and ii neutrons? Explain your answers b Which sub-atomic particle (electron, proton or neutron) would be deviated the most? Explain your answer Figure 2.6 J J Thomson calculated the charge to mass ratio of electrons He used results from experiments with electrons in cathode-ray tubes Fact file Atomic scientists now believe that elementary particles called quarks and leptons are the building blocks from which most matter is made They think that protons and neutrons are made up from quarks and that an electron is a type of lepton Masses and charges: a summary Electrons, protons and neutrons have characteristic charges and masses The values of these are too small to be very useful when discussing general chemical properties For example, the charge on a single electron is −1.602 × 10−19 coulombs We therefore compare their masses and charges by using their relative charges and masses These are shown in Table 2.2 Atomic structure 27 Sub-atomic particle electron Symbol Relative mass e 1836 neutron proton n p 1 Relative charge −1 +1 Atom vanadium strontium phosphorus Nucleon number 51 84 31 Proton number 23 38 15 Table 2.2 Comparing electrons, neutrons and protons 2.3 Numbers of nucleons Proton number and nucleon number The number of protons in the nucleus of an atom is called the proton number (Z) It is also known as the atomic number Every atom of the same element has the same number of protons in its nucleus It is the proton number which makes an atom what it is For example an atom with a proton number of 11 must be an atom of the element sodium The Periodic Table of elements is arranged in order of the proton numbers of the individual elements (see Appendix 1, page 497) The nucleon number (A) is the number of protons plus neutrons in the nucleus of an atom This is also known as the mass number How many neutrons? We can use the nucleon number and proton number to find the number of neutrons in an atom Since nucleon number = number of protons + number of neutrons Then number of neutrons = nucleon number − number of protons = A−Z For example, an atom of aluminium has a nucleon number of 27 and a proton number of 13 So an aluminium atom has 27 − 13 = 14 neutrons Isotopes All atoms of the same element have the same number of protons However, they may have different numbers of neutrons Atoms of the same element which have differing numbers of neutrons are called isotopes Isotopes are atoms of the same element with different nucleon (mass) numbers Isotopes of a particular element have the same chemical properties because they have the same number of electrons They have slightly different physical properties, such as small differences in density We can write symbols for isotopes We write the nucleon number at the top left of the chemical symbol and the proton number at the bottom left The symbol for the isotope of boron with protons and 11 nucleons is written: nucleon number → 11 proton number → B Hydrogen has three isotopes The atomic structure and isotopic symbols for the three isotopes of hydrogen are shown in Figure 2.8 When writing generally about isotopes, chemists also name them by omitting the proton number and placing Check-up Fact file Use the information in the table to deduce the number of electrons and neutrons in a neutral atom of: a vanadium b strontium c phosphorus continued 28 Atomic structure Isotopes can be radioactive or non-radioactive Specific radioisotopes (radioactive isotopes) can be used to check for leaks in oil or gas pipelines and to check the thickness of paper They are also used in medicine to treat some types of cancer and to check the activity of the thyroid gland in the throat protium electron deuterium tritium 1 neutron proton protons neutrons isotopic symbol 1 H 2 H H Figure 2.8 The atomic structure and isotopic symbols for the three isotopes of hydrogen the nucleon number after the name For example, the isotopes of hydrogen can be called hydrogen-1, hydrogen-2 and hydrogen-3 Check-up Use the Periodic Table on page 497 to help you a Write isotopic symbols for the following neutral atoms: i bromine-81 ii calcium-44 iii iron-58 iv palladium-110 b What is the number of protons and neutrons in each of these atoms? 2.4 How many protons, neutrons and electrons? In a neutral atom the number of positively charged protons in the nucleus equals the number of negatively charged electrons outside the nucleus When an atom gains or loses electrons, ions are formed which are electrically charged For example: The chloride ion has a single negative charge because there are 17 protons (+) and 18 electrons (−) Mg → magnesium atom 12 protons 12 electrons Mg2+ + magnesium ion 12 protons 10 electrons 2e− electrons removed The magnesium ion has a charge of 2+ because it has 12 protons (+) but only 10 electrons (−) The isotopic symbol for an ion derived from sulfur33 2− 33 is 16 S This sulfide ion has 16 protons, 17 neutrons (because 33 − 16 = 17) and 18 electrons (because 16 + = 18) Check-up Deduce the number of electrons in each of these ions: a 40 19 b 15 c 18 d 71 31 K+ N3− O2− Ga3+ Cl + e− Cl− → chlorine atom electron chloride ion gained 17 protons 17 protons 18 electrons 17 electrons Atomic structure 29 Summary Every atom has an internal structure with a nucleus in the centre and the negatively charged electrons arranged in ‘shells’ outside the nucleus Most of the mass of the atom is in the nucleus, which contains protons (positively charged) and neutrons (uncharged) Beams of protons and electrons are deflected by electric fields but neutrons are not All atoms of the same element have the same number of protons This is the proton number (Z), which is also called the atomic number The nucleon number, which is also called the mass number (A), is the total number of protons and neutrons in an atom The number of neutrons in an atom is found by subtracting the proton number from the nucleon number (A − Z) In a neutral atom, number of electrons = number of protons When there are more protons than electrons the atom becomes a positive ion When there are more electrons than protons, a negatively charged ion is formed Isotopes are atoms with the same atomic number but different nucleon numbers They only differ in the number of neutrons they contain End-of-chapter questions 30 Boron is an element in Group III of the Periodic Table a Boron has two isotopes What you understand by the term isotope? [1] 11 b State the number of i protons, ii neutrons and iii electrons in one neutral atom of the isotope 5B [3] c State the relative masses and charges of: i an electron [2] ii a neutron [2] iii a proton [2] Total = 10 Zirconium, Zr, and hafnium, Hf, are metals An isotope of zirconium has 40 protons and 91 nucleons a i Write the isotopic symbol for this isotope of zirconium [1] ii How many neutrons are present in one atom of this isotope? [1] 180 2+ b Hafnium ions, 72Hf , are produced in a mass spectrometer How many electrons are present in one of these hafnium ions? [1] Atomic structure c The sub-atomic particles present in zirconium and hafnium are electrons, neutrons and protons A beam of protons is fired into an electric field produced by two charged plates, as shown in the diagram + beam of protons – i Describe how the beam of protons behaves when it passes through the gap between the charged plates Explain your answer ii Describe and explain what happens when a beam of neutrons passes through the gap between the charged plates a Describe the structure of an atom, giving details of the sub-atomic particles present b Explain the terms atomic number and nucleon number c Copy and complete the table: Neutral atom Mg Al Atomic number 12 13 Nucleon number 24 27 [2] Total = [6] [2] Numbers of each sub-atomic particle present d Explain why atoms are neutral e An oxygen atom has protons in its nucleus Explain why it cannot have protons f When calculating the mass of an atom, the electrons are not used in the calculation Explain why not [2] [2] [1] [1] [1] Total = 13 The symbols below describe two isotopes of the element uranium 235 92 U 238 92 U a State the meaning of the term isotope b i In what ways are these two isotopes of uranium identical? ii In what ways they differ? c In a mass spectrometer uranium atoms can be converted to uranium ions, U2+ State the number of electrons present in one U2+ ion [1] [2] [2] [1] Total = Atomic structure 31 The table below shows the two naturally occurring isotopes of chlorine a Copy and complete the table 35 17 Cl 37 17 Cl number of protons number of electrons number of neutrons [3] b The relative atomic mass of chlorine is 35.5 What does this tell you about the relative abundance of the two naturally occurring isotopes of chlorine? c Magnesium chloride contains magnesium ions, Mg2+, and chloride ions, Cl− i Explain why a magnesium ion is positively charged ii Explain why a chloride ion has a single negative charge 32 Atomic structure [2] [1] [2] Total = Test yourself Chapter An isotope of a neutral gallium atom has the symbol 69 31 Ga Which particles are present in one atom of this isotope? A 31 protons, 69 neutrons and 31 electrons B 31 protons, 38 neutrons and 31 electrons C 31 protons, 38 neutrons and 38 electrons D 69 protons, 38 neutrons and 69 electrons What are the relative masses of an electron, a proton and a neutron? A proton = 1, electron = , neutron = 1836 B proton = 1, electron = , neutron = 183 C proton = 1, electron = 1, neutron = D proton = 1, electron = Which one of these statements best describes isotopes? A Atoms of the same element with different numbers of neutrons B Atoms of different elements with different numbers of nucleons C Atoms of the same element with different numbers of electrons D Atoms of different elements with the same number of nucleons AS and A Level Chemistry © Cambridge University Press , neutron = 1836 Test yourself: Chapter 2 Which one of these statements is correct? A If an element occurs naturally it only has a single isotope B Relative isotopic masses are always whole numbers C The accurate relative atomic mass of an atom of carbon-12 is 12.001 D The relative atomic mass of an element is the mass of an atom of the element compared with the mass of an atom of carbon-12 Which one of the following values is the correct relative molecular mass of iron(III) sulfate, Fe2(SO4)3? (Ar values: Fe = 55.8, S = 32.1, O = 16.0) A 399.9 B 376.2 C 344.1 D 303.8 An electron beam is deflected at a greater angle than a proton beam when both are exposed to the same strong electric field Why is this? A An electron has a negative charge while a proton has a positive charge B An electron has a much larger mass than a proton C An electron has a much smaller mass than a proton D An electron has more electrical energy than a proton Test yourself: Chapter AS and A Level Chemistry © Cambridge University Press A sample of magnesium obtained from a meteorite has the isotopic composition, 24Mg (74%), 25Mg (10%), 26Mg (16%) Which one of the following is most likely to be the correct relative atomic mass of the sample of magnesium, to 1 decimal place? A 24.2 B 24.3 C 24.4 D 24.5 31 3− The number of electrons in the phosphide ion 15 P is: A 12 B 15 C 18 D 28 Which one of these statements about the mass spectrum for krypton, shown below, is correct? 78 79 80 81 82 83 17.37% 11.55% 11.56% 0.35% 2.27% 56.90% 84 85 86 A The vertical axis shows the mass of the ions produced B The horizontal axis shows the mass/charge ratio of the ions produced C The accurate relative atomic mass of krypton is 85.1 D There are five isotopes of krypton AS and A Level Chemistry © Cambridge University Press Test yourself: Chapter 10 35 Chlorine has two isotopes, 17 Cl and 37 17 Cl A chlorine molecule has the formula Cl2 Deduce the maximum number of different chlorine molecule ions, Cl2+, that can appear in a mass spectrum of chlorine A B C D Test yourself: Chapter AS and A Level Chemistry © Cambridge University Press ... 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... 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... 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