Preview Chemistry An Introduction to Organic, Inorganic Physical Chemistry by Catherine E. Housecroft, Edwin C. Constable (2006)

Preview Chemistry An Introduction to Organic, Inorganic Physical Chemistry by Catherine E. Housecroft, Edwin C. Constable (2006) Preview Chemistry An Introduction to Organic, Inorganic Physical Chemistry by Catherine E. Housecroft, Edwin C. Constable (2006) Preview Chemistry An Introduction to Organic, Inorganic Physical Chemistry by Catherine E. Housecroft, Edwin C. Constable (2006) Preview Chemistry An Introduction to Organic, Inorganic Physical Chemistry by Catherine E. Housecroft, Edwin C. Constable (2006)

Housecroft CvrAW 01/07/2005 8:34 Page 3rd Edition Provides robust coverage of the different branches of chemistry – with unique depth in organic chemistry in an introductory text – helping students to develop a solid understanding of chemical principles, how they interconnect, and how they can be applied to our lives KEY FEATURES • Chapters are grouped into inorganic, physical and organic components, with integrating themes highlighted throughout This provides flexibility and highlights the connections between different areas • Topic boxes throughout the book highlight the applications of chemistry in industry, biology, the environment and the laboratory These help students to understand the relevance of chemistry to everyday life • The text is four-colour throughout with three-dimensional computer-generated artwork, and is supported by CHIME graphics on the companion website, helping students to visualise chemical structures • Definition boxes and end-of-chapter checklists provide excellent revision aids • End-of-chapter problems reinforce learning and develop subject knowledge; in this edition, answers to non-descriptive problems have been added at the end of the book • A companion website at www.pearsoned.co.uk/housecroft features multiple-choice questions, rotatable three-dimensional molecular structures, and a Mathematics Tutor For full information and details of the OneKey resource available with this book, see prelim page xx Catherine E Housecroft and Edwin C Constable are both Professors of Chemistry at the University of Basel, Switzerland They have extensive international teaching experience and their research interests include supramolecular chemistry, nanotechnology, organometallic and cluster chemistry In 1997 Professor Constable was awarded the prestigious Howard lectureship for pre-eminence in organic chemistry or related disciplines by the University of Sydney, Australia CHEMISTRY CHEMISTRY Catherine E Housecroft and Edwin C Constable Catherine E Housecroft and Edwin C Constable CHEMISTRY 3rd Edition 3rd Edition Cover illustration by Gary Thompson www.pearson-books.com Catherine E Housecroft Edwin C Constable Visit the Chemistry, Third Edition Companion Website at www.pearsoned.co.uk/housecroft to find valuable student learning material including: Multiple choice questions to help test your learning Rotatable 3D structures taken from the book Mathematics Tutor Annotated links to relevant sites on the web CHEMISTRY An Introduction to Organic, Inorganic and Physical Chemistry 3rd edition Catherine E Housecroft Edwin C Constable Pearson Education Limited Edinburgh Gate Harlow Essex CM20 2JE England and Associated Companies throughout the world Visit us on the World Wide Web at: www.pearsoned.co.uk First published under the Longman imprint 1997 Second edition 2002 Third edition 2006 # Pearson Education Limited 1997, 2006 The rights of Catherine E Housecroft and Edwin C Constable to be identified as authors of this work have been asserted by them in accordance with the Copyright, Designs and Patents Act 1988 All rights reserved No part of this publication may be reproduced, stored in a retrieval system, or transmitted in any form or by any means, electronic, mechanical, photocopying, recording or otherwise, without either the prior written permission of the publisher or a licence permitting restricted copying in the United Kingdom issued by the Copyright Licensing Agency Ltd, 90 Tottenham Court Road, London W1T 4LP All trademarks used herein are the property of their respective owners The use of any trademark in this text does not vest in the author or publisher any trademark ownership rights in such trademarks, nor does the use of such trademarks imply any affiliation with or endorsement of this book by such owners ISBN 131 27567 British Library Cataloguing-in-Publication Data A catalogue record for this book is available from the British Library Library of Congress Cataloging-in-Publication Data A catalog record for this book is available from the Library of Congress 10 09 08 07 06 05 Typeset in 10pt Times by 60 Printed by Ashford Colour Press Ltd, Gosport The publisher’s policy is to use paper manufactured from sustainable forests Summary of contents Preface About the authors Acknowledgements xxi xxiii xxviii Some basic concepts Thermochemistry 56 Atoms and atomic structure 79 Homonuclear covalent bonds 118 Heteronuclear diatomic molecules 167 Polyatomic molecules: shapes 200 Polyatomic molecules: bonding 236 Ions 263 Elements 301 10 Mass spectrometry 330 11 Introduction to spectroscopy 344 12 Vibrational and rotational spectroscopies 356 13 Electronic spectroscopy 387 14 NMR spectroscopy 401 15 Reaction kinetics 428 16 Equilibria 487 17 Thermodynamics 530 18 Electrochemistry 575 19 The conductivity of ions in solution 597 20 Periodicity 609 21 Hydrogen and the s-block elements 623 22 p-Block and high oxidation state d-block elements 668 23 Coordination complexes of the d-block metals 737 24 Carbon compounds: an introduction 792 25 Acyclic and cyclic alkanes 831 vi Summary of contents 26 Alkenes and alkynes 839 27 Polar organic molecules: an introduction 904 28 Halogenoalkanes 911 29 Ethers 943 30 Alcohols 961 31 Amines 987 32 Aromatic compounds 1007 33 Carbonyl compounds 1060 34 Aromatic heterocyclic compounds 1111 35 Molecules in nature 1150 Appendices 1184 Answers to non-descriptive problems 1221 Index 1235 Contents Preface About the authors Acknowledgements 1.1 1.2 1.3 1.4 1.5 1.6 1.7 1.8 1.9 1.10 1.11 1.12 1.13 1.14 1.15 1.16 1.17 1.18 1.19 2.1 2.2 2.3 2.4 2.5 2.6 2.7 2.8 2.9 2.10 Some basic concepts xxi xxiii xxviii What is chemistry and why is it important? What is the IUPAC? SI units The proton, electron and neutron The elements States of matter Atoms and isotopes The mole and the Avogadro constant Gas laws and ideal gases The periodic table Radicals and ions Molecules and compounds: bond formation Molecules and compounds: relative molecular mass and moles Concentrations of solutions Reaction stoichiometry Oxidation and reduction, and oxidation states Empirical, molecular and structural formulae Basic nomenclature Final comments Problems 3 11 14 14 24 26 28 29 30 32 39 45 48 52 53 Thermochemistry 56 Factors that control reactions Change in enthalpy of a reaction Measuring changes in enthalpy: calorimetry Standard enthalpy of formation Calculating standard enthalpies of reaction Enthalpies of combustion Hess’s Law of Constant Heat Summation Thermodynamic and kinetic stability Phase changes: enthalpies of fusion and vaporization An introduction to intermolecular interactions Summary Problems 56 57 58 63 64 66 68 71 71 74 76 76 viii Contents 3.1 3.2 3.3 3.4 3.5 3.6 3.7 3.8 3.9 3.10 3.11 3.12 3.13 3.14 3.15 3.16 3.17 3.18 3.19 3.20 3.21 4.1 4.2 4.3 4.4 4.5 4.6 4.7 4.8 4.9 4.10 4.11 4.12 4.13 4.14 4.15 4.16 4.17 4.18 Atoms and atomic structure 79 The importance of electrons The classical approach to atomic structure The Bohr atom – still a classical picture Quanta Wave–particle duality The uncertainty principle The Schroădinger wave equation Probability density The radial distribution function, 4r2 RðrÞ2 Quantum numbers Atomic orbitals Relating orbital types to the principal quantum number More about radial distribution functions Applying the Schroădinger equation to the hydrogen atom Penetration and shielding The atomic spectrum of hydrogen and selection rules Many-electron atoms The aufbau principle Electronic configurations The octet rule Monatomic gases Summary Problems 79 80 81 82 82 83 84 86 87 88 90 94 94 96 98 100 106 106 108 111 112 116 117 Homonuclear covalent bonds 118 Introduction Measuring internuclear distances The covalent radius of an atom An introduction to bond energy: the formation of the diatomic molecule H2 Bond energies and enthalpies The standard enthalpy of atomization of an element Determining bond enthalpies from standard heats of formation The nature of the covalent bond in H2 Lewis structure of H2 The problem of describing electrons in molecules Valence bond (VB) theory Molecular orbital (MO) theory What the VB and MO theories tell us about the molecular properties of H2 ? Homonuclear diatomic molecules of the first row elements – the s-block Orbital overlap of p atomic orbitals Bond order Relationships between bond order, bond length and bond enthalpy Homonuclear diatomic molecules of the first row p-block elements: F2 and O2 118 119 123 125 127 129 133 134 135 136 137 139 144 144 146 148 149 149 Contents 4.19 4.20 4.21 4.22 4.23 5.1 5.2 5.3 5.4 5.5 5.6 5.7 5.8 5.9 5.10 5.11 5.12 5.13 5.14 5.15 5.16 6.1 6.2 6.3 6.4 6.5 6.6 6.7 6.8 6.9 6.10 6.11 6.12 6.13 ix Orbital mixing and – crossover Homonuclear diatomic molecules of the first row p-block elements: B2 , C2 and N2 Periodic trends in the homonuclear diatomic molecules of the first row elements The diatomic species O2 , [O2 ]ỵ , [O2 ]ÿ and [O2 ]2ÿ Group trends among homonuclear diatomic molecules Summary Problems 152 158 159 160 164 165 Heteronuclear diatomic molecules 167 Introduction Lewis structures for HF, LiF and LiH The valence bond approach to the bonding in HF, LiF and LiH The molecular orbital approach to the bonding in a heteronuclear diatomic molecule The molecular orbital approach to the bonding in LiH, LiF and HF Bond enthalpies of heteronuclear bonds Electronegativity – Pauling values ( P ) The dependence of electronegativity on oxidation state and bond order An overview of the bonding in HF Other electronegativity scales Polar diatomic molecules Isoelectronic species The bonding in CO by the Lewis and valence bond approaches The bonding in carbon monoxide by MO theory [CN] and [NO]ỵ : two ions isoelectronic with CO [NO]ỵ , NO and [NO] Summary Problems 167 168 168 170 172 177 179 181 182 182 183 186 187 189 193 196 197 197 Polyatomic molecules: shapes 200 Introduction The geometries of triatomic molecules Molecules larger than triatomics described as having linear or bent geometries Geometries of molecules within the p-block: the first row Heavier p-block elements Mid-chapter problems The valence-shell electron-pair repulsion (VSEPR) model The VSEPR model: some ambiguities The Kepert model Application of the Kepert model An exception to the Kepert model: the square planar geometry Stereoisomerism Two structures that are close in energy: the trigonal bipyramid and square-based pyramid Shape and molecular dipole moments 200 203 154 205 206 208 211 212 218 220 220 222 222 226 228 Oxidation and reduction, and oxidation states 41 The oxidation state of a group metal in a compound is ỵ2 Metals from the d-block will usually have positive oxidation states (exceptions are some low oxidation state compounds – see Section 23.14) Added to these rules are the facts that most elements in groups 13, 14, 15 and 16 can have variable oxidation states In reality, it is essential to have a full picture of the structure of a compound before oxidation states can be assigned Worked example 1.31 Working out oxidation states What are the oxidation states of each element in the following: KI, FeCl3 , Na2 SO4 ? KI: FeCl3 : Na2 SO4 : The group metal is typically in oxidation state þ1 This is consistent with the iodine being in oxidation state ÿ1, and the sum of the oxidation states is Chlorine is usually in oxidation state ÿ1, and since there are Cl atoms, the oxidation state of the iron must be ỵ3 to give a neutral compound Of the three elements, S can have a variable oxidation state and so we should deal with this element last Na is in group and usually has an oxidation state of ỵ1 Oxygen is usually in an oxidation state of ÿ2 The oxidation state of the sulfur atom is determined by ensuring that the sum of the oxidation states is 0: Oxidation state of Naị ỵ Oxidation state of Sị ỵ Oxidation state of Oị ẳ ỵ2ị ỵ Oxidation state of Sị ỵ 8ị ẳ Oxidation state of S ẳ ỵ ẳ ỵ6 Worked example 1.32 Variable oxidation states of nitrogen Determine the oxidation state of N in each of the following species: NO, [NO2 ] , [NO]ỵ , [NO3 ]ÿ , NO2 , ONF Each of the compounds or ions contains oxygen; O is usually in oxidation state ÿ2 NO: This is a neutral compound, therefore: ðOxidation state of Nị ỵ Oxidation state of Oị ẳ Oxidation state of Nị ỵ 2ị ẳ Oxidation state of N ẳ ỵ2 [NO2 ] : The overall charge is 1, therefore: Oxidation state of Nị ỵ Oxidation state of Oị ẳ Oxidation state of Nị ỵ 4ị ẳ Oxidation state of N ẳ ỵ ẳ ỵ3 42 CHAPTER Some basic concepts [NO]ỵ : The overall charge is +1, therefore: ðOxidation state of Nị ỵ Oxidation state of Oị ẳ ỵ1 Oxidation state of Nị ỵ 2ị ẳ ỵ1 Oxidation state of N ẳ ỵ1 ỵ ẳ ỵ3 [NO3 ] : The overall charge is 1, therefore: Oxidation state of Nị ỵ Oxidation state of Oị ẳ Oxidation state of Nị ỵ 6ị ẳ Oxidation state of N ẳ ỵ ẳ ỵ5 NO2 : This is a neutral compound, therefore: Oxidation state of Nị ỵ  Oxidation state of O) ¼ ðOxidation state of Nị ỵ 4ị ẳ Oxidation state of N ẳ þ4 ONF: This is a neutral compound; in a compound, F is always in oxidation state ÿ1, therefore: ðOxidation state of Nị ỵ Oxidation state of Oị ỵOxidation state of Fị ẳ Oxidation state of Nị ỵ 2ị ỵ 1ị ẳ Oxidation state of N ẳ ỵ3 Changes in oxidation states An increase in the oxidation state of an atom of an element corresponds to an oxidation process; a decrease in the oxidation state of an atom of an element corresponds to a reduction process " Nomenclature for oxidation states: see Section 1.18 In equation 1.45, the change in oxidation state of the sulfur from to ÿ2 is reduction, and in reaction 1.46, the change in oxidation state of the zinc from to ỵ2 is oxidation In a reduction–oxidation (redox) reaction, the change in oxidation number for the oxidation and reduction steps must balance In reaction 1.47, iron is oxidized to iron(III) and chlorine is reduced to chloride ion The net increase in oxidation state of the iron for the stoichiometric reaction must balance the net decrease in oxidation state for the chlorine 2Fe(s) + 3Cl2(g) Oxidation state: 2FeCl3(s) (1.47) +3 −1 Change = +3 Oxidation Change = × (−1) = −3 Reduction In equation 1.48, Fe2ỵ ions are oxidized by [MnO4 ] ; at the same time, [MnO4 ]ÿ is reduced by Fe2ỵ ions The [MnO4 ] ion is the oxidizing agent, Oxidation and reduction, and oxidation states 43 and Fe2ỵ is the reducing agent In the reaction, the oxidation states of H and O remain as ỵ1 and 2, respectively The oxidation state of Mn decreases from ỵ7 to ỵ2, and the oxidation state of Fe increases from ỵ2 to ỵ3 The balanced equation shows that ve Fe2ỵ ions are involved in the reaction and therefore the overall changes in oxidation states balance as shown in equation 1.48 [MnO4]– + 5Fe2+ + 8H+ Mn2+ + 5Fe3+ + 4H2O +7 +2 +2 (1.48) +3 Change = –5 Reduction Change = × (+1) = +5 Oxidation Worked example 1.33 Reduction of iron(III) oxide When heated with carbon, Fe2 O3 is reduced to Fe metal: heat Fe2 O3 sị ỵ 3Csị 2Fesị þ 3COðgÞ " Identify the oxidation and reduction processes Show that the oxidation state changes in the reaction balance The oxidation state of O is ÿ2 C(s) and Fe(s) are elements in oxidation state Fe2 O3 is a neutral compound: Oxidation state of Fe) ỵ Oxidation state of Oị ẳ Oxidation state of Feị ẳ Oxidation state of Fe ẳ ỵ6 ẳ ỵ3 In CO, the oxidation state of C ẳ ỵ2 The oxidation process is C going to CO; the reduction process is Fe2 O3 going to Fe The oxidation state changes balance as shown below: Fe2O3 + 3C +3 heat 2Fe + 3CO +2 Change = × (–3) = –6 Reduction Change = × (+2) = +6 Oxidation Worked example 1.34 Reaction of sodium with water When sodium metal is added to water, the following reaction occurs: 2Nasị ỵ 2H2 Olị 2NaOHaqị ỵ H2 gị " State which species is being oxidized and which is being reduced Show that the oxidation state changes balance In Na(s), Na is in oxidation state In H2 (g), H is in oxidation state 44 CHAPTER Some basic concepts H2 O is a neutral compound: ð2  Oxidation state of HÞ ỵ Oxidation state of Oị ẳ Oxidation state of Hị ỵ 2ị ẳ Oxidation state of H ẳ ỵ2 ẳ ỵ1 NaOH is a neutral compound; the usual oxidation states of Na and O are þ1 and ÿ2, respectively: ðOxidation state of NaÞ þ ðOxidation state of Oị ỵ Oxidation state of Hị ẳ ỵ1ị ỵ 2ị ỵ Oxidation state of Hị ẳ Oxidation state of H ẳ ẳ ỵ1 Using these oxidation states, we can write the following equation and changes in oxidation states: 2Na + 2H2O 2NaOH + H2 +1 +1 +1 Change = × (+1) = +2 Oxidation No change Change = × (–1) = –2 Reduction The equation shows that Na is oxidized; H is reduced on going from H2 O to H2 , but remains in oxidation state ỵ1 on going from H2 O to NaOH The changes in oxidation state balance Worked example 1.35 Hydrogen peroxide as an oxidizing agent Hydrogen peroxide reacts with iodide ions in the presence of acid according to the following equation: 2I ỵ H2 O2 þ 2Hþ ÿÿ I2 þ 2H2 O " Show that the changes in oxidation states in the equation balance Confirm that H2 O2 acts as an oxidizing agent The structure of H2 O2 is shown in the margin The molecule contains an OÿO bond (i.e a bond between like atoms – a homonuclear bond) This bond has no influence on the oxidation state of O H2 O2 is a neutral compound, therefore: Oxidation state of Oị ỵ Oxidation state of Hị ẳ Oxidation state of Oị ỵ 2ỵ1ị ẳ Oxidation state of O ¼ ÿ2 ¼ ÿ1 In H2 O, the oxidation states of H and O are ỵ1 and 2, respectively In Hỵ , hydrogen is in oxidation state ỵ1 In Iÿ and I2 , iodine is oxidation states ÿ1 and 0, respectively The oxidation state changes in the reaction are shown below; H remains in oxidation state +1 and does not undergo a redox reaction Empirical, molecular and structural formulae 2I– + H2O2 + 2H+ I2 + 2H2O –1 –1 45 –2 Change = × (+1) = +2 Oxidation Change = × (–1) = –2 Reduction The equation shows that Iÿ is oxidized by H2 O2 , and therefore H2 O2 acts as an oxidizing agent 1.17 Empirical, molecular and structural formulae Empirical and molecular formulae The empirical formula of a compound is the simplest possible formula (with integer subscripts) that gives the composition of the compound The molecular formula of a compound is the formula consistent with the relative molecular mass Worked example 1.36 The empirical formula of a compound gives the ratio of atoms of elements that combine to make the compound However, this is not necessarily the same as the molecular formula which tells you the number of atoms of the constituent elements in line with the relative molar mass of the compound The relationship between the empirical and molecular formulae of a compound is illustrated using ethane, in which the ratio of carbon : hydrogen atoms is : This means that the empirical formula of ethane is CH3 The relative molecular mass of ethane is 30, corresponding to two CH3 units per molecule – the molecular formula is C2 H6 In methane, the empirical formula is CH4 and this corresponds directly to the molecular formula Empirical and molecular formulae A compound has the general formula Cn H2n þ This compound belongs to the family of alkanes It contains 83.7% carbon by mass Suggest the molecular formula of this compound Let the formula of the compound be Cx Hy The percentage composition of the compound is 83.7% C and 16.3% H " Ar values: see front inside cover of the book % of C by mass ¼ Mass of C in g  100 Total mass in g % of H by mass ¼ Mass of H in g  100 Total mass in g For a mole of the compound, the total mass in g ¼ relative molecular mass ¼ Mr % C ¼ 83:7 ¼ ð12:01  xÞ Â 100 Mr % H ẳ 16:3 ẳ 1:008 yị 100 Mr We not know Mr , but we can write down the ratio of moles of C : H atoms in the compound – this is the empirical formula From above: 46 CHAPTER Some basic concepts 12:01x 1:008y  100 ¼  100 83:7 16:3 y 100  16:3  12:01 ¼ ¼ 2:32 x 100  83:7  1:008 Mr ¼ The empirical formula of the compound is CH2:32 or C3 H7 The compound must fit into a family of compounds of general formula Cn H2n þ , and this suggests that the molecular formula is C6 H14 The working above sets the problem out in full; in practice, the empirical formula can be found as follows: % C ¼ 83:7 Ar C ¼ 12:01 % H ¼ 16:3 Ar H ¼ 1:008 Ratio C : H ¼ Worked example 1.37 83:7 16:3 : % 6:97 : 16:2 % : 2:32 ¼ : 12:01 1:008 Determining the molecular formula of a chromium oxide A binary oxide of chromium with Mr ¼ 152:02 contains 68.43% Cr Determine the empirical formula and the molecular formula [Data: Ar Cr ¼ 52:01; O ¼ 16:00] The binary chromium oxide contains only Cr and O The composition by mass of the compound is: 68:43% Cr ð100 68:43ị% O ẳ 31:57% O The ratio of moles of Cr : O ¼ ¼ %Cr %O : Ar Crị Ar Oị 68:43 31:57 : 52:01 16:00 ẳ 1:316 : 1:973 ¼ 1: 1:973 1:316 ¼ : 1:5 or : The empirical formula of the chromium oxide is therefore Cr2 O3 To find the molecular formula, we need the molecular mass, Mr : Mr ¼ 152:02 For the empirical formula Cr2 O3 : ½2  Ar Crị ỵ ẵ3 Ar Oị ẳ 52:01ị ỵ 16:00ị ẳ 152:02 This value matches the value of Mr Therefore, the molecular formula is the same as the empirical formula, Cr2 O3 Worked example 1.38 Determining the molecular formula of a sulfur chloride A binary chloride of sulfur with Mr ¼ 135:02 contains 52.51% Cl What are the empirical and molecular formulae of the compound? [Data: Ar S ¼ 32.06; Cl ¼ 35.45] Empirical, molecular and structural formulae 47 The binary sulfur chloride contains only S and Cl The composition by mass of the compound is: 52:51% Cl 100 52:51ị% S ẳ 47:49% S The ratio of moles of S : Cl ¼ ¼ %S %Cl : Ar ðSÞ Ar ðClÞ 47:49 52:51 : 32:06 35:45 ¼ 1:481 : 1:481 ¼ 1:1 The empirical formula of the sulfur chloride is therefore SCl The molecular mass, Mr , of the compound is 135.02 For the empirical formula SCl: Ar Sị ỵ Ar Clị ẳ 32:06 ỵ 35:45 ẳ 67:51 This mass must be doubled to obtain the value of Mr Therefore, the molecular formula is S2 Cl2 Structural formulae and ‘ball-and-stick’ models " Covalent bonding and structure: see Chapters 4–7 " Isomers: see Sections 6.11, 23.5 and 24.6–24.8 Neither the empirical nor the molecular formula provides information about the way in which the atoms of a molecule are connected The molecular formula H2 S does not indicate the arrangement of the three atoms in a molecule of hydrogen sulfide, but structure 1.3 is informative We could also have arrived at this structure by considering the number of valence electrons available for bonding For some molecular formulae, it is possible to connect the atoms in more than one reasonable way and the molecule is said to possess isomers An example is C4 H10 for which two structural formulae 1.4 and 1.5 can be drawn More detailed structural information can be obtained from ‘balland-stick’ models Models 1.6–1.8 correspond to formulae 1.3–1.5 We expand the discussion of drawing structural formulae for organic molecules in Section 24.3 (1.3) (1.4) (1.5) (1.6) (1.7) (1.8) 48 CHAPTER Some basic concepts 1.18 Basic nomenclature " Trivial names for selected organic and inorganic compounds are given in Appendix 13; this can be found on the companion website, www.pearsoned.co.uk/ housecroft In this section we outline some fundamental nomenclature and set out some important IUPAC ground rules for organic and inorganic compounds We summarize some widely used ‘trivial’ names of compounds and you should become familiar with these as well as with their IUPAC counterparts Detailed nomenclature rules and the specifics of organic chain numbering are found in Chapters 24–35 Basic organic classifications: straight chain alkanes The general formula for an alkane with a straight chain backbone is Cn H2n ỵ ; the molecule is saturated and contains only C–C and C–H bonds The simplest member of the family is methane CH4 , 1.9 The name methane carries with it two pieces of information: (1.9) meth- is the prefix that tells us there is one C atom in the carbon chain; the ending -ane denotes that the compound is an alkane The names of organic compounds are composite with the stem telling us the number of carbon atoms in the main chain of the molecule These are listed in the middle column of Table 1.8 For a straight chain alkane, the name is completed by adding -ane to the prefix in the table Compounds 1.10 (C3 H8 ) and 1.11 (C7 H16 ) are both alkanes Using the stems from Table 1.8, 1.10 with a 3carbon chain is called propane, and 1.11 with a 7-carbon chain is called heptane Basic organic classifications: functional groups A functional group in a molecule imparts a characteristic reactivity to the compound The functional group in an alkene is the C¼C double bond Table 1.8 Numerical descriptors Number Stem used to give the number of C atoms in an organic carbon chain Prefix used to describe the number of groups or substituents 10 11 12 13 14 15 16 17 18 19 20 methethpropbutpenthexheptoctnondecundecdodectridectetradecpentadechexadecheptadecoctadecnonadecicos- monoditritetrapentahexaheptaoctanonadecaundecadodecatridecatetradecapentadecahexadecaheptadecaoctadecanonadecaicosa- Basic nomenclature 49 and, in an alcohol, the functional group is the ÿOH unit The organic functional groups that we will describe in this book are listed in Table 1.9 The presence of most of these groups is recognized by using an instrumental technique such as infrared, electronic or nuclear magnetic resonance spectroscopy – see Chapters 12–14 (1.10) (1.11) Basic inorganic nomenclature The aim of the IUPAC nomenclature is to provide a compound or an ion with a name that is unambiguous One problem that we face when dealing with some inorganic elements is the possibility of a variable oxidation state A simple example is that of distinguishing between the two common oxides of carbon, i.e carbon monoxide and carbon dioxide The use of ‘mono-’ and ‘di-’ indicates that the compounds are CO and CO2 respectively The accepted numerical prefixes are listed in Table 1.8 Note that ‘di-’ should be used in preference to ‘bi-’ Binary compounds A binary compound is composed of only two types of elements We deal first with cases where there is no ambiguity over oxidation states of the elements present, e.g s-block metal Examples of such binary compounds include NaCl, CaO, HCl, Na2 S and MgBr2 The formula should be written with the more electropositive element (often a metal) placed first The names follow directly: A binary compound is composed of two types of elements NaCl Na2 S CaO MgBr2 HCl sodium chloride sodium sulfide calcium oxide magnesium bromide hydrogen chloride Anions The endings ‘-ide’, ‘-ite’ and ‘-ate’ generally signify an anionic species Some examples are listed in Table 1.10 The endings ‘-ate’ and ‘-ite’ tend to indicate the presence of oxygen in the anion (i.e an oxoanion) and are used for anions that are derived from oxoacids; e.g the oxoanion derived from sulfuric acid is a sulfate There is more than one accepted method of distinguishing between the different oxoanions of elements such as sulfur, nitrogen and phosphorus (Table 1.10) Older names such as sulfate, sulfite, nitrate and nitrite are still accepted within the IUPAC guidelines It is more informative, however, to incorporate the oxidation state of the element that is combining with oxygen, and an alternative name for sulfate is tetraoxosulfate(VI) This shows not only the oxidation state of the sulfur atom, but the number of oxygen atoms as well A third accepted option is to use the name tetraoxosulfate(2ÿ) In Chemistry, we have made every effort to stay within the 50 CHAPTER Some basic concepts Table 1.9 Selected functional groups for organic molecules Name of functional group Functional group Example; where it is in common use, a trivial name is given in pink Alcohol Ethanol (CH3 CH2 OH) Aldehyde Ethanal (CH3 CHO) Acetaldehyde Ketone Propanone (CH3 COCH3 ) Acetone Carboxylic acid Ethanoic acid (CH3 CO2 H) Acetic acid Ester Ethyl ethanoate (CH3 CO2 C2 H5 ) Ethyl acetate Ether R ¼ R’ or R 6¼ R’ Diethyl ether (C2 H5 OC2 H5 ) Amine Ethylamine (CH3 CH2 NH2 ) Amide Ethanamide (CH3 CONH2 ) Acetamide Halogenoalkane Bromoethane (CH3 CH2 Br) Acid chloride Ethanoyl chloride (CH3 COCl) Acetyl chloride Nitrile Ethanenitrile (CH3 CN) Acetonitrile Nitro Nitromethane (CH3 NO2 ) Thiol Ethanethiol (CH3 CH2 SH) Basic nomenclature 51 Table 1.10 Names of some common anions In some cases, more than one name is accepted by the IUPAC Formula of anion Name of anion Hÿ [OH]ÿ Fÿ Clÿ Brÿ Iÿ O2ÿ S2ÿ Se2ÿ N3ÿ N3 ÿ P3ÿ [CN]ÿ [NH2 ]ÿ [OCN]ÿ [SCN]ÿ [SO4 ]2ÿ [SO3 ]2ÿ [NO3 ]ÿ [NO2 ]ÿ [PO4 ]3ÿ [PO3 ]3ÿ [ClO4 ]ÿ [CO3 ]2ÿ Hydride Hydroxide Fluoride Chloride Bromide Iodide Oxide Sulfide Selenide Nitride Azide Phosphide Cyanide Amide Cyanate Thiocyanate Sulfate or tetraoxosulfate(VI) Sulfite or trioxosulfate(IV) Nitrate or trioxonitrate(V) Nitrite or dioxonitrate(III) Phosphate or tetraoxophosphate(V) Phosphite or trioxophosphate(III) Perchlorate or tetraoxochlorate(VII) Carbonate or trioxocarbonate(IV) IUPAC recommendations while retaining the most common alternatives, e.g sulfate Oxidation states The oxidation state is very often indicated by using the Stock system of Roman numerals The numeral is always an integer and is placed after the name of the element to which it refers; Table 1.10 shows its application to some oxoanions The oxidation number can be zero, positive or negative.§ An oxidation state is assumed to be positive unless otherwise indicated by the use of a negative sign Thus, (III) is taken to read (ỵIII); but for the negative state, write (ÿIII) In a formula, the oxidation state is written as a superscript (e.g [MnVII O4 ]ÿ ) but in a name, it is written on the line (e.g iron(II) bromide) Its use is important when the name could be ambiguous (see below) Binary compounds We look now at binary compounds where there could be an ambiguity over the oxidation state of the more electropositive element (often a metal) Examples of such compounds include FeCl3 , SO2 , SO3 , ClF, ClF3 and SnCl2 Simply writing ‘iron chloride’ does not distinguish between the chlorides of iron(II) and iron(III), and for FeCl3 it is necessary to write iron(III) chloride Another accepted name is iron trichloride The oxidation state of sulfur in SO2 can be seen immediately in the name sulfur(IV) oxide, but also acceptable is the name sulfur dioxide Similarly, SO3 can be named sulfur(VI) oxide or sulfur trioxide § A zero oxidation state is signified by 0, although this is not a Roman numeral 52 CHAPTER Some basic concepts Table 1.11 The names of some common, non-metallic cations Formula of cation Name of cation Hỵ [H3 O]ỵ [NH4 ]ỵ [NO]ỵ [NO2 ]ỵ [N2 H5 ]ỵ Hydrogen ion Oxonium ion Ammonium ion Nitrosyl ion Nitryl ion Hydrazinium ion Accepted names for ClF, ClF3 and SnCl2 are: ClF ClF3 SnCl2 chlorine(I) fluoride or chlorine monofluoride chlorine(III) fluoride or chlorine trifluoride tin(II) chloride or tin dichloride Cations Cations of metals where the oxidation state does not usually vary, notably the s-block elements, may be named by using the name of the metal itself (e.g sodium ion, barium ion), although the charge may be indicated (e.g sodium(I) ion or sodium(1ỵ) ion, barium(II) ion or barium(2ỵ) ion) Where there may be an ambiguity, the charge must be shown (e.g iron(II) or iron(2ỵ) ion, copper(II) or copper(2ỵ) ion, thallium(I) or thallium(1ỵ) ion).Đ The names of polyatomic cations are introduced as they appear in the textbook but Table 1.11 lists some of the most common inorganic, non-metallic cations with which you may already be familiar Look for the ending ‘-ium’; this often signifies the presence of a cation, although remember that ‘-ium’ is a common ending in the name of elemental metals (see Section 1.5) Many metal ions, in particular those in the d-block, occur as complex ions; these are described in Chapter 23 1.19 Final comments The aim of this first chapter is to provide a point of reference for basic chemical definitions, ones that you have probably encountered before beginning a first year university chemistry course If you find later in the book that a concept appears to be ‘assumed’, you should find some revision material to help you in Chapter Section 1.18 gives some basic guidelines for naming organic and inorganic compounds, and more detailed nomenclature appears as the book progresses We have deliberately not called Chapter 1: ‘Introduction’ There is often a tendency to pass through chapters so-labelled without paying attention to them In this text, Chapter is designed to help you and to remind you of basic issues § An older form of nomenclature which is commonly encountered still uses the suffix ‘-ous’ to describe the lower oxidation state and ‘-ic’ for the higher one Thus, copper(I) is cuprous and copper(II) is cupric This system is unambiguous only when the metal exhibits only two oxidation states 53 Problems PROBLEMS Use values of Ar from the front inside cover of the book 1.1 What is 0.0006 m in (a) mm, (b) pm, (c) cm, (d) nm? 1.2 A typical C¼O bond distance in an aldehyde is 122 pm What is this in nm? 1.3 The relative molecular mass of NaCl is 58.44 and its density is 2.16 g cmÿ3 What is the volume of mole of NaCl in m3 ? 1.4 The equation E ¼ h# relates the Planck constant (h) to energy and frequency Determine the SI units of the Planck constant 1.5 Kinetic energy is given by the equation: E ¼ 12 mv By going back to the base SI units, show that the units on the left- and right-hand sides of this equation are compatible 1.6 Calculate the relative atomic mass of a sample of naturally occurring boron which contains 19.9% 10 11 B and 80.1% B Accurate masses of the isotopes to sig fig are 10.0 and 11.0 1.7 The mass spectrum of molecular bromine shows three lines for the parent ion, Br2 ỵ The isotopes for 81 bromine are 79 35 Br (50%) and 35 Br (50%) Explain why there are three lines and predict their mass values and relative intensities Predict what the mass spectrum of HBr would look like; isotopes of hydrogen are given in Section 1.7 (Ignore fragmentation; see Chapter 10.) 1.8 Convert the volume of each of the following to conditions of standard temperature (273 K) and pressure (1 bar ¼ 1:00  105 Pa) and give your answer in m3 in each case: (a) 30.0 cm3 of CO2 at 290 K and 101 325 Pa (1 atm) (b) 5.30 dm3 of H2 at 298 K and 100 kPa (1 bar) (c) 0.300 m3 of N2 at 263 K and 102 kPa (d) 222 m3 of CH4 at 298 K and 200 000 Pa (2 bar) 1.9 The partial pressure of helium in a 50.0 dm3 gas mixture at 285 K and 105 Pa is 4:0  104 Pa How many moles of helium are present? 1.10 A 20.0 dm3 sample of gas at 273 K and 2.0 bar pressure contains 0.50 moles N2 and 0.70 moles Ar What is the partial pressure of each gas, and are there any other gases in the sample? (Volume of one mole of ideal gas at 273 K, 1:00  105 Pa (1 bar) ¼ 22.7 dm3 ) 1.11 Determine the amount (in moles) present in each of the following: (a) 0.44 g PF3 , (b) 1.00 dm3 gaseous PF3 at 293 K and 2:00  105 Pa, (c) 3.480 g MnO2 , (d) 0.0420 g MgCO3 (Volume of mole of ideal gas at 273 K, 105 Pa ¼ 22.7 dm3 ) 1.12 What mass of solid is required to prepare 100.0 cm3 of each of the following solutions: (a) 0.0100 mol dmÿ3 KI; (b) 0.200 mol dmÿ3 NaCl; (c) 0.0500 mol dmÿ3 Na2 SO4 ? 1.13 With reference to the periodic table, write down the likely formulae of compounds formed between: (a) sodium and iodine, (b) magnesium and chlorine, (c) magnesium and oxygen, (d) calcium and fluorine, (e) lithium and nitrogen, (f ) calcium and phosphorus, (g) sodium and sulfur and (h) hydrogen and sulfur 1.14 Use the information in the periodic table to predict the likely formulae of the oxide, chloride, fluoride and hydride formed by aluminium 1.15 Give balanced equations for the formation of each of the compounds in problems 1.13 and 1.14 from their constituent elements 1.16 What you understand by each of the following terms: proton, electron, neutron, nucleus, atom, radical, ion, cation, anion, molecule, covalent bond, compound, isotope, allotrope? 1.17 Suggest whether you think each of the following species will exhibit covalent or ionic bonding Which of the species are compounds and which are molecular: (a) NaCl; (b) N2 ; (c) SO3 ; (d) KI; (e) NO2 ; (f ) Na2 SO4 ; (g) [MnO4 ]ÿ ; (h) CH3 OH; (i) CO2 ; (j) C2 H6 ; (k) HCl; (l) [SO4 ]2ÿ ? 1.18 Determine the oxidation state of nitrogen in each of the following oxides: (a) N2 O; (b) NO; (c) NO2 , (d) N2 O3 ; (e) N2 O4 ; (f ) N2 O5 1.19 In each reaction below, assign the oxidation and reduction steps, and, for (b)–(g), show that the changes in oxidation states for the oxidation and reduction processes balance: (a) Cu2ỵ aqị ỵ 2e Cusị (b) Mgsị ỵ H2 SO4 aqị MgSO4 aqị ỵ H2 gị (c) 2Casị ỵ O2 gị 2CaOsị (d) 2Fesị ỵ 3Cl2 gị 2FeCl3 sị (e) Cu(s) ỵ 2AgNO3 (aq) Cu(NO3 )2 (aq) ỵ 2Ag(s) (f ) CuOsị ỵ H2 gị Cusị ỵ H2 Ogị (g) ẵMnO4 aqị ỵ 5Fe2ỵ aqị ỵ 8Hỵ aqị Mn2ỵ aqị ỵ 5Fe3ỵ aqị ỵ 4H2 OðlÞ " " " " " " " 1.20 (a) In a compound, oxygen is usually assigned an oxidation state of ÿ2 What is the formal oxidation state in the allotropes O2 and O3 , in the compound H2 O2 , and in the ions [O2 ]2 and [O2 ]ỵ ? (b) Determine the oxidation and reduction steps during the decomposition of hydrogen peroxide which occurs by the following reaction: 2H2 O2 lị 2H2 Olị ỵ O2 gị " 1.21 Give a systematic name for each of the following compounds: (a) Na2 CO3 ; (b) FeBr3 ; (c) CoSO4 ; (d) BaCl2 ; (e) Fe2 O3 ; (f ) Fe(OH)2 ; (g) LiI; (h) KCN; (i) KSCN; (j) Ca3 P2 54 CHAPTER Some basic concepts 1.22 Write down the formula of each of the following compounds: (a) nickel(II) iodide; (b) ammonium nitrate; (c) barium hydroxide; (d) iron(III) sulfate; (e) iron(II) sulfite; (f ) aluminium hydride; (g) lead(IV) oxide; (h) tin(II) sulfide 1.23 How many atoms make up the carbon chain in (a) octane, (b) hexane, (c) propane, (d) decane, (e) butane? 1.24 Balance the following equations: (a) C4 H10 ỵ O2 CO2 ỵ H2 O (b) SO2 ỵ O2 SO3 (c) HCl ỵ Ca H2 ỵ CaCl2 (d) Na2 CO3 ỵ HCl NaCl ỵ CO2 ỵ H2 O (e) HNO3 ỵ Mg MgNO3 ị2 ỵ H2 (f ) H3 PO4 ỵ NaOH Na2 HPO4 ỵ H2 O " " " " 1.30 Balance each of the following equations: (a) Fe3ỵ þ H2 ÿÿ Fe2þ þ Hþ (b) Cl2 þ Brÿ Cl ỵ Br2 (c) Fe2ỵ ỵ ẵCr2 O7 þ Hþ ÿÿ Fe3þ þ Cr3þ þ H2 O 3þ (d) NH2 OH ỵ Fe N2 O ỵ Fe2ỵ þ H2 O þ Hþ (e) ½S2 O3 Š2ÿ þ I2 ẵS4 O6 ỵ I (f ) ẵMoO4 ỵ ẵPO4 ỵ Hỵ ẵPMo12 O40 þ H2 O (g) HNO3 þ H2 SO4 ÿÿ ½H3 Oỵ ỵ ẵNO2 ỵ ỵ ẵHSO4 " " " " " " " 1.31 Balance the following equation for the reaction of magnesium with dilute nitric acid: " " 1.25 Find x, y and z in the following reactions: (a) 2CO ỵ O2 2COx (b) N2 ỵ xH2 yNH3 (c) Mg ỵ 2HNO3 MgNO3 ịx ỵ H2 (d) xH2 O2 yH2 O ỵ zO2 (e) xHCl ỵ CaCO3 CaCly ỵ CO2 ỵ H2 O (f ) xNaOH ỵ H2 SO4 Nay SO4 ỵ zH2 O (g) MnO2 ỵ xHCl MnCl2 ỵ Cl2 þ yH2 O (h) xNa2 S2 O3 þ I2 ÿÿ yNaI ỵ zNa2 S4 O6 " " Mgsị ỵ HNO3 aqị MgNO3 ị2 aqị ỵ H2 gị " Use the balanced equation to determine the mass of Mg that will completely react with 100.0 cm3 0.50 M nitric acid " " " " " 1.32 Balance the following equation for the reaction of aqueous phosphoric acid with sodium hydroxide: H3 PO4 aqị ỵ NaOHaqị " " 1.26 Balance the following equations (a) Fe ỵ Cl2 FeCl3 (b) SiCl4 ỵ H2 O SiO2 ỵ HCl (c) Al2 O3 ỵ NaOH ỵ H2 O Na3 AlOHị6 (d) K2 CO3 ỵ HNO3 KNO3 ỵ H2 O ỵ CO2 heat (e) Fe2 O3 ỵ CO Fe ỵ CO2 (f ) H2 C2 O4 ỵ KOH K2 C2 O4 ỵ H2 O " " " " Na3 PO4 aqị ỵ H2 Olị 15.0 cm aqueous phosphoric acid of concentration 0.200 mol dmÿ3 is added to 50.0 cm3 aqueous NaOH of concentration 2.00 mol dmÿ3 Which reagent is in excess, and how many moles of this reagent remain unreacted? " " 1.27 Balance the following equations (a) AgNO3 ỵ MgCl2 AgCl ỵ MgNO3 ị2 (b) PbO2 CCH3 ị2 ỵ H2 S PbS ỵ CH3 CO2 H (c) BaCl2 ỵ K2 SO4 BaSO4 ỵ KCl (d) PbNO3 ị2 ỵ KI PbI2 ỵ KNO3 (e) CaHCO3 ị2 ỵ CaOHị2 CaCO3 ỵ H2 O " 1.33 Balance the following equation for the precipitation of silver chromate: AgNO3 aqị ỵ K2 CrO4 aqị " Ag2 CrO4 sị ỵ KNO3 aqị " " " " 1.28 Balance the following equations (a) C3 H8 ỵ Cl2 C3 H5 Cl3 ỵ HCl (b) C6 H14 ỵ O2 CO2 ỵ H2 O (c) C2 H5 OH þ Na ÿÿ C2 H5 ONa þ H2 (d) C2 H2 ỵ Br2 C2 H2 Br4 (e) CaC2 ỵ H2 O CaOHị2 ỵ C2 H2 " " 5.00 g of K2 CrO4 is dissolved in water and the volume of the solution is made up to 100.0 cm3 25.0 cm3 of a 0.100 mol dmÿ3 solution of AgNO3 is added to the solution of K2 CrO4 Determine the mass of Ag2 CrO4 that is formed 1.34 Balance the following equation: " " ẵC2 O4 aqị ỵ ẵMnO4 aqị ỵ Hỵ aqị " 2ỵ Mn aqị ỵ CO2 gị ỵ H2 Olị " 1.29 In each of the following, mixing the aqueous ions shown will produce a precipitate Write the formula of the neutral product, and then balance the equation (a) Agỵ aqị ỵ Cl aqị (b) Mg2ỵ aqị ỵ ẵOH aqị (c) Pb2ỵ aqị ỵ S2 aqị (d) Fe3ỵ aqị ỵ ẵOH aqị (e) Ca2ỵ aqị ỵ ẵPO4 aqị (f ) Agỵ aqị ỵ ẵSO4 aqị ÿÿ " What volume of 0.200 M aqueous KMnO4 will react completely with 25.0 cm3 0.200 M aqueous K2 C2 O4 in the presence of excess acid? 1.35 5.00 g of solid CaCO3 is thermally decomposed in the following reaction: " " " " " heat CaCO3 ðsÞ ÿÿ CaOðsÞ þ CO2 ðgÞ " What mass of CaO is formed? This oxide reacts with water to give Ca(OH)2 Write a balanced equation for this process Problems 1.36 Balance the following equation for the reaction of sodium thiosulfate with diiodine: Na2 S2 O3 ỵ I2 Na2 S4 O6 þ NaI " Diiodine is insoluble in water, but dissolves in aqueous potassium iodide solution 0.0250 dm3 of a solution of I2 in aqueous KI reacts exactly with 0.0213 dm3 0.120 M sodium thiosulfate solution What is the concentration of the diiodine solution? 1.37 An organic compound A contains 40.66% C, 23.72% N and 8.53% H The compound also contains oxygen Determine the empirical formula of the compound The molecular mass of A is 59.07 What is its molecular formula? 1.38 (a) A chloride of platinum, PtClx , contains 26.6% Cl What is the oxidation state of the platinum in this compound? 55 (b) Two oxides of iron, A and B, contain 69.94 and 72.36% Fe, respectively For A, Mr ¼ 159:70; for B, Mr ¼ 231:55 What are the empirical formulae of A and B? 1.39 Crystalline copper(II) sulfate contains water The formula of the crystalline solid is CuSO4 :xH2 O Determine x if the crystals contain 12.84% S and 4.04% H 1.40 (a) Glucose, Cx Hy Oz , contains 40.00% C and 6.71% H What is the empirical formula of glucose? If the molecular mass of glucose is 180.16, determine its molecular formula (b) A fluoride of tungsten, WFx , contains 38.27% F What is the oxidation state of tungsten in this compound? ... Mathematics Tutor Annotated links to relevant sites on the web CHEMISTRY An Introduction to Organic, Inorganic and Physical Chemistry 3rd edition Catherine E Housecroft Edwin C Constable Pearson... scientific world, an understanding of chemical concepts is essential It is traditional to split chemistry into the three branches of inorganic, organic and physical; theoretical chemistry may be... is chemistry and why is it important? What is the IUPAC? SI units The proton, electron and neutron The elements States of matter Atoms and isotopes The mole and the Avogadro constant Gas laws and