CGP A-Level Chemistry Edexcel Year 1 & 2 Revision PDF

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CGP CGP CGP P G C A-Level Chemistry Edexcel The new A-Levels are seriously tough... A-Level But don’t worry — CGP have come to the rescue with this fantastic all-in-one book! Short, sharp revision notes for every topic... Chemistry Exam Board: Edexcel No pointless rambling, just the relevant info Exam-style questions to test your skills... Of course — all with detailed answers included 100% matched to the new A-Level course... Perfect for both years of Edexcel A-Level Chemistry Plus a brilliant free Online Edition... So you can read the whole book on your computer or tablet! Complete Revision & Practice If you can find a better revision guide from anyone else, we’ll eat our hats ☻ Complete Revision & Practice CGP Price Promise: If you can buy this book cheaper somewhere else, then ‘hurrah for you’ ISBN 978 1 78294 301 3 P G C 9 781782 943013 CER72 £19.99 (Retail Price) www.cgpbooks.co.uk  C G P CGP or browse our books at are now available on computer or tablet! Our bestselling books See our full Kindle range at cgpbooks.co.uk/kindle......perfect for revision on your The Periodic Table 0 1.0 4.0 H hydrogen He 2 helium 1 1 3 4 5 6 7 2 9.0 relative 10.8 12.0 14.0 16.0 19.0 20.2 6.9 Be atomic mass Li lithium beryllium B boron C carbon N nitrogen O oxygen F fluorine Ne neon 3 4 5 6 7 8 9 10 23.0 24.3 atomic (proton) number 27.0 28.1 35.5 31.0 32.1 39.9 Na Mg sodium Al aluminium Si silicon P phosphorus S sulfur Cl chlorine Ar argon magnesium 11 12 13 14 15 16 17 18 39.1 40.1 45.0 47.9 50.9 52.0 54.9 55.8 58.9 58.7 63.5 65.4 69.7 72.6 74.9 79.0 79.9 83.8 potassium K Ca calcium Sc scandium Ti titanium V Cr vanadium chromium manganese Mn Fe iron Co cobalt Ni nickel Cu copper Zn zinc Ga germanium Ge arsenic gallium As selenium Se bromine Kr Br krypton 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 85.5 87.6 88.9 91.2 92.9 95.9 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 Rb Sr Y Zr Nb Mo Tc Pd Ag Cd In Sn Sb Te I Xe 0915 - 13690 rubidium strontium yttrium zirconium Ru rhodium niobium molybdenum technetium ruthenium Rh palladium silver cadmium Indium tin antimony tellurium Iodine xenon 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 132.9 137.3 138.9 178.5 180.9 183.8 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 Cs caesium barium Ba lanthanum La* hafnium Ta tungsten Hf tantalum W rhenium Os Re osmium Ir iridium Pt platinum Au gold Hg Tl thallium Pb lead Bi bismuth Po polonium At astatine Rn radon mercury 55 56 57 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86 Elements with atomic numbers 112–116 have francium Ra Ac † Rf Fr radium Db Sg Bh Hs Mt Ds Rg been reported but not fully authenticated actinium rutherfordium dubnium seaborgium bohrium hassium meitnerium darmstadtium roentgenium 87 88 89 104 105 106 107 108 109 110 111 140 141 144 150 152 157 159 163 165 167 169 173 175 * Lanthanides Ce cerium Pr praseodymium Nd Pm Sm Eu Gd neodymium promethium samarium europium gadolinium Tb terbium Dy dysprosium Ho holmium Er erbium thulium Tm ytterbium Lu Yb lutetium 58 59 60 61 62 63 64 65 66 67 68 69 70 71 232 238 † Actinides Th thorium Pa protactinium U uranium Np Pu Am Cm neptunium plutonium americium Bk Cf Es Fm Md nobelium curium berkelium californium einsteinium fermium Lr No lawrencium mendelevium 90 91 92 93 94 95 96 97 98 99 100 101 102 103  A-Level Chemistry Exam Board: Edexcel Revising for Chemistry exams is stressful, that’s for sure — even just getting your notes sorted out can leave you needing a lie down. But help is at hand... This brilliant CGP book explains everything you’ll need to learn (and nothing you won’t), all in a straightforward style that’s easy to get your head around. We’ve also included exam questions to test how ready you are for the real thing. P G C A-Level revision? It has to be CGP! Published by CGP Editors: Katie Braid, Katherine Faudemer, Robin Flello, Emily Howe, Paul Jordin, Sophie Scott and Ben Train. Contributors: Mike Bossart, Rob Clarke, Ian H. Davis, John Duffy, Max Fishel, Emma Grimwood, Paddy Gannon, Lucy Muncaster, Jane Simoni, Derek Swain, Paul Warren and Christopher Workman. CER72DK With thanks to Jamie Sinclair for the proofreading. With thanks to Jan Greenway for the copyright research. Cover Photo © Laguna Design/Science Photo Library Clipart from Corel® Based on the classic CGP style created by Richard Parsons. Text, design, layout and original illustrations © Coordination Group Publications Ltd. (CGP) 2015 All rights reserved. 0800 1712 712 www.cgpbooks.co.uk  Contents If you’re revising for the AS exams, you’ll need Topics 1 – 10, and the Practical Skills section at the back. If you’re revising for the A-Level exams, you’ll need the whole book. The Scientific Process AS Topic 3 — Redox I AS The Scientific Process 2 Oxidation Numbers 38 Redox Reactions 40 Topic 1 — Atomic Structure AS and the Periodic Table Topic 4 — Inorganic Chemistry AS The Atom 4 and the Periodic Table Relative Mass 6 Group 2 42 More on Relative Mass 8 Group 1 and 2 Compounds 44 Electronic Structure 10 Halogens 46 Atomic Emission Spectra 12 Reactions of Halogens 48 Ionisation Energies 14 Reactions of Halides 50 Periodicity 16 Tests for Ions 52 Topic 2 — Bonding AS Topic 5 — Formulae, Equations AS and Structure & Amounts of Substances Ionic Bonding 19 The Mole 54 Covalent Bonding 22 Empirical and Molecular Formulae 56 Shapes of Molecules 24 Chemical Equations 58 Giant Covalent and Metallic Structures 26 Calculations with Gases 60 Electronegativity and Polarisation 28 Acid-Base Titrations 62 Intermolecular Forces 30 Titration Calculations 64 Hydrogen Bonding 32 Uncertainty and Errors 66 Solubility 34 Atom Economy and Percentage Yield 68 Predicting Structures and Properties 36 Topic 6 — Organic Chemistry I AS Topic 10 — Equilibrium I AS The Basics 70 Dynamic Equilibrium 118 Organic Reactions 72 Le Chatelier’s Principle 120 Isomerism 74 Alkanes 76 Crude Oil 78 Fuels 80 Topic 11 — Equilibrium II Alkenes 82 Calculations Involving Kc 122 Stereoisomerism 83 Gas Equilibria 124 Reactions of Alkenes 86 Le Chatelier’s Principle Polymers 88 and Equilibrium Constants 126 Halogenoalkanes 90 More on Halogenoalkanes 92 Alcohols 94 Oxidation of Alcohols 96 Topic 12 — Acid-Base Equilibria Organic Techniques 98 Acids and Bases 128 pH 130 The Ionic Product of Water 132 Topic 7 — Modern AS Experiments Involving pH 134 Analytical Techniques I Titration Curves and Indicators 136 Mass Spectrometry 100 Buffers 139 Infrared Spectroscopy 102 Topic 8 — Energetics I AS Topic 13 — Energetics II Enthalpy Changes 104 Lattice Energy 142 More on Enthalpy Changes 106 Polarisation 144 Hess’s Law 108 Dissolving 146 Bond Enthalpy 110 Entropy 148 More on Entropy Change 150 Free Energy 152 Topic 9 — Kinetics I AS Collision Theory 112 Reaction Rates 114 Catalysts 116 Topic 14 — Redox II Topic 18 — Organic Chemistry III Electrochemical Cells 154 Aromatic Compounds 205 Electrode Potentials 156 Electrophilic Substitution Reactions 208 The Electrochemical Series 158 Phenols 211 Storage and Fuel Cells 160 Amines 212 Redox Titrations 162 Amides 215 More on Redox Titrations 165 Condensation Polymers 216 Amino Acids 218 Grignard Reagents 220 Topic 15 — Transition Metals Organic Synthesis 221 Practical Techniques 224 Transition Metals 168 More Practical Techniques 226 Complex Ions 170 Empirical and Molecular Formulae 228 Complex Ions and Colour 172 Chromium 174 Reactions of Ligands 176 Transition Metals and Catalysis 178 Topic 19 — Modern Analytical Techniques II High Resolution Mass Spectrometry 230 Topic 16 — Kinetics II NMR Spectroscopy 231 Reaction Rates 180 Proton NMR Spectroscopy 234 Orders of Reactions 182 Chromatography 236 The Initial Rates Method 184 Combined Techniques 238 Rate Equations 186 The Rate-Determining Step 188 Halogenoalkanes and Reaction Mechanisms 190 Practical Skills AS Activation Energy 192 Planning Experiments 240 Practical Techniques 242 Presenting Results 244 Topic 17 — Organic Chemistry II Analysing Results 246 Evaluating Experiments 248 Optical Isomerism 194 Aldehydes and Ketones 196 Reactions of Aldehydes and Ketones 198 Do Well In Your Exams 250 Carboxylic acids 200 Answers 252 Esters 202 Index 273 Acyl Chlorides 204 2 The Scientific Process The Scientific Process These pages are all about the scientific process — how we develop and test scientific ideas. It’s what scientists do all day, every day (well except at coffee time — never come between scientists and their coffee). Scientists Come Up with Theories — Then Test Them... Science tries to explain how and why things happen. It’s all about seeking and gaining knowledge about the world around us. Scientists do this by asking questions and suggesting answers and then testing them, to see if they’re correct — this is the scientific process. 1) Ask a question — make an observation and ask why or how whatever you’ve observed happens. E.g. Why does sodium chloride dissolve in water? 2) Suggest an answer, or part of an answer, by forming a theory or a model (a possible explanation of the observations or a description of what you think is happening actually happening). E.g. Sodium chloride is made up of charged particles || | | | | | | | | | | | | | | | | | | | | || | which are pulled apart by the polar water molecules. A theory is only | | | | | | | scientific | | | | | | | 3) Make a prediction or hypothesis — a specific testable statement, if it can be tested. | | | | | | | | | | | | | | || based on the theory, about what will happen in a test situation. || | | | | | | | E.g. A solution of sodium chloride will conduct electricity much better than water does. 4) Carry out tests — to provide evidence that will support the prediction or refute it. E.g. Measure the conductivity of water and of sodium chloride solution....Then They Tell Everyone About Their Results... The results are published — scientists need to let others know about their work. Scientists publish their results in scientific journals. These are just like normal magazines, only they contain scientific reports (called papers) instead of the latest celebrity gossip. 1) Scientific reports are similar to the lab write-ups you do in school. And just as a lab write-up is reviewed (marked) by your teacher, reports in scientific journals undergo peer review before they’re published. Scientists use standard terminology when writing their reports. This way they know that other scientists will understand them. For instance, there are internationally agreed rules for naming organic compounds, so that scientists across the world will know exactly what substance is being referred to. See page 70. 2) The report is sent out to peers — other scientists who are experts in the same area. They go through it bit by bit, examining the methods and data, and checking it’s all clear and logical. When the report is approved, it’s published. This makes sure that work published in scientific journals is of a good standard. 3) But peer review can’t guarantee the science is correct — other scientists still need to reproduce it. 4) Sometimes mistakes are made and bad work is published. Peer review isn’t perfect but it’s probably the best way for scientists to self-regulate their work and to publish quality reports....Then Other Scientists Will Test the Theory Too 1) Other scientists read the published theories and results, and try to test the theory themselves. This involves: Repeating the exact same experiments. Using the theory to make new predictions and then testing them with new experiments. 2) If all the experiments in the world provide evidence to back it up, the theory is thought of as scientific ‘fact’. 3) If new evidence comes to light that conflicts with the current evidence the theory is questioned all over again. More rounds of testing will be carried out to try to find out where the theory falls down. This is how the scientific process works — evidence supports a theory, loads of other scientists read it and test it for themselves, eventually all the scientists in the world agree with it and then bingo, you get to learn it. When looking at experiments that give conflicting results, it’s important to look at all the evidence to work out whether a theory is supported or not — this includes looking at the methodology (the techniques) used in the experiments and the data collected. This is how scientists arrived at the structure of the atom (see page 4) — and how they came to the conclusion that electrons are arranged in shells and orbitals. As is often the case, it took years and years for these models to be developed and accepted. The Scientific Process  3 The 1862_RG_MainHead Scientific Process If the Evidence Supports a Theory, It’s Accepted — for Now Our currently accepted theories have survived this ‘trial by evidence’. They’ve been tested over and over again and each time the results have backed them up. BUT, and this is a big but (teehee), they never become totally indisputable fact. Scientific breakthroughs or advances could provide new ways to question and test the theory, which could lead to changes and challenges to it. Then the testing starts all over again... And this, my friend, is the tentative nature of scientific knowledge — it’s always changing and evolving. Evidence Comes From Lab Experiments... || | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | || | | | | For example, if you’re investigating how | | | | | | | | | | | | 1) Results from controlled experiments in laboratories are great. temperature affects the rate of a reaction, you | | | | | | | | | | 2) A lab is the easiest place to control variables so that they’re need to keep everything but the temperature all kept constant (except for the one you’re investigating). constant, e.g. the pH of the solution, the 3) This means you can draw meaningful conclusions. concentration of the solution, etc. || || | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | |...But You Can’t Always do a Lab Experiment There are things you can’t study in a lab. And outside the lab, controlling the variables is tricky, if not impossible. Are increasing CO2 emissions causing climate change? There are other variables which may have an effect, such as changes in solar activity. You can’t easily rule out every possibility. Also, climate change is a very gradual process. Scientists won’t be able to tell if their predictions are correct for donkey’s years. Does drinking chlorinated tap water increase the risk of developing certain cancers? There are always differences between groups of people. The best you can do is to have a well-designed study using matched groups — choose two groups of people (those who drink tap water and those who don’t) which are as similar as possible Samantha thought her (same mix of ages, same mix of diets, etc). But you still can’t rule out every possibility. study was very well designed — especially Taking newborn identical twins and treating them identically, except for the fitted bookshelf. making one drink gallons of tap water and the other only pure water, might be a fairer test, but it would present huge ethical problems. Science Helps to Inform Decision-Making Lots of scientific work eventually leads to important discoveries that could benefit humankind — but there are often risks attached (and almost always financial costs). Society (that’s you, me and everyone else) must weigh up the information in order to make decisions — about the way we live, what we eat, what we drive, and so on. Information is also used by politicians to devise policies and laws. Chlorine is added to water in small quantities to disinfect it (see page 49). Some studies link drinking chlorinated water with certain types of cancer. But the risks from drinking water contaminated by nasty bacteria are far, far greater. There are other ways to get rid of bacteria in water, but they’re heaps more expensive. Scientific advances mean that non-polluting hydrogen-fuelled cars can be made. They’re better for the environment, but are really expensive. And it’d cost a lot to adapt filling stations to store hydrogen. Pharmaceutical drugs are really expensive to develop, and drug companies want to make money. So they put most of their efforts into developing drugs that they can sell for a good price. Society has to consider the cost of buying new drugs — the NHS can’t afford the most expensive drugs without sacrificing something else. So there you have it — how science works... Hopefully these pages have given you a nice intro to how science works. You need to understand it for the exam, and for life. Once you’ve got it sussed it’s time to move on to the really good stuff — the chemistry. Bet you can’t wait... The Scientific Process 4 Topic 1 — Atomic Structure and the Periodic Table The Atom This stuff about atoms and elements should be ingrained in your brain from GCSE. You do need to know it perfectly though if you are to negotiate your way through the field of man-eating tigers and pesky atoms... Atoms are made up of Protons, Neutrons and Electrons Atoms are the stuff all elements and compounds are made of. They’re made up of 3 types of subatomic particle — protons, neutrons and electrons. Electrons Nucleus 1) Electrons have –1 charge. 1) Most of the mass of the atom is 2) They whizz around the p n concentrated in the nucleus. n n p p n nucleus in orbitals. 2) The diameter of the nucleus is rather The orbitals take up most titchy compared to the whole atom. of the volume of the atom. 3) The nucleus is where you find the protons and neutrons. The mass and charge of these subatomic particles are tiny, so relative mass and relative charge are used instead. | || | | | | | | | | | | | | | | | Subatomic particle Relative mass Relative charge | | | | | | | | | | | | The mass of an electron is || | | | | | | | | | | || | | | | | | | | | Proton 1 +1 negligible compared to a | Neutron 1 0 proton or a neutron — this means you can usually ign Electron, e – 0.0005 –1 ore it. | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | Nuclear Symbols Show Numbers of Subatomic Particles You can figure out the number of protons, neutrons and electrons from the nuclear symbol, which is found in the periodic table. || | | | | | | | | | | | | | | | | Mass number || A | | | | | | | | | | | X Sometimes the atomic number is | | | | | | | | | | || || This tells you the total number of Element symbol | | | | | | | | | | | | left out of the nuclear symbol, protons and neutrons in the nucleus. e.g. 7Li. You don’t really need Z it because the element’s symbol tells you its value. Atomic (proton) number || | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | 1) This is the number of protons in the nucleus — it identifies the element. 2) All atoms of the same element have the same number of protons. 1) For neutral atoms, which have no overall charge, the number of electrons is the same as the number of protons. 2) The number of neutrons is just mass number minus atomic number, | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | To work out the number of each | | | | | | | | || | | | | | | | | | | | i.e. ‘top minus bottom’ in the nuclear symbol. subatomi c particle present in a molecule, just work out how many there are in Nuclear Atomic Mass Protons Electrons Neutrons each atom and then add them all up. symbol number, Z number, A || | | | || | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | 7 3 Li 3 7 3 3 7–3=4 19 9F 9 19 9 9 19 – 9 = 10 24 12 Mg 12 24 12 12 24 – 12 = 12 Ions have Different Numbers of Protons and Electrons “Hello, I’m Newt Ron...” Negative ions have more electrons than protons......and positive ions have fewer electrons than protons. The negative charge means that there’s The 2+ charge means that there are F– 1 more electron than there are protons. 2 fewer electrons than there are protons. F has 9 protons (see table above), so F– must have 10 electrons. Mg2+ Mg has 12 protons (see table above), so Mg2+ must have 10 electrons. The overall charge = +9 – 10 = –1. The overall charge = +12 – 10 = +2. Topic 1 — Atomic Structure and the Periodic Table  5 The Atom Isotopes are Atoms of the Same Element with Different Numbers of Neutrons Isotopes of an element are atoms with the same number of protons but different numbers of neutrons. Chlorine-35 and chlorine-37 are examples of isotopes: 35 – 17 = 18 neutrons Different mass numbers mean different 37 – 17 = 20 neutrons masses and different numbers of neutrons. 35 17 Cl The atomic numbers are the same. Both isotopes have 17 protons and 17 electrons. 37 17 Cl 1) It’s the number and arrangement of electrons that decides the chemical properties of an element. Isotopes have the same configuration of electrons (see pages 10-11), so they’ve got the same chemical properties. 2) Isotopes of an element do have slightly different physical properties though, such as different densities, rates of diffusion, etc. This is because physical properties tend to depend more on the mass of the atom. Here’s another example — naturally | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | || occurring magnesium consists of 3 isotopes. The periodic table gives the atomic number | | | | | | | | | | | | | || | | | | | | | | | | | | | | | | for each element. The other number isn’t the 24 Mg (79%) 25 Mg (10%) 26 Mg (11%) mass number — it’s the relative atomic mass 12 protons 12 protons 12 protons (see page 6). They’re a bit different , but you 12 neutrons 13 neutrons 14 neutrons can often assume they’re equal — it doesn’t matter unless you’re doing really accurate work. 12 electrons 12 electrons 12 electrons || | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | Practice Questions Q1 Draw a diagram showing the structure of an atom, labelling each part. Q2 Where is the mass concentrated in an atom, and what makes up most of the volume of an atom? Q3 Draw a table showing the relative charge and relative mass of the three subatomic particles found in atoms. Q4 Using an example, explain the terms ‘atomic number’ and ‘mass number’. Exam Questions Q1 Hydrogen, deuterium and tritium are all isotopes of each other. a) Identify one similarity and one difference between these isotopes. [2 marks] b) Deuterium can be written as 21 H. Determine the number of protons, neutrons and electrons in a deuterium atom. [1 mark] c) Write the nuclear symbol for tritium, given that it has 2 neutrons. [1 mark] 32 2- 40 30 42 Q2 This question relates to the atoms or ions A to D: A 16 S B 18Ar C 16 S D 20 Ca a) Identify the similarity for each of the following pairs, justifying your answer in each case. i) A and B. [1 mark] ii) A and C. [1 mark] iii) B and D. [1 mark] b) Which two of the atoms or ions are isotopes of each other? Explain your reasoning. [2 marks] 1 16 12 Q3 A molecule of propanol, C3H7OH, is made up of 1H , 8O and 6 C atoms. Calculate the number of electrons, protons and neutrons in one molecule of propanol. [2 marks] Got it learned yet? — Isotope so... This is a nice page to ease you into things. Remember that positive ions have fewer electrons than protons, and negative ions have more electrons than protons. Get that straight in your mind or you’ll end up in a right mess. Topic 1 — Atomic Structure and the Periodic Table 6 Relative Mass Relative mass... What? Eh?... Read on... Relative Masses are Masses of Atoms Compared to Carbon-12 The actual mass of an atom is very, very tiny. Don’t worry about exactly how tiny for now, but it’s far too small to weigh with a normal pair of scales in your classroom. So, the mass of one atom is compared to the mass of a different atom. This is its relative mass. Here are some definitions for you to learn: The relative atomic mass, Ar, 1) Relative atomic mass is an average is the weighted mean mass of all the relative isotopic masses, of an atom of an element, so it’s not usually a whole number. compared to 1/12th of the 2) Relative isotopic mass is usually mass of an atom of carbon-12. a whole number. E.g. a natural sample of chlorine Relative isotopic mass is the contains a mixture of 35Cl (75%) mass of an atom of an isotope, and 37Cl (25%), so the relative compared with 1/12th of the isotopic masses are 35 and 37. mass of an atom of carbon-12. But its relative atomic mass is 35.5. Jason’s shirt was isotropical... Relative Molecular Masses are Masses of Molecules The relative molecular mass (or relative formula mass), Mr, is the average mass of a molecule or formula unit, compared to 1/12th of the mass of an atom of carbon-12. Don’t worry, this is one definition that you don’t need to know for the exam. But... you do need to know how to work out the relative molecular mass, and the relative formula mass, so it’s probably best if you learn what they mean anyway. 1) Relative molecular mass is used when 1) Relative formula mass is used for compounds referring to simple molecules. that are ionic (or giant covalent, such as SiO2). 2) To find the relative molecular mass, just add up the relative 2) To find the relative formula mass, add up the atomic mass values of all the atoms in the molecule. relative atomic masses (Ar) of all the ions or atoms in the formula unit. (Ar of ion = Ar of atom. E.g. Mr(C2H6O) = (2 × 12.0) + (6 × 1.0) + 16.0 = 46.0 The electrons make no difference to the mass.) || | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | || | | | | || | | | | | || See page 22 for more on simple molecules, and E.g. Mr(CaF2) = 40.1 + (2 × 19.0) = 78.1 pages 20 and 26-27 for more on giant structures. | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | || Ar Can Be Worked Out from Isotopic Abundances You need to know how to calculate the relative atomic mass (Ar) of an element from its isotopic abundances. 1) Different isotopes of an element occur in different quantities, or isotopic abundances. 2) To work out the relative atomic mass of an element, you need to work out the average mass of all its atoms. 3) If you’re given the isotopic abundances in percentages, all you need to do is follow these two easy steps: Step 1: Multiply each relative isotopic mass by its % relative isotopic abundance, and add up the results. Step 2: Divide by 100. Example: Find the relative atomic mass of boron, given that 20.0% of the boron atoms found on Earth have a relative isotopic mass of 10.0, while 80.0% have a relative isotopic mass of 11.0. Step 1: (20.0 × 10) + (80.0 × 11) = 1080 Step 2: 1080 ÷ 100 = 10.8 Topic 1 — Atomic Structure and the Periodic Table  7 Relative Mass Mass Spectrometry Can Tell Us About Isotopes Mass spectra are produced by mass spectrometers — devices which are used to find out what samples are made up of by measuring the masses of their components. Mass spectra can tell us dead useful things, e.g. the relative isotopic masses and abundances of different elements. Mass spectra can be used to work out the relative atomic masses of different elements. % abundance Cl isotopes This is the mass spectra for chlorine. 100 | || | | | | | | | | | | | | | | | | | | | | | | | This spectrum show || | | | | | | | | | The y-axis gives the abundance of s that | | | | | | | | | | || 80 75.5% chlorine exists as 2 ions, often as a percentage. For an isotopes. 75.5% of chlorine is 35 element, the height of each peak 60 Cl, and 24.5% is 37Cl. | gives the relative isotopic abundance. 40 | | | | | | | | | | | | | | | | | | | | | | | | | | 24.5% The x-axis units are given as a ‘m/z’ value, 20 which is a mass/charge ratio. Since the charge on the ions is mostly +1, you can often assume 0 34 35 36 37 38 the x-axis is simply the relative isotopic mass. m/z The method for working out the relative atomic mass from a graph is a bit different to working it out from percentages (see previous page), but it starts off in the same way. Step 1: Multiply each relative isotopic mass by its Mass Spectrum of Ne relative isotopic abundance, and add up the results. 114.0 Relative abundance Step 2: Divide by the sum of the isotopic abundances. Example: Use the data from this mass spectrum to work out the relative atomic mass of neon. Give your answer to 1 decimal place. Step 1: (20 × 114.0) + (21 × 0.2) + (22 × 11.2) = 2530.6 11.2 Step 2: (114.0 + 0.2 + 11.2 = 125.4) 0.2 2530.6 ÷ 125.4 = 20.2 20 21 22 23 m/z Practice Questions Q1 Explain what relative atomic mass (Ar) and relative isotopic mass mean. Q2 Explain the difference between relative molecular mass and relative formula mass. Q3 Explain what relative isotopic abundance means. Exam Questions Q1 Copper exists in two main isotopic forms, 63Cu and 65Cu. Mass Spectrum of Cu Relative abundance a) Calculate the relative atomic mass of copper using the information 120.8 from the mass spectrum. [2 marks] b) Explain why the relative atomic mass of copper is not a whole number. [2 marks] 54.0 Q2 The percentage make-up of naturally occurring potassium is: 93.1% 39K, 0.120% 40K and 6.77% 41K. 61 63 65 67 Use the information to determine the relative atomic mass of potassium. [2 marks] m/z You can’t pick your relatives, you just have to learn them... Isotopic masses are a bit frustrating. Why can’t all atoms of an element just be the same? But the fact is they’re not, so you’re going to have to learn how to use those spectra to work out the relative atomic masses of different elements. The actual maths is pretty simple. A pinch of multiplying, a dash of addition, some division to flavour and you’re away. Topic 1 — Atomic Structure and the Periodic Table 8 More on Relative Mass “More relative mass?! How much more could there possibly be?” I hear you cry. Well, as you’re about to see, there’s plenty more. This is all dead useful to scientists and (more importantly) to you in your exams. You Can Calculate Isotopic Masses from Relative Atomic Mass If you know the relative atomic mass of an element, and you know all but one of the abundances and relative isotopic masses of its isotopes, you can work out the abundance and isotopic mass of the final isotope. Neat huh? Example: Silicon can exist in three isotopes. 92.23% of silicon is 28Si and 4.67% of silicon is 29Si. Given that the Ar of silicon is 28.1, calculate the abundance and isotopic mass of the third isotope. Step 1: First, find the abundance of the third isotope. You’re dealing with percentage abundances, so you know they need to total 100%. So, the abundance of the final isotope will be 100% – 92.23% – 4.67% = 3.10% Step 2: You know that the relative atomic mass (Ar) of silicon is 28.1, and you know two of the three isotopic masses. So, you can put all of that into the equation you use to work out the relative atomic mass from relative abundances and isotopic masses (see page 6), which you can then rearrange to work out the final isotopic mass, X. || || | | | | | | | | | | | | | | | | | | | | | Remember — isotopic || | | | | | | | | | | | | 28.1 = ((28 × 92.23) + (29 × 4.67) + (X × 3.10)) ÷ 100 | | | | | | | | | | | | | masses are usually whole 28.1 = (2717.87 + (X × 3.10)) ÷ 100 numbers, so you should 2810 – 2717.87 = X × 3.10 round your answer to the nearest whole number. 29.719 = X So the isotopic mass of the third isotope is 30 — 30Si. || | | | | | | | | | | | | | | | | | | | | | | | | You Can Predict the Mass Spectra for Diatomic Molecules Now, this is where it gets even more mathsy and interesting (seriously — I love it). You can use your knowledge to predict what the mass spectra of diatomic molecules (i.e. molecules containing two atoms) look like. Example: Chlorine has two isotopes. 35Cl has an abundance of 75% | || | | | | | | | | | | | | | and 37Cl has an abundance of 25%. Predict the mass spectrum of Cl2. || | | | | | | | | || To convert a || | | | | | | | | | percentage to 1) First, express each of the percentages as a decimal: 75% = 0.75 and 25% = 0.25. a decimal, just divide by 100. 2) Make a table showing all the 35 Cl 37 Cl || | | | | | | | | | | | | || different Cl2 molecules. For each 35 Cl – 35Cl: 0.75 × 0.75 35 Cl – 37Cl: 0.25 × 0.75 35 Cl molecule, multiply the abundances = 0.5625 = 0.1875 (as decimals) of the isotopes to get 37 Cl – 35Cl: 0.25 × 0.75 37 Cl – 37Cl: 0.25 × 0.25 the relative abundance of each one. 37 Cl = 0.1875 = 0.0625 3) Look for any molecules in the table that are the same and add up their abundances. In this case, 37Cl–35Cl and 35Cl–37Cl are the same, so the actual abundance for this molecule is: 0.1875 + 0.1875 = 0.375. 4) Divide all the relative abundances by the smallest relative abundance to get the smallest whole number ratio. And by working out the relative molecular Mass Spectrum of Cl2 mass of each molecule, you can predict the mass spectrum for Cl2: relative abundance 9 Relative Molecule Relative abundance Molecular Mass 6 35 Cl – 35Cl 35 + 35 = 70 0.5625 ÷ 0.0625 = 9 1 35 Cl – 37Cl 35 + 37 = 72 0.375 ÷ 0.0625 = 6 68 70 72 74 76 m/z 37 Cl – Cl 37 37 + 37 = 74 0.0625 ÷ 0.0625 = 1 Topic 1 — Atomic Structure and the Periodic Table  9 More on Relative Mass Mass Spectrometry Can Also Help to Identify Compounds 1) You’ve seen how you can use a mass spectrum showing the relative isotopic abundances of an element to work out its relative atomic mass. You need to make sure you can remember how to do this. You can also get mass spectra for molecules made up from more than one element. || | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | || | | | | | Assuming the ion has a 1+ charge, | | | | | | | 2) When the molecules in a sample are bombarded with electrons, which it normally will have. an electron is removed from the molecule to form a molecular ion, M+(g). || | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | | 3) To find the relative molecular mass of a compound, you look at the molecular ion peak (the M peak) on the mass spectrum. This is the peak with the highest m/z value (ignoring any small M+1 peaks that occur due to the presence of any atoms of carbon-13). The mass/charge value of the molecular ion peak is the molecular mass. Th