CHM 001 Access Chemistry PDF

CHM 001: Access Chemistry Course Development Course Developers: Dr A. 0. Oyewale Mrs. F M Folarinmi Course Co-ordinators: Mr. Adakole Ape Dr Femi Peters Designer: Mrs. M J. Afolabi...

CHM 001: Access Chemistry Course Development Course Developers: Dr A. 0. Oyewale Mrs. F M Folarinmi Course Co-ordinators: Mr. Adakole Ape Dr Femi Peters Designer: Mrs. M J. Afolabi National Open University of Nigeria Kaduna Campus Kaduna External Course Editor: Prof S. T Bajah Education Consultant Madan. NATIONAL OPEN UNIVERSITY OF NIGERIA CHM 001: Foundation Chemistry National Open University of Nigeria Headquarters National Open University of Nigeria 14/16 Ahmadu Bello Way Victoria Island Lagos Abuja Annex ice 245 Samuel Adesujo Ademulegun Street Central Business District Opposite Arewa Suites Abuja E-mail: [email protected] URL: www.nou.edu.ng National Open University of Nigeria, 2004 First published 2004 ISBN 978-058-240-1 All Rights Reserved. Published by Folubee Prints for National Open University of Nigeria  Course Guide Introduction Welcome to this chemistry course. We suppose that many of you are taking this course because you want to strengthen your 'background in chemistry. Armed with this strong background in chemistry, you can then proceed to advanced programmes in chemistry. The training of biologists, geoscientists, medical doctors, nurses, soil scientists, food scientists, chemical engineers, petroleum engineers and many others requires that students in those areas have a good exposure in chemistry. We have packaged this chemistry course in such a way that you will learn chemistry using a technique peculiar to the open learning system. This technique has been adopted by the National Open University of Nigeria ; a most of the courses offered to students. You will here learn the content of this chemistry course at a reasonable pace. You will need to master the language chemists use to describe the world around us. The language is simple, interesting and specific to the subject, chemistry. We wish you success with the course and hope that you will fmd it both interesting and useful. What you will learn in this course This course is titled "foundation chemistry". It assumes that you had taken a chemistry course before but probably you did not make the grade you desired. So you are re-taking the course with the aim of doing better. There is no doubt therefore that you are familiar with the language of chemistry. In this present course, you will be presented information in chemistry in a structured way to make learning easier. All the units follow the same pattern and so after the first few units, the rest will become easy to follow. The whole range of senior school certificate examination (SSCE) syllabus has been covered in this course. Learning Outcomes-Aims and Objectives The broad aims of this foundation chemistry course can be summarised thus. The course aims to provide you with chemistry content that will be sufficient for you to have the equivalent of the SSCE. Thus you will have solid foundation in chemistry which will enable you go into an advanced science course needing a background of chemistry. The objectives of this course are set out below. On completion of the course, you should be able to: 1. Distinguish between chemistry and the other science subjects. 2. Discuss the role of chemistry in our every day living. 3. Apply the language of chemistry in describing the world around you. 4. Carry out simple chemical calculations. 5. Identify chemical process in what goes on in your environment. Course Materials Three different sets of course materials are provided: A Course Guide which spells out the broad details of the foundation chemistry course including the aims and objectives.  The Study Units with detailed learning information. Each study unit has a set of performance objectives along with other relevant learner guide. There are forty five (45) study units grouped into three main sections referred to as volumes: Volume One: General and Physical Chemistry Volume Two: Physical and Inorganic Chemistry Volume Three: Organic Chemistry. A set of recommended Chemistry textbooks is listed at the end of each study unit. iv  Table of Contents Course Guide iii Volume One: General and Physical Chemistry xxv Unit 1: Elementary Units in Chemical Reactions 1 1.0 Introduction 1.1 Objectives 1 1.2 Elements compounds and mixtures 2 1.3 The particulate nature of matter 3 1.3.1 The concept of atoms and molecules 3 1.3.2 The atomic theory and chemical reaction 3 1.3.3 The constituents of the atom 3 1.3.4 Relative atomic and molecular masses 4 1.3.5 Chemical symbols and formulae 4 1.3.6 Other laws of chemical combination. 6 1.4 Chemical reactions and equations 7 1.5 Conclusion 7 1.6 Summary 7 1.7 Tutor-Marked Assignments 7 1.8 References. 8 Unit 2: Electronic Configuration 1— Static Model 9 2.0 Introduction 9 2.1 Objectives 9 2.2 The electrical nature of the atom 9 2.2.1 Historical evidences 9 2.2.2 Atomic and mass numbers 10 2.3 Atomic models 11 2.3.1 The nuclear atom 11 2.3.2 Electronic energy levels 12 2.3.3 Electronic congifuration 12 2.3.4 Ion formation 14 2.4 Conclusion 14 2.5 Summary 14 2.6 Tutor-Marked Assignments 14 2.7 References 15 Unit 3: The Nucleus and Radioactivity 16 3.0 Introduction 16  3.1 Objectives 16 3.2 The nucleus 16 3.3 Nuclear reactions 17 3.3.1 Radioactivity 17 3.3.2 Nuclear radiations 18 3.3.3 Nuclear fusion 19 3.3.4 Nuclear fission 19 3.3.5 Other examples of nuclear transformation reactions 19 3.4 Uses of radioactivity 19 3.5 Hazards of radioactivity 20 3.6 Conclusion 20 3.7 Summary 20 3.8 Tutor-Marked Assignments 20 3.9 References 21 Unit 4: Chemical Bonding 1: Electrovalent, Covalent and Co-ordinate Covalent 22 4.0 Introduction 22 4.1 Objectives 22 4.2 Electrovalent (Ionic) bonding 23 4.2.1 Structure of electrovalent compounds 24 4.2.2 Properties of electrovalent compounds 24 4.3 Covalent bonding 24 4.3.1 Properties of covalent compounds 25 4.4 Co-ordinate (dative) covalent bonding 25 4.5 Conclusion 26 4.6 Summary 27 4.7 Tutor-Marked Assignments 27 4.8 References 27 Unit 5: Chemical Bonding II: Metallic and Intermolecular Bonding 28 5.0 Introduction 28 5.1 Objectives 28 5.2 Metallic Bonding 28 5.2.1 Uses of Metals 29 5.3 Intermolecular bonding 30 5.3.1 Van der Waal's forces 30 5.3.2 Dipole-dipole attractions 30 5.3.3 Hydrogen bonding 31 5.4 Conclusion 31 5.5 Summary 32 5.6 Tutor-Marked Assignments 32 vi  5.7 References 32 Unit 6: Periodic Table I: Classification of Elements 33 6.0. Introduction 33 6.1 Objectives 33 6.2 The Periodic law and periodic table 33 6.3 Description of the periodic table 34 6.3.1 The periods 35 6.3.2 The main groups 35 6.3.3 The transition elements 37 6.4 Conclusion 37 6.5 Summary 37 6.6 Tutor-Marked Assignments 38 6.7 References 38 Unit 7: Electronic Configuration II: Atomic Orbital Model 39 7.0 Introduction 39 7.1 Objectives 39 7.2 Quantum theory of atomic orbitals 40 7.2.1 The principal quantum number (n) 40 7.2.2 The subsidiary or azimuthal quantum number (4) 40 7.2.3 The magnetic quantum number (m) 40 7.2.4 The spin quantum number (s) 40 7.3 The shape of atomic orbitals 41 7.4 Electronic configuration of atoms — orbital model 41 7.5 Electronic configuration and periodic classification of elements 42 7.6 Conclusion 43 7.7 Summary 43 7.8 Tutor-Marked Assignments 44 7.9 References 44 Unit 8: Periodic Table II: Gradations of Atomic properties 45 8.0 Introduction 45 8.1 Objectives 45 8.2 Brief revision of electronic configuration 45 8.3 Electronic configuration and the periodic table 46 8.4 The periodic table and atomic properties 48 8.4.1 Atomic size 48 8.4.2 Ionic radius 49 8.4.3 Ionization energy 49 8.4.4. Electron affinity 51 8.4.5 Electronegativity 52  8.5 Conclusion 53 8.6 Summary 53 8.7 Tutor-Marked Assignments 53 8.8 References 53 Unit 9: Mole Concept 1 54 9.0 Introduction 54 9.1 Objectives 54 9.2 The mole defined 54 9.3 The mole and the molar mass 55 9.4 Finding the formula of a compound 56 9.5 Calculating the percentages of elements in a compound 58 9.6 Using empirical formula to calculate unknown atomic mass of an element. 58 9.7 The mole in chemical reactions (yield and percentage yield) 59 9.7.1 Calculation of percentage yield 59 9.7.2 Molar volume for gases 60 9.8 Conclusion 60 9.9 Summary 60 9.10 Tutor-Marked Assignments 61 9.11 References 61 Unit 10: Mole Concept II 62 10.0 Introduction 62 10.1 Objectives 62 10.2 Mole concept in solutions 62 10.2.1 The molar and molal solutions 62 10.2.2 Calculation of mole and mass in solution 63 10.3 The mole concept in volumetric analysis 64 10.3.1 Determination of concentration of a solution 64 10.3.2 Determination of percentage purity 65 10.3.3 Mole concept in solution dilution 65 10.4 The mole concept in electrolysis 67 10.4.1 The mole concept and the Faraday 67 10.4.2 The mole concept and the efficiency of an electrolytic process. 68 10.5 Conclusion 68 10.6 Summary 68 10.7 Tutor-Marked Assignments 69 10.8 References 69 Unit 11: The Kinetic Theory and States of Matter ------ 70 11.0 Introduction 70  11 1 Objectives 70 11.2 The kinetic theory 70 11.2.1 The postulate of the kinetic theory 70 11.2.2 The kinetic theory and practical evidences 71 11.3 The kinetic theory and the state of matter. 71 11.4. The kinetic theory and change of state 72 11.4.1 Melting and melting point 72 11.4.2 Vaporisation and boiling point 72 11.4.3 Evaporation 72 11.4.4 Sublimation 72 1.5 Heating and cooling graphs 73 11.6 Conclusion 73 11.7 Summary 73 11.8 Tutor-Marked Assignments. 74 11.9 References 74 Unit 12: The Gas Laws (I): Boyle's and Charles' law and the General Gas Equations 75 12.0 Introduction 75 12.1 Objectives 75 12.2 The kinetic molecular theory of gases. 75 12.3 Boyle's law 76 12.3.1 Statement of Boyle's law 76 12.3.2 How kinetic theory explains Boyle's law 77 12.4 Charles' Law 78 12.4.1 Charles' law and the kelvin temperature scale. 78 12.4.2 The statement of Charles' law 78 12.4.3 How the kinetic theory explains Charles' law 80 12.5 The general gas equation 80 12.6 Conclusion 81 12.7 Summary 81 12.8 Tutor-Marked Assignments 81 12.9 References 81 Unit 13: The Gas Laws 11: Dalton's. Graham's, Avogadro's and Gay Lussac's Laws 13.0 Introduction I nci ;lama: 1. )aitot1k. - ____  Avogadro's law and its applications 84 13.4 Gay Lussac's law of combining volumes 85 13.5 Conclusion 85 13.6 Summary 85 13.7 Tutor-Marked Assignments 85 13.8 References 85 13.9 Liquids ----------------------- -- ----------- 86 Unit 14: Introduction 86 14.0 Objectives 86 14.1 The vapour pressure of liquids 86 14.2 Boiling and boiling point 87 14.3 Boiling point - a criterion of purity 87 14.4 Methods of boiling point determination 88 14.5 14.6 Conclusion 89 Summary 89 14.7 Tutor-Marked Assignments 89 14.8 References 89 14.9 90 Unit 15: 90 15.0 Introduction Objectives 90 15.1 Solid classification 90 15.2 Solid properties 91 15.3 The molecular solid 91 15.3.1 The metallic solid 91 15.3.2 The ionic solid 91 15.3.3 15.3.4 The covalent solid 91 92 15.4 The structure of diamond and graphite Melting and melting point 93 15.5 94 15.6 Conclusion 94 15.7 Summary 15.8 Tutor-Marked Assignments 15.9 References  Volume Two: Physical and Inorganic Chemistry Unit 1: Energy and Energy Changes in Chemical and Physical Processes (1) ---- 96 1.0 Introduction 96 1.1 Objectives 96 1.2 Energy changes in chemical processes 96 1.2.1 Heat content and heat of reaction 96 1.2.2 Exothermic and endothermic reactions 97 1.2.3 Energy level diagrams 97 1.2.4 The standard state 97 1.2.5 The thermochemical equations 98 1.3 Types of heat of reaction 98 1.3.1 Standard heat (enthalpy) of formation 98 1.3.2 Standard heat (enthalpy) of combustion 99 1.3.3 Standard heat (enthalpy) of neutralisation 99 1.3.4 Standard heat (enthalpy) of solution 99 1.3.5 Enthalpy of change of state 100 1.4 Determination of the heat of reaction 100 1.5 Conclusion 102 1.6 Summary 102 1.7 Tutor-Marked Assignments 103 1.8 References 103 Units 2: Changes in Chemical and Physical Processes (II) The Free Energy and Spontaneous Reactions. 104 2.0 Introduction 104 2.1 Objectives 105 2.2 Entropy and entropy change 105 2.3 Free energy and the free energy change 106 2.4 Applications of the free energy equation 106 2.4.1 An exothermic reaction accompanied by an increase in entropy 106 2.4.2 An endothermic reaction accompanied by a decrease in entropy 106 2.4.3 An exothermic reaction accompanied by a decrease in entropy 106 2.4.4 An endothermic reaction accompanied by an increase in entropy 107 2.4.5 A reaction in which the enthalpy change is zero. 107 2.5 Conclusion 107 2.6 Summary 107 2.7 Tutor-Marked Assignments 108 2.8 References 108 xi  IVY Unit 3: Chemical Kinetic (I) — Rates of Reactions 109 3.0 Introduction 109 3.1 Objectives 109 3.2 Rate of a Chemical reaction 109 3.2.1 Meaning of rate of reaction 110 3.2.2. Monitoring the rates of Chemical reactions 112 ' 3.2.3 Methods of monitoring the rates of reactions 112 3.3 Factors that affect reactions rates 112 3.3.1 Nature of the reactant 113 3.3.2 Concentration of the reactant 113 3.3.3 Temperature 113 3.3.4 Pressure / volume for gases 113 3.3.5 Surface area 113 3.3.6 Catalyst 114 3.3.7 Light 114 3.4 Conclusion 114 3.5 Summary 114 3.6 Tutor-Marked Assignments 114 3.7 References Chemical Kinetic (II) — Collision Theory and the Activation Energy --------- 115 Unit 4: 115 4.0 Intrroduction 115 4.1 Objectives 115 4.2 Activation energy and the reaction rate 117 4.3 The collision theory 117 4.3.1 The collision theory discussed 117 4.3.2 The collision theory and the nature of the reactant 117 4.3.3 The collision theory and the reactant concentration. 117 4.3.4 The collision theory and temperature 118 4.3.5 The collision theory and pressure 118 4.3.6 The collision theory and catalyst 118 4.4 Conclusion 118 4.5 Summary 118 4.6 Tutor-Marked Assignments 118 4.7 References 119 Unit 5 Chemical Equilibrium (I) — Reversible Reactions and Le Chatelier's Principle 119 5.0 Introduction 119 5.1 Objectives 120 5.2 Reversible reactions  5.2.1 Common examples of reversible reactions 120 5.2.2 Reversible reactions and the equilibrium state 121 5.2.3 The reversible reaction and the equilibrium constant 121 5.3 Factors affecting an equilibrium state 123 5.3.1 Le chatelier's principles 123 5.3.2 Le Chatelier's principle and temperature effect 123 5.3.3. Le chatelier's principle and pressure change 123 5.3.4. Le chatelier's principle and concentration change 123 5.4. Effect of catalyst on reversible reactions 124 5.5 Conclusion 124 5.6 Summary 124 5.7 Tutor-Marked Assignments 124 5.8 References 125 Unit 6: Chemical Equilibrium (II) Applications 126 6.0 Introduction 126 6.1 Objectives 126 6.2 Applications of the equilibrium law to gas reactions 126 6.2.1 The Haber process 126 6.2.2 The contact process 127 6.3 Applications of the equilibrium law to aqueous equilibria 127 6.3.1 Dissociation of weak acids and bases 127 6.3.2 Dissociation of water and the pH scale 128 6.3.3 Hydrolysis of salts 129 6.3.4 Buffer solutions 129 6.3.5 Common ion effect 129 6.3.6 Solubility product 130 6.4 Conclusion 131 6.5 Summary 131 6.6 Tutor-Marked Assignments 131 6.7 References 131 Unit 7: Acid, Bases and Salts: General Properties 132 7.0 Introduction 132 7.1 Objectives 132 7.2 Acids 132 7.2.1 Definition, examples and preparation 132 7.2.2 General properties of acids 133 7.2.3 Uses of acids 133 7.3 Bases 133  7.3.1 Definition, examples and preparation 133 7.3.2 General properties of bases 134 7.3.3 Uses of bases 134 7.4 Salts 134 7.4.1 Definition, examples and preparation 134 7.4.2 Types of salts 135 7.5 General characteristics of acids bases and salts 135 7.5.1 Ionisation 135 7.5.2 Deliquescence 136 7.5.3 Hygroscopy 136 7.5.4 Efflorescence 136 7.6 Conclusion 136 7.7 Summary 136 7.8 Tutor-Marked Assignments 136 7.9 References 137 Unit 8: Acids, Bases and Salts: Applications in Volumetric Analysis 138 8.0 Introduction 138 8.1 Objectives 138 8.2 Solutions 139 8.2.1 Types of solutions 139 8.2.2 Standard solutions 139 8.2.3 Primary standard substances 139 8.2.4 Preparing a standard solution 140 8.3 Acid-base titration 140 8.3.1 The equivalent points and end point of a titration 140 8.3.2 Acid-base indicators 141 8.3.3 Neutralisation curves and choice of indicator 141 8.3.4 The use of sodium trioxocarbonate (iv) in acid-base titration. 142 8.4 Calculations in volumetric analysis 142 8.5 Conclusion 144 8.6 Summary 144 8.7 Tutor-Marked Assignments 145 8.8 References 145 Unit 9: Electrolysis and Redox Reactions 146 9.0 Introduction 146 9.1 Objectives 146 9.2 Electrolysis 147 9.2.1 Electrolytes and non electrolytes 147 xiv  9.2.2 E.ecnlyric conduction and electrolysis 147 9.2.3 Precutntial discharge of ions during electrolysis 148 9.2.4 Examples of electrolysis of some salts 149 9.3 Redox reactions 149 9.3.1 Oxidation and reduction 149 9.3.2. Assignment of oxidation number 150 9.3.3 Redox reactions and the electrochemical cell 150 9.3.4 Comparison of the electrolytic and electrochemical cell. 151 9.0 Conclusion 151 9.5 Summary 152 9.6 Tutor-Marked Assignments 152 9.7 References 152 Unit 10: Practical Applications of Electrolysis 153 10.0 Introduction 153 10.1 Objectives 153 10.2 Industrial Applications of electrolysis 153 10.2.1 Purification of copper 153 10.2.2 Electroplating 154 10.2.3 Isolation of elements or extraction of ores 154 10.2.4 Preparation of sodium hydroxide 155 10.3 Quantitative electrolysis 155 10.4 Corrosion of metals 157 10.5 Rusting of iron 157 10.6 Methods used to prevent corrosion 157 10.7 Conclusion 158 10.8 Summary 158 10.9 Tutor-Marked Assignments 158 10.10 References 158 Unit 11: Carbon and its Compounds 159 11.0 Introduction 159 11.1 Objectives 159 11.2 Allotropes of Carbon 159 11.2.1 Diamond 160 11.2.2 Graphite 160 11.2.3 Amorphous carbon 161 11.3 General Properties of Carbon 162 11.4 Chemical Properties of Carbon 162 11.4.1 Combustion reactions 162 xv  11.4.2 Combination reactions 162 11.4.3 Carbon as reducing agent 162 11.4.4 Reactions with oxidizing agents 163 11.5 Oxides of Carbon 163 11.5.1 Carbon (IV) oxide 163 11.5.2 Carbon (II) oxide 164 11.6 Trioxocarbonates (IV) and hydrogen trioxocarbonates (IV) 166 11.6.1 Preparation of soluble trioxocarbonates (IV) salts 166 11.6.2 Preparation of insoluble trioxocarbonates (IV) salts 166 11.6.3 Hydrogen trioxocarbonate (IV) salts 167 11.6.4 Properties of trioxocarbonates (IV) salts 167 11.6.5 Uses of some important trioxocarbonates (IV) salts 167 11.7 Carbon cycle in nature 167 11.8 Conclusion 168 11.9 Summary 169 11.10 Tutor-Marked Assignments 169 11.11 References 169 Unit 12: The Chemistry of Important Industrial Gases 170 12.0 Introduction 170 12.1 Objectives 170 12.2 Hydrogen 170 12 2 1 Industrial preparation of hydrogen 171 12.2.2 Uses of hydrogen 172 12.3 Oxygen 172 12.3.1 Industrial preparation of oxygen 173 12.3.2 Uses of oxygen 174 12.4 Nitrogen 174 12.4.1 Industrial preparation of nitrogen 175 12.4.2 Uses of nitrogen 175 12.5 Chlorine 175 12.6 Oxides of sulphur 176 12.6.1 Sulphur (iv) oxide 176 12.6.2 Uses of sulphur(IV) oxide 177 12.6.3 Sulphur(VI) oxide 177 12.7 Conclusion 177 12.8 Summary 177 12.9 Tutor-Marked Assignments 177 12.10 References 177 xvi Unit 13: Metals I: General Characteristics 178 13.0 Introduction 178 13.1 Objectives 178 13.2 Physical properties of metals 178 13.3 Chemical properties of metals 179 13.4 Occurrence of metals in nature 180 13.5 Electrochemical or activity series 180 13.6 Extraction of metals 182 13.6.1 Electrolytic reduction 182 13.6.2 Chemical reduction 183 13.6.3 Thermal reduction 183 13.7 Conclusion 184 13.8 Summary 184 13.9 Tutor Marked Assignments 184 13.10 References 184 Unit 14: Metals II The transition metals - 185 14.0 Introduction 185 14.1 Objectives 185 14.2 Electronic configuration of 1st transition series metals 186 14.3 Physical properties of 1st transition series metals 186 14.4 Chemical reactivity 187 14.5 Variable oxidation states 188 14.6 Complex ion formation 189 14.7 Colour of transition metal ions 190 14.8 Catalytic activity 190 14.9 Industrial uses of transition metals 191 14.10 Alloys 191 14.11 Conclusion 192 14.12 Summary 192 14.13 Tutor-Marked Assignments 192 14.14 References 192 Unit 15: The Chemical Industry An Overview - 193 15.0 Introduction 193 15.1 Objectives 193 15.2 Sources of raw materials 194 15.3 Types of chemical industries 195 15.3.1 Heavy chemicals 195 15.3.2 Fine chemicals 195 15.3.3 Fertilizers 196 xvii  196 15.3.4 Plastics 196 15.3.5 Soaps and detergents 196 15.3.6 Pharmaceuticals and drugs 196 15.3.7 Other chemical industries The economics of industrial processes 197 15.4 197 15.5 The Nigerian chemical industry 198 15.6 Conclusion 198 15.7 Summary 198 15.8 Tutor-Marked Assignments 198 15.9 References 199 Volume Three: Organic Chemistry Introduction to Organic Chemistry and classification of Organic Molecules — 200 Unit 1: 200 1.0 Introduction 200 1.1. Objectives Definition and scope of organic chemistry 201 1.2. 201 1.3. Uniqueness of carbon atom Representation of organic molecules 202 1.4. Classification of organic compounds 202 1.5. Open-chain or aliphatic compounds 203 1.5.1 Saturated and unsaturated compounds 203 1.5.2 203 1.5.3 Aromatic compounds 203 1.5.4 Alicyclic compounds 203 1.5.5 Heterocyclic compound 204 1.6 Conclusion 204 1.7 Summary 204 1.8 Tutor-Marked Assignments 204 1.9 References The Homologous Series, Functional Groups and Isomerism 205 Unit 2: 205 2.0 Introduction 205 2.1 Objectives 205 2.2 The Homologous Series 206 2.3 Functional Groups 207 2.4 Isomerism 207 2.4.1 Structural isomerism 208 2.4.2 Geometric isomerism 209 2.5 Conclusion 209 2.6 Summary 209 2.7 Tutor-Marked Assignments  2.8 Re' acts 209 Unit 3: IUPAC Nomeclature of Organic Compounds 210 3.0 Introduction 210 3.1 Objectives 210 3.2 Nomenclature of organic compounds 211 3.2.1 The hydrocarbons - 211 3.2.2 Rules for IUPAC nomenclature 213 3.2.3 Naming non-hydrocarbon compounds 216 3.3 Conclusion 217 Summary 217 3.5 Tutor-Marked Assignments 217 3.6 References 217 Unit 4: Purification Methods and the Determination of Empirical, Molecular and Structural Formula Organic Compounds 218 4.0 Introduction 218 4.1 Objectives 218 4.2 Purification methods 219 4.2.1 Distillation 219 4.2.2 Crystallization 219 4.2.3 Chromatography 220 4.3 Empirical and molecular formula 221 4.3.1 Empirical formula 221 4.3.2 Molecular formula 223 4.4 Structural formula 224 4.5 Conclusion 227 4.6 Summary 227 4.7 Tutor-Marked Assignments 227 4.8 References 228 Unit 5: The Chemistry of Alkanes 229 5.0 Introduction 229 5.1 Objectives 229 5.2 Natural sources of alkanes 229 5.2.1 Natural gas 229 5.2.2 Petroleum or crude oil 230 5.2.3 Vegetable origin 230 5.3 Laboratory preparation of alkanes 230 5.4 Properties of alkanes 230 5.4.1 Physical properties 230 5.4.2 Chemical properties 230 xix  231 5.5 Uses of allcanes 232 5.6 Conclusion 232 5.7 Summary 232 5.8 Tutor-Marked Assignments 232 5.9 References The Chemistry of Alkanes and Alkynes 233 Unit 6: 233 6.0 Introduction 233 6.1 Objectives 234 6.2 Sources of alkenes and al cynes 234 6.2.1 Natural sources of alkenes 234 6.2.2 Manufacture of ethyne Laboratory preparation of ethene and ethyne 234 6.3. 235 6.4 Isomerism in alkenes 236 6.5 Chemical Properties 235 6.5.1 Alkenes 236 6.5.2 Alkynes 237 6.6 Test for unsaturation 237 6.7 Conclusion 237 6.8 Summary 237 6.9 Tutor-Marked Assignments 237 6.10 References Introduction to the Chemsitry of Benzene 238 Unit 7: 238 7.0 Introduction 238 7.1 Objectives 239 7.2 Structure and bonding in benzene 239 7.2.1 Resonance in benzene 240 7.3 Chemical properties of benzene 240 7.3.1 Substitution reactions of benzene 241 7.3.2 Addition reactions of benzene 241 7.4 Conclusion 242 7.5 Summary 242 7.6 Tutor-Marked Assignments 242 7.7 References 243 Unit 8: Introduction to Petroleum Chemistry 243 8.0 Introduction 243 8.1 Objectives 243 8.2 Crude oil reserves 244 8.3 The refining process  8.3.1 Fractional distillation 244 8.3.2 Quality of petrol — octane number 245 8.4 The conversion processes 246 8.4.1 Cracking 246 8.4.2 Isomerisation 246 8.4.3 Reforming 246 8.5 Petrochemicals 247 8.6 Oil production and the Nigerian economy and environment 247 8.7 Conclusion 247 8.8 Summary 248 8.9 Tutor-Marked Assignments 248 8.10 References 248 Unit 9: Alkanols I Classification, sources and Uses of Alkanols — 249 9.0 Introduction 249 9.1 Objectives 250 9.2 Classification of alkanols 250 9.3 Sources of alkanols 251 9.3.1 General methods of preparation 251 9.3.2 Manufacture of alkanols 251 9.4 Uses of ethanol 253 9.5 Conclusion 253 9.6 Summary 253 9.7 Tutor-Marked Assignments 253 9.8 References 253 Unit 10: Alkanols II Characteristics of Alkanols — 254 10.0 Introduction 254 10.1 Objectives 254 10.2 Intermolecular hydrogen bonding in allcanols 254 10.3 Physical properties of alkanols 255 10.4 Chemical properties of allcanols 255 10.4.1 Oxidation reactions 255 10.4.2 Elimination (dehydration) reactions 256 10.4.3 Substitution reactions 256 10.4.4 Iodofonn test 257 10.5 Laboratory test for alkanols 257 10.6 Conclusion 257 10.7 Summary 257 10.8 Tutor-Marked Assignments 257 xxi  257 10.9 References 258 Unit 11: Alkanoic Acids 258 11.0 Introduction 259 11.1 Objectives 259 11.2 Sources of alkanoic acids 259 11.2.1 Natural sources 259 11.2.2 General methods of preparation 259 11.3 Characteristics of alkanoic acids 259 11.3.1 Physical properties 260 11.3.2 Chemical properties 261 11.4 laboratory test for alkanoic acids 261 11.5 Uses of alkanoic acids 261 11.6 Conclusion 261 11.7 Summary 262 11.8 Tutor-Marked Assignments 262 11.9 References 263 Unit 12: Alkanoates 263 12.0 Introduction 263 12.1 Objectives Preparation of alkanoates — esterification 264 12.2 264 12.3 General characteristics of alkanoates 264 12.3.1 Physical properties 264 12.3.2 Chemical properties 265 12.4 Uses of alkanoates 265 12.5 Conclusion 265 12.6 Summary 265 12.7 Tutor-Marked Assignments 265 12.8 References 266 Unit 13: Fats, Oils and Amino Acids 266 13.0 Introduction 267 13.1 Objectives 267 13.2 Sources of fats and oils 268 13.3 Hardening of oils 268 13.4 Soap manufacture — saponification 269 13.5 Detergents 269 13.6 Uses of fats and oils 270 13.7 Tests for fats and oils 270 13.8 Amino Acids 271 13.9 Conclusion nil  13.10 Summary 271 13.11 Tutor Marked Assignments 271 13.12 References 271 Unit 14: An Introduction to Polymer Chemistry 272 14.0 Introduction 272 14.1 Objectives 272 14.2 Polymerisation processes 173 14.2.1 Addition polymerisation 273 14.2.2 Condensation polymerisation 274 14.3 Plastics and resins 275 14.4 Natural polymers 276 14.4.1 Carbohydrates 276 14.4.2 Proteins 278 14.5 Synthetic polymers 278 14.6 Conclusion 278 14.7 Summary 278 14.8 Tutor-Marked Assignments 279 14.9 References 279 Unit 15: Environmental Impact of the Chemical Industry Air and Water Pollution — — 280 15.0 Introduction 280 15.1 Objectives 280 15.2 Water pollution 280 15.2.1 Phosphates 281 15.2.2 Industrial discharges 281 15.2.3 Pesticides 281 15.2.4 Hot water 281 15.3 Air pollution 281 15.3.1 Solid particles 281 15.3.2 Oxides of carbon 282 15.3.3 Oxides of nitrogen and sulphur 282 15.3.4 Hydrocarbons 282 15.3.5 Chlorofluorocarbons 282 15.4 Biodegradable and non-biodegradable pollutatnts 282 15.5 Pollution control 283 15.6 Conclusion 283 15.7 Summary 283 15.8 Tutor-Marked Assignments 283 15.9 References 283 Answers 284  XXiV  Volume One General Physical Chemistry XXV  Unit 1 Elementary Units in Chemical Reactions 1.0 Introduction Chemistry is the study of matter and chemists in their investigations, study the properties and transformations of matter. Many materials that we use everyday, directly or indirectly are products of chemical research and examples of useful products of chemical reactions are limitless. What then is matter? Matter is anything that has mass and occupies space. Matter is classified into solid, liquid and gas. Solids have fixed shapes and sizes. Glass, sand and most metals are examples of solids. Liquids have fixed volumes but no fixed shapes Liquids take the shape of the containing vessel. Water and kerosene are examples of liquids. Gases flow and fill the entire space available. Air and cooking gas are examples of gases. The above classification is commonly referred to as physical classification. Matter can also be classified into elements, compounds and mixtures. This latter classification is referred to as chemical classification. It is remarkable that all these substances, solids, liquids, gases, elements, compounds and mixtures are built up from simple basic units. What is the basic unit of matter? What are the building blocks of matter and what laws govern the interaction of matter? The above are some of the questions that will be answered in this unit. 1.1 Objectives By the end of this unit, you should be able to: Define atoms, molecules, elements and compounds. Differentiate between atoms and molecules as well as elements and compounds. Explain why matter is said to be electrical in nature. State the postulates of Dalton's atomic theory and explain modifications to it. List and give relative masses and charges of the subatomic particles. State the laws of chemical combinations. Write chemical symbols and formulae for common elements and compounds. Determine chemical formulae from experimental data. Write and balance simple chemical equations. 1.2 Elements, Compounds and Mixtures An element is a pure substance, which cannot be split up into simple substances by a chemical reaction. A pure substance that can be broken down into elements is called a compound. There are over 100 chemical elements. Some occur naturally as free elements, mixed with other elements or compounds. Some are very rare, while most occur in combined state in compounds. The table 1.1 below gives a list of some el its and their total abundance either as free elements or in combined states. Table 1.1 List of some elements and their abundance Element Abundance in percentage by mass Oxygen 49.1 Silicon 26.1 Aluminium 7.5 Iron 4.7 Calcium 3.4 Sodium 2.6 Hydrogen 0.88 Chlorine 0.19 Carbon 0.09 Elements can be broadly classified into two groups, metals and non-metals. Examples of metals are copper and iron. The general characteristics of metals are lustre, good conductor of heat and electricity.Metals can be rolled and hammered into sheets and drawn into wires. They are used for roofing and electrical cables respectively. All metals are solids at room temperature except mercury which is a liquid at room temperature. About 75 percent of the elements are metals. Unlike metals, non metals do not have characteristic lustre. Many are gases at room temperature and others are solids except bromine which is a red brown liquid at room temperature. Non metals are non-conductors of heat and electricity. They cannot be rolled into sheets or drawn into wires like the metals. Oxygen, nitrogen, carbon and iodine are examples of non metals. A compound is formed when two or more elements combine chemically in fixed proportion by mass. The properties of compounds are different from those of the elements from which they are formed. A lot of energy is often required to split compounds into the constituent elements. There are limitless number of compounds. Sodium chloride, (common salt), water and calcium trioxocarbonate(iv), (marble) are examples of compounds. A mixture is a physical combination of elements or compounds. The composition of a mixture varies and the components are separated by physical methods. Such physical methods include, heating, cooling, dissolution, filtration and distillation. Air and petroleum are examples of mixtures. 2 1.3 The Particulate Nature of Matter 1.3.1 The concept of atoms and molecules Experiments show that very small quantities of matter have the same chemical properties or characteristics as larger ones of the same type e.g. a small iron nail and a big iron rod have the same chemical characteristics. For example both will rust when exposed to air and moisture and both will conduct heat and electricity. The smallest unit of an element that has the properties and characteristics of the element is the atom. The atom is the smallest unit of an element that can take part in a chemical reaction. The molecule is the smallest unit of a compound that has the characteristics of the compound. It is the smallest unit of the compound that can take part in a chemical reaction. The atom is to the element as molecule is to the compound. The behaviour of matter is explained using the above concept and chemical reactions are explained as combinations and rearrangement of atoms. The basic unit of matter in chemical reactions is the atom. Atoms and molecules are the building blocks of matter. 13.2 The atomic theory and chemical reaction A theory is a tested proposal to explain an observed statement of facts. A theory serves as a guide to new experiments. When proved incorrect or inadequate by experiment, a theory is discarded or modified so that new experimental facts can be accounted for. This latter statement is true of Dalton's atomic theory. Dalton's atomic theory form the basis of theory of the atom. It has been modified in the light of new experimental facts about the atom. The postulates of Dalton's atomic theory are: All elements are made up of small, indivisible particles called atoms. Atoms can neither be created nor destroyed. Atoms of the same element are alike in every respect and differ from atoms of all other elements. When atoms combine they do so in simple ratios. All chemical changes result from the combination or the separation of atoms. Some important modifications to the theory are: Matter is composed of more fundamental particles, some of which are electrically neutral, some carry positive charge and some negative charge. This implies that the atom is divisible. All atoms of the same element are not identical. The existence of atoms of the same element having different masses have been proved with the use of an instrument called mass spectrometer. Atoms of the same element having different masses are called isotopes. The smallest particle of an element that can take part in a chemical change is still the atom and in a chemical change atoms are neither created nor destroyed. Irrespective of masses, atoms of the same element have same chemical properties. 133 The constituents of the atom Matter is electrical in nature. Evidences for this assertion came from results of experiments of early scientists like Faraday, Thompson and Millikan. The negatively charged particle in matter is the electron. It has negligible mass. The proton is the positively charged particle. It carries the same magnitude of charge as the electron and is very much heavier than the electron. The third particle is the neutron, a neutral particle with a mass approximately equal to that of the proton. These three particles are constituents of the atom except the hydrogen atoms that do not contain neutrons. The number of each particle present in the atom varies from one element to another. For the atom (matter) to be electrically neutral, the number of protons must equal the number of electrons. Atoms of the same element will have the same number of protons and electrons but may have different numbers of neutrons. Such atoms will have different masses and are called isotopes. 3  Relative Mass Relative Charge Particle 5.5 x 10 -1 Electron 1.00 727 +1 Proton Neutron 1.00867 0 1.3.4 Relative atomic and molecular masses Chemistry is a quantitative science and it is always desired to know the relative masses of substances that react as pure elements or compounds. The measurement of masses of atoms is not possible because they are very small. Their masses can however be compared to give relative atomic and molecular masses of elements and compounds. The relative atomic mass of an element is the mass of one atom of the element compared to one-twelfth (1/12) of the mass of one atom of carbon - 12 isotope. On this scale 1 atom of carbon-12 isotope is given a mass of 12 atomic mass units. With the use of the mass spectrometer it has been possible to determine fairly accurately the relative atomic masses of elements. The relative molecular mass of a compound is the sum of the relative atomic masses of the elements present in the chemical formula of the compound. 13.5 Chemical symbols and formulae Chemical symbols and formulae are abbreviations used to represent elements and compounds. A chemical equation uses these symbols and formulae to summarise a chemical reaction. Chemical symbols consist of the first one or two letters of the name of the element. Some symbols do not correspond with the elements names; these symbols are derived from the Latin names of the elements. It is important that you know the symbols for as many of the common elements as possible. Table1.3 gives some examples of chemical symbols of elements. Table 1.3 Elements, symbols and relative atomic masses Symbol Relative atomic mass Element Hydrogen H 1.00797 Oxygen 0 15.9998 Chlorine Cl 35.453 Sodium Na 22.990 Iron Fe 55.847 Magnesium Mg 24.305 Zinc Zn 65.38 Cobalt Co 58.933 Copper Cu (cuprium) 64.456 Lead Pb (plumbuim) 207.19 Gold Au(Aurium) 196.967 4 Hydrogen, oxygen and chlorine elements exist in nature as diatomic molecules. i.e. units consisting of two atoms that are chemically bound together. The Latin names for copper, lead and gold are in bracket. The formula of a compound gives the proportion of the different elements present in the compound by mass. By the law of constant composition the proportion by mass of the different elements present is fixed for a pure sample of the compound irrespective of the method of preparations. Example 1.1 Analysis of 1.630g pure sample of a compound of calcium and sulphur gives 0.906g Ca and 0.724 sulphur. mass ofCa x 100.906 x 100 % Ca — = 55.6% mass of sample 1.630 mass of sulphur Yo S x 100 ( mass of sample.724 = 44.4% 1.630 Relative mass of Ca and S in the compound 55.6 44.4 Relative atomic mass of Ca relative atomic mass of S. 55.6. 44.4 40 32 1.39 : 1.39 i.e. 1:1 The formula of the compound is CaS. Having determined the relative atomic mass of the elements that make up the compound, you can proceed to determine the empirical formula of the compound. Since the relative masses of the elements in a compound depend partly on the masses of the atoms and also on the relative number of each atom of each element involved in the combination to form the compound. Example 1.2 2.83g of a compound of lead (Pb) and surphur gives 2.45g lead and 0.38g sulphur. Determine the formula of the compound. 2.45 x 100 = 86.6 "1/0 P b = 2.83 0.38 S= x 100 = 13.4 2.83 5  ▪ Relative number of Pb : S 86.6 3.4 207 32 0.418 :.419 i.e 1 :1 Formula of compound is PbS. Now determine a formula of a compound of calcium, carbon and oxygen with 40 percent Ca, and 48 percent oxygen. 1.3.6 Other laws of chemical combination In the previous section, the law of constant composition (proportion) is stated. Two more laws will be discussed in this section. (a) The law of conservation of matter This states that matter is neither created nor destroyed in a chemical reaction. A consequence of this law is that a chemical equation must always be balanced to account for all atoms present on the reactant side, and on the product side of the reaction. For example 2E12 02 —> 2H20 2Na C12 —> 2NaC1 H2 Cl2 -' 2HC1 2C0 + 02 2CO2 (b) The law of multiple proportions This states that when two different compounds are formed from the same two elements e.g. (CO and CO 2 or SO2 and S03 ) the masses of one element which react with a fixed mass of the other are in a ratio of small whole numbers. Example 1.3 Analysis of two compounds of carbon shows that 1.33g oxygen combines with 1.00g of carbon, while in the second 6.64g oxygen combines with 2.49g of carbon. Show that this is in agreement with the law of multiple proportion. For the 1st compound 1.00g carbon combines with 1.33g oxygen For the second compound 2.49g carbon combines with 6.64g oxygen 2.49g 6.64 carbon will combine with 2.49g g oxygen 2.49g 6  1.00g carbon will combine with 2.667g oxygen. Ratio of 0 mass combining with 1.0gC = 1.33: 2.667 1 : 2 simple ratio. 1.4 Chemical Reactions and Equations A chemical reaction involves the reshuffling of atoms from one set of combinations to another. In the reshuffling, one compound is converted to another. The smallest units that can take part in chemical reactions are the atoms and molecules. A chemical equation is often used to summarise the reaction that has taken place. A chemical equation gives the reactants and products of the reaction and the quantities of the reactants and products in correct ratio in accordance with the law of conservation of matter. Sometimes chemical equations give the physical states of the reactants and products. (s) for solid (1) for liquid (g) for gas (aq) for aqueous (solution in water) The equation does not give you the rate of the reaction, the energy effect and the conditions necessary for the reaction to occur. 1.5 Conclusion A chemical reaction involves elements, compounds and mixtures and leads to rearrangement or reshuffling of atoms. The elementary particles that form the basic units of elements and compounds in chemical reactions are atoms and molecules. Though atoms are composed of more fundamental particles, they are not split in chemical reactions. Dalton's atomic theory is the basis of modem atomic theory and explains satisfactorily the laws that govern chemical reactions of elements and compounds. Atoms and molecules are the building blocks of matter. The atom has a substructure of its own. This is the subject of the next unit. 1.6 Summary The concept of atoms and molecules is introduced and explained. Dalton's atomic theory is discussed and the modification in the light of new experimental evidences highlighted. The atom has a substructure of its own and it is electrical in nature. In chemical reactions, atoms are not split but are reshuffled. They are neither created nor destroyed. Pure compounds contain fixed proportions of elements by mass. A chemical equation is a summary of the reaction. Chemical symbols and formulae allow for summary of a reaction in short hand form. A chemical equation must be balanced for the law of conservation energy to be satisfied. Chemistry is a quantitative science and as such, amounts of matter used in reactions must be known. relative atomic and molecular masses allow us to see on the atomic level. the relative amounts of to products in chemical reactions. -P (c) Calculate the percentage by mass of hydrogen in the following compounds: (i) water (H2 O) (ii) Hydrogen sulphide (H 2 S) (iii) hydrogen chloride (HO) and (iv) methane (CH4) 2. (a) What are isotopes? (b) Natural chlorine consists of 75 percent of chlorine-35 and 25 percent of chlorine-37 isotopes. Calculate the relative atomic mass of chlorine element. (c) Explain why matter is neutral though it contains negative particles in its structure. 1.8 References Bajah, S T., Teibo, B. 0. ; Onwu, G and Obikwere, A. (2002). Senior Chemistry Textbooks I and 2. Lagos Longman Publishers. Osei Yaw Ababio (2002). New School Chemistry. Onitsha Africana-Fep Publishers. 8  Unit 2 Electronic Configuration 1— Static Model 2.0 Introduction In the last unit the concept of atoms and molecules was introduced and discussed. Dalton's theory proposed that atoms are indivisible units of matter. The atom is the smallest unit of matter that can take part in a chemical change. Dalton's atomic theory satisfactorily explained the laws of chemical combination but could not explain why substances react the way they do. Why is oxygen able to react with maximum of two atoms of hydrogen as in water? Why do some elements exist only as diatomic molecules? Why are some elements very reactive and some inert? Dalton's law could not explain electrolysis neither could it explain the different masses of atoms of the same element. Today we believe that the atom has a substructure of its own. The atom consists of much smaller particles that we call protons, neutrons and electrons. The relative masses and charges of these particles are given in unit 1. What are the evidences in support of this new picture of the atom? How many of these particles are present in the atom of elements? How are these particles arranged in the atom? These are some of the questions that will be answered in this unit. 2.1 Objectives At the end of this unit, you should be able to: Discuss the scientific evidences for the electrical nature of the atom. Give the number of subatomic particles in atoms of given elements. Recall the relative masses and charges of the subatomic particles. Define isotope Explain the nuclear model of the atom proposed by Rutherford. State the limitations of Rutherford's theory. Explain the origin of electronic energy levels in atoms. Write electronic shell configuration for elements and ions. 2.2 The Electrical Nature of the Atom 2.2.1 Historical evidences A very early evidence for the electrical nature of the atom came from Faraday. The result of Faraday's experiment on electrolysis showed that chemical change could be caused by the passage of electricity through 9 aqueous solutions of chemical compounds. This evidence was closely followed by the discharge tube experiment. A heated metal cathode emitted negatively charged particles. This beam of particles is called cathode rays and the particles are the electrons. J.J. Thompson worked on cathode rays and confirmed that they are negatively charged. Their charge mass ratio was determined and found to be —1.76 x coulomb/g'. R. A. Millikan determined the electronic charge in his famous oil drop experiment in 1908. The charge of the electron is —1.602 x 10 19 coulomb. The mass of the electron was calculated. It is 9.11 x 10 -28g. Fig. 2.1: Millikan oil drop experiment The atom is electrically neutral, so it is reasonable to expect positively charged particles in the atom. The existence of positively charged particles was confirmed in a modified version of the set up used by J. J. Thompson by another scientist by name Goldstein. The charge/mass ratio of the positive particles are much smaller than for the electrons. The largest charge/mass was for hydrogen ion (H1 which now represents the fundamental particle of positive charge in the atom. The proton is about 1835 times as heavy as the electron and carry a charge equal but of opposite sign to that of the electron. Protons and electrons are present in all atoms. The evidence of radioactivity by H. Becquerel further demonstrated the existence of subatomic particles. Becquerel observed that certain substances spontaneously emit radiations. The most important of these radiations are the alpha, beta and gamma radiations. Chadwick later confirmed the existence of a neutral particle in the atom and called this a neutron. This neutron has a mass approximately equal to that of the proton. Table 2.1: Properties of subatomic particles Particles Mass Charge Grams a.mu Coulombs electron charge Proton 1.67x1 0-24 1.007274 +1.602x10-19 +1 Neutron 1.68x10-24 1.008665 0 Electron 9.11x10-28.000549 —1.602x10-' 9 —1 2.2.2 Atomic and mass numbers Atoms of different elements have varying numbers of protons, electrons and neutrons. The atomic number is the number of protons in an atom of the element and for a neutral atom, the atomic number is also the number of electrons. The sum of the number of protons and neutrons is the mass number. Isotopes are atoms of the same element with different mass numbers. We shall see in the subsequent units that the chemical property of an element depends on the number as well as the arrangement of the atomic electrons. This explains why atoms of the same element with different masses have the same chemical properties. I2 C and 14c These are isotopes of carbon, mass number 12 and 14 and neutron numbers 6, 6 6 and8respctivly. 10 23 Atomic Models 23.1 The nuclear atom The presence of subatomic particles posed challenge to early scientists. It was necessary to propose a model for the atom. J. J. Thompson proposed that the atom could be viewed as positive matter in which electrons are uniformly distributed to make it neutral at every point. This view was dropped because of the findings of two other scientists, Rutherford and Marsden. They bombarded a thin gold foil with fast moving alpha particles. They found that most of the alpha particles pass through the foil undeflected. Some were deflected as large angles while very few were sent back on their paths. Alpha particles are positively charged and many of them passing through the foil undeflected suggested that most of the gold foil was empty space. (a) (b) Fig. 2.2: Scattering ofa-particles by a metal foil. (b) Deflection ofd particles by atomic nuclei Rutherford, using the findings in the above experiment proposed a model for the atom. He proposed that the atom consisted of a tiny positively charged nucleus. The nucleus is centrally placed in the atom and the electrons surround it. The very small number of deflections of alpha particles suggest that the nucleus occupies a very small portion of the atom. For heavy particles such as the alpha particle to be so deflected suggests that the nucleus is a centre of heavy mass and positive charge. The protons and neutrons occupy the nucleus while the electrons are arranged around the nucleus and move in orbits around it, as planets around the sun. This is Rutherfords nuclear model of the atom. By counting the number of alpha particles deflected in various directions, Rutherford was able to show that the diameter of a nucleus is about 1/100,000 times the diameter of the atom. Fig. 2.3: Schematic of a discharge tube. When a voltage is applied across electrodes that are sealed in a partially evacuated glass tube, the space between the electrodes glows. 11  2.3.2 Electronic energy levels A major problem of Rutherford's model of the atom is that the electron (—vely charged) rotating around the nucleus (+vely charged) will lose energy continuously because of the electrostatic force of attraction of the nucleus on the electrons. This is not observed in practice. Intact energy absorption and emission by elements is discontinuous. Have you ever seen the rainbow in the sky? The colours you see range from violet to red with no sharp line separating one colour from the other. This is a continuous spread of colours and is called a continuous spectrum. The different colours are component colours of light. In the laboratory the separation of light into its component colours also happens when light passes through a glass prism. Light from the vapour of an element does not give a continuous spectrum. Each element has its own characteristic bright lines in particular positions. This is a line spectrum suggesting that light energy absorption or emissons by elements is only at particular energies characteristic of the element. On the basis of the above observations Niel's Bohr proposed a model for the atom in which electrons move round the nucleus only in allowed orbits numbered serially. The orbit closest to the nucleus is assigned number 1 and is the orbit of lowest energy. Allowed transitions are transitions from one orbit to another and will lead to emission or absorption of the energy difference between the orbits. In the light of Bohr's model, there are electronic energy levels in atoms corresponding to different orbits of electron motion that are allowed. These energy levels are sometimes referred to as electron shells and designated K, L, M, N etc. corresponding to orbit numbers 1, 2, 3, 4, etc respectively. Bohr's model gave satisfactory explanation of the hydrogen spectrum. The theory is limited however in its explanation for multi electron atoms. The wave mechanical treatment of the atom overcomes this limitation of Bohr's theory and is the subject of a subsequent unit. Fig. 2.4: The radii of Bohr 1. orbits. 2.33 Electronic configuration Electronic configuration gives the arrangement of electrons in the energy levels in the atom. In the ground state (most stable state of lowest energy) of the atom, electron assignment to energy levels is according to the following rules: 12  the order of filling is K , L , M.N etc. The first shell is filled first before filling the second the K shell can accommodate maximum of 2 electrons the L shell can accommodate maximum of 8 electrons. the M shell accommodates 18 electrons. When 8 electrons are accommodated