Year 13 Chemistry Textbook.pdf

CHEMISTRY FOR YEAR 13 Curriculum Advisory Services Ministry of Education, Heritage & Arts Fiji 2018 MEHA CHEMISTRY FOR YEAR 13 Page i The CAS Section of the Ministry of Education, Heritage and Arts owns the copyright to this text book, Chemistry for Year 13, which is based on the Year 13 Chemistry Syllabus 2018. Schools may reproduce this in part or in full for classroom purposes only. Acknowledgement of the CAS Section of the Ministry of Education, Heritage and Arts copyright must be included on any reproductions. Any other use of this book must be referred to the Permanent Secretary for Education, through the Director CAS, Ministry of Education, Heritage and Arts, Fiji. Ministry of Education, Heritage & Arts, Fiji, 2018 Published by Ministry of Education, Heritage and Arts Marela House Private Mail Bag Suva Fiji Tel: (679) 3313050 Website: www.education.gov.fj MEHA CHEMISTRY FOR YEAR 13 Page ii ACKNOWLEDGEMENT Chemistry for Year 13 textbook has been produced for use at the Year 13 level by the Curriculum Advisory Services (CAS) of the Ministry of Education, Heritage and Arts. The following CAS officers are acknowledged for their contribution to the development of this textbook: Mr. Sunny Prasad Senior Education Officer [Chemistry] Miss. Poonam Chand Research Officer [Chemistry] Miss. Tejal Maharaj Research Officer [Chemistry] Acknowledgement is also extended to:  Mr. Vimlesh Chand, the Director – Curriculum Advisory Services, for the final feedback, comments and proof reading of this textbook.  Prof. Surendra Prasad (USP), Dr. David Rohindra (USP), Dr. Romila Gopalan (USP), Dr. Francis Mani (USP) and Ms. Veena Bilimoria (USP), who were consulted before finalising the textbook.  Prof. Rajendra Prasad (FNU), who was also consulted before finalising the textbook. The members of the Chemistry Curriculum Workgroup Committee (2017) are also acknowledged for assisting in the vetting of this textbook. The members consisted of Ms. Tarisi Tawake (AOG High School), Ms. Shradha Prasad (Kalabu Secondary School), Ms. Sahindra Kumar (DAV Girls College), Ms. Shaleni Kiran Prasad (Sacred Heart College), Mr. Salendra Lal (Ballantine Memorial School), Ms. Ashmir Ali (RSMS), Ms. Rahida Begum (Suva Muslim College), Ms. Seruwaia Koto (AOG High School), Mr. Atinesh Kumar (Vunimono High School), Ms. Miliana Lawa Savua (Nasinu Secondary School), Ms. Shelin Prasad (Pt. Shreedhar College) and Mr. Ronil Singh (William Cross College). MEHA CHEMISTRY FOR YEAR 13 Page iii PREFACE Welcome to Chemistry for Year 13. This book was prepared as a course material for Year 13 Chemistry and related activities. It covers the Year 13 Chemistry Syllabus 2018. The syllabus is the framework for the learning and teaching of Chemistry. This textbook is a resource for students and teachers and it is hoped that it will be useful in implementing the syllabus and to complement lessons prepared by teachers. Students and teachers are encouraged to use other resource materials as well. Suggestions for amendments are welcome. This is an amended copy of the Chemistry for Year 13 textbook which was developed in 2018. MINISTRY OF EDUCATION, HERITAGE & ARTS SUVA 24th March, 2020 MEHA CHEMISTRY FOR YEAR 13 Page iv CONTENTS Strand Title Page No. 1 General Chemistry……………………………………… 1 1.1 Scientific Skills…………………………………………….…………. 2 Graphs………………………………………………………………….… 2 Line graphs…………………………………………………….… 2 Relationship and trend from line graphs…………...……. 5 Making estimates of unmeasured data……………………. 7 Safety in the laboratory……………………………………….…….. 10 Chemical hazards………………………………………..…….. 12 Safety measures of hazardous substances……..………… 13 Disposal of wastes……………………………………………… 15 Safety Data Sheet (SDS)………………………………….…… 16 Experimental techniques………………….………….……………… 17 Reflux………………………………….………………………….. 17 Difference between reflux and distillation….………… 18 Melting point determination……………………….………… 20 1.2 Green Chemistry……………………………………………………… 22 Sustainable development………………...………………………….. 22 Green chemistry…………………………………………………….… 22 Atom economy……………………………………………………….... 23 2 Investigating Matter……………………………………. 25 2.1 Atomic Structure and Bonding…………………………………... 26 Introduction…………………………………………………………..… 27 Calculation of Relative Atomic Mass………………………………. 28 Electron configuration………………………………………………... 29 Quantum numbers………………………………………………. 31 Filling of orbitals…………………………………………………. 33 Electron configuration of ions………………………………… 37 Abbreviated electron configuration……………………….…. 38 Trends in the periodic table……………………………………….… 40 Atomic Radii………………………………………………………. 40 Ionic Radii…………………………………………………………. 41 Ionisation Energy………………………………………………… 42 MEHA CHEMISTRY FOR YEAR 13 Page v Electronegativity……………………………………………..….. 44 Chemical Bonding………………………………………………..…… 45 Ionic bonding…………………………………………………….. 45 Covalent bonding……………………………………………..… 46 Lewis structure diagrams…………………………………….. 46 Predicting molecular geometry………………………………. 49 Sigma (σ) and Pi (π) Bond Formation…………………….... 52 Polarity of Molecules…………………………………………… 54 Dative (Co-ordinate Covalent) Bonding……………………. 56 Inter-Molecular Attractions………………………………….. 57 Dipole-dipole interactions………………………….….… 58 Hydrogen bonding…………………………………………. 58 Induced-dipole interactions…………………………..... 59 Ion-dipole interactions…………………………………… 60 Physical properties due to intermolecular attractions… 61 2.2 States of Matter…………………………………………………....... 64 Introduction…………………………………………………………….. 65 Gases……………………………………………………………………... 67 The Kinetic Theory of Gases………………………………….. 67 Ideal Gas and Real Gas……………………………………….. 67 Important Gas Laws……………………………………….…… 67 Boyle’s Law……………………………………………..…… 67 Charle’s Law………………………………………………… 69 Dalton’s Law of Partial Pressure……………………….. 72 Avogadro’s Principle……………………………………….. 73 Gay-Lussac’s Law of Combining Volumes……..…….. 74 Combined Gas Law………………………………………… 74 Ideal Gas Law……………………………………………….. 75 Liquids and Solutions…………………………………………...…… 78 The Kinetic-Molecular Description of Liquids……………. 78 Solutions………………………………………………………….. 79 Mole fraction…………………………………………………. 79 Molarity…………………………………………………….…. 80 Molality………………………………………………….……. 80 Weight Percent (w/w %)………………………………...... 81 Weight/Volume Percent (w/V %)……………………….. 82 Volume percent (V/V %)…………………………………. 83 MEHA CHEMISTRY FOR YEAR 13 Page vi 3 Physical Chemistry…………………….………….…... 85 3.1 Electrochemistry……..……………………………………….…….. 86 Redox reactions………………………………………………………... 87 Oxidation number……………………………………………… 89 Oxidants and Reductants…………………………………….. 90 Balancing redox-equations (ion-electron method)…….… 96 Acidic Solutions……………………………………………. 96 Basic Solutions…………………………………………….. 98 Redox titration………………………………………………….. 100 Limiting reagent………………………………………………... 103 Electrochemical cells………………………………………………….. 105 Cell notation…………………………………………………….. 109 Cell potential /Standard cell potential……………………. 112 Standard reduction potential………………………………… 112 Calculating standard cell potential………………………… 115 Spontaneity of a reaction……………………………………... 120 3.2 Thermochemistry……………………………………….……………. 124 System and Surrounding…………………………………….. 125 Heat of reactions………………………………………………... 125 Enthalpy……………………………………………………….…. 128 Calorimetry………………………………………………….…… 132 Standard Heat of Reactions…………………………….……. 136 Thermochemical equations……………………………….…. 136 Calculating Standard Heat of Reaction using Standard Heat of Formation…………………………………. 137 Hess’s Law…………………………………………………..…… 138 Bond Energy………………………………………………….…. 144 3.3 Aqueous Solution Chemistry……….……………………………. 150 Chemical equilibrium……………………………………….…. 151 Equilibrium constant……………………………………….…. 152 Acids and Bases…………………………………….………….. 159 Dissociation constant of acids and pKa ………………….. 162 Base dissociation constant and pKb………………………. 165 pH and pOH Calculations…………………………………….. 169 pH calculations for strong acids and bases…….….… 171 pH calculations for weak acids and bases………..….. 173 Acid base titration……………………………………………… 180 MEHA CHEMISTRY FOR YEAR 13 Page vii Titration curves…………………………………………….. 180 Acid – base indicators………………………………….…. 184 Buffer solutions……………………………………………….… 187 Solubility and precipitation reactions……………………… 188 Solubility Product ……………………………………….…….. 189 Solubility Product Expressions……………………………… 190 Relationship between solubility and Ksp………….……….. 191 Ionic Product …………………………………………….……... 196 The Common - Ion Effect ………………………………..…… 198 4 Materials…………………………………………………... 200 4.1 Inorganic Chemistry………………………………………………… 201 Hydrides………………………………………………………………….. 202 Ionic hydrides……………………………………………………. 202 Covalent hydrides………………………………………………. 204 General trends in the properties of hydrides…………….. 206 Oxides…………………………………………………………………….. 211 Chlorides…………………………………………………………………. 214 Transition metals………………………………………………………. 216 Electron configuration of transition metals…………….… 216 Properties of transition metals…………………………….… 220 Oxidation states…………………………………………… 220 Coloured compounds………………………………….…. 221 Catalysts…………………………………………………….. 222 Paramagnetism……………………………………….……. 223 Complex ions………………………………………….……. 223 Formula of complex ions…………………….……. 224 Naming of complex ions……………………….….. 225 Writing the formula of complex ions…………… 230 4.2 Organic Chemistry……….………………………………………….. 235 Isomerism………………………………………………….…………….. 236 Structural isomerism………………………………….………. 236 Stereoisomerism…………………………………..……………. 238 Hydrocarbons………………………………………………….…….…. 243 Aliphatic hydrocarbons………………………….................. 243 Aromatic hydrocarbons……………………………………….. 245 Alkyl halides…………………………………………………………..… 248 Physical properties of alkyl halides……………………..…. 250 MEHA CHEMISTRY FOR YEAR 13 Page viii Chemical reactivity of alkyl halides………………………… 252 Reactions of alkyl halides…………………………………….. 253 Substitution reactions……………………..…….………. 253 Elimination reactions…………………………………….. 255 Amines……………………………………………………………..……. 257 IUPAC Nomenclature of primary amines………………….. 259 Physical properties of amines……………………………….. 260 Making amines from alkyl halides…………………….……. 260 Basicity of amines………………………………………….…… 260 Basicity of ammonia and primary amines…………... 261 Alcohols……………………………………………………………….….. 263 Properties of alcohols……………………………………….…. 267 Reactions of alcohols………………………………………….. 268 Test to distinguish between alcohols………………………. 271 Aromatic alcohols (Phenol)…………………………….…….. 272 Functional group isomers of alcohols (ethers)…………… 273 Aldehydes and Ketones…………………………………………….…. 277 Aldehydes………………………………………………………... 278 Physical properties of aldehydes………………..….. 278 Ketones…………………………………………………………... 279 Physical properties of ketones……………………….. 279 Tests to distinguish between aldehydes and ketones….. 281 Carboxylic acids…………………………………………………….…. 285 Physical properties of carboxylic acids…………………… 286 Preparation of carboxylic acids…………………………….. 287 Oxidation of primary alcohols…………………………. 287 Oxidation of aldehydes………………………………….. 288 Reactions of carboxylic acids………………………………. 289 Reactions with metals………………………………….. 289 Reactions with metal hydroxides…………………….. 290 Reactions with carbonates and bicarbonates……… 290 Reduction of carboxylic acids…………………………. 291 Carboxylic acid derivatives…………………………………………… 293 Amides…………………………………………………………….. 293 Preparing amides from carboxylic acids…………….. 294 Acid anhydrides………………………………………………… 295 Preparing acid anhydrides from carboxylic acids…. 295 MEHA CHEMISTRY FOR YEAR 13 Page ix Acid chlorides………………………………………………..… 296 Preparing acid chlorides from carboxylic acids…... 297 Esters……………………………………………………………………... 298 Physical properties of esters…………………………………. 300 Preparation of esters……………………………………….….. 301 Reactions of esters………………………………………….….. 302 Priority of organic functional groups……………………….. 303 5 Consumer Chemistry…………………..……………….. 307 Polymers……………………………………………………………….….. 308 Natural polymers…………………………………………….…. 308 Synthetic polymers……………………………………….….… 308 Properties of polymers…………………………………………. 309 Condensation polymers……………………………………………..… 310 Addition polymers………………………………………………………. 311 Biodegradable and non-biodegradable polymers………………... 313 Problems associated with non-biodegradable polymers……….. 314 Bibliography…………………………………………………………………………... 316 Glossary………………………………………………………………………………….. 317 Appendices Appendix 1: The Periodic Table……..………………….……………………… 327 Appendix 2: Safety Data Sheet (SDS).……..………………………………… 328 Appendix 3: Possible Quantum Number Values…………………………… 334 Appendix 4: Average Bond Energies………..………………………………… 335 Appendix 5: Summary of Formulae.………..………………………………… 336 MEHA CHEMISTRY FOR YEAR 13 Page x STRAND 11 STRAND GENERAL GENERAL CHEMISTRY CHEMISTRY Adapted from: https://play.google.com STRAND OUTCOME Demonstrate knowledge and skills on the experimental techniques and understand the importance of green chemistry. SUB – STRANDS 1.1: Scientific Skills 1.2: Green Chemistry CHEMISTRY FOR YEAR 13 STRAND 1 Page 1 1.1 Scientific Skills Achievement Indicators: Upon completion of this sub-strand, students will be able to:  Represent data in meaningful and useful ways.  Organise and process data to identify trends, patterns and relationships.  Demonstrate laboratory safety practices during experiments.  Describe the hazards of chemical substances.  Describe the proper disposal of used or waste chemicals.  Explain the importance of a Safety Data Sheet (SDS).  Examine the importance and use of some laboratory techniques. Graphs Graphs are often the best way to represent numerical data. This is because graphs make it easy to read, compare and identify trends in the data. Graphs can also be used to make predictions. There are many different types of graphs which are used to present the different types of statistical information. However, in Chemistry, line graphs are most commonly used. Line graphs Line graphs display data that changes continuously over time. Line graphs are drawn by plotting points by their x and y coordinates, then joining them together or drawing a line through the middle. X-axis  This is the horizontal axis (runs across the bottom).  The independent variable goes in the x-axis. These are the values that can be changed/manipulated in the experiment to give an output (dependent variable).  In other words, the independent variable is what causes the results in an experiment. Y-axis  This is the vertical axis (runs up the left side).  The dependent variable goes in the y-axis. These are the values that result from the independent variables (input). CHEMISTRY FOR YEAR 13 STRAND 1 Page 2  This is the variable you want to see if it was affected or not in the experiment.  In other words, the dependent variable is what is obtained from the experiment. Note: 1. Variables are things that vary and can be changed. 2. When drawing tables for the results, the independent variable goes in the first column. 3. It is important to show the correct units of the variables while labelling the axis. Example of Independent and Dependent Variables in Experiments 1. When the effect of altering the temperature of a gas on the volume is studied:  The independent variable is temperature.  The dependent variable is volume. Example y Volume (cm3) (Dependent variable) Temperature (oC) x (Independent variable) 2. An experimenter decides to determine the effect of temperature on rates of reaction.  The independent variable is temperature.  The dependent variable is the rate of reactions. 3. An experimenter studies the impact of a drug on cancer causing cells.  The independent variable is the dosage of the drug.  The dependent variable is the impact of the drug on cancer causing cells. CHEMISTRY FOR YEAR 13 STRAND 1 Page 3 Constructing a Line Graph 1. Put the independent variable on the horizontal (x-axis) and the responding dependent variable on the vertical (y-axis). 2. Divide each axis into equal segments. Number the segment marks so that they include the complete range of values of the variables. 3. Label each axis, including the units of measurement (example: cm, mL, g mol-1 and ℃). 4. Put a dot at the location of each pair of variable values. 5. If the data points lie in a straight line or in a smooth curve, draw a line that goes through each point. 6. If the data points are scattered, you must decide whether you are most interested in the individual fluctuations of the data, or in the overall trend. If you want to show the fluctuations, then connect each data point to the next with a single straight line. If you want to show the overall trend, then draw a line of best fit. A line of best fit is a straight line or a smooth curve that averages the data points. It goes through the collections of data points so that there are approximately equal number of points on each side of the line. Also a line of best fit passes through or is close to as many points as possible. 7. Give the graph a descriptive title. An example of a graph showing the line of best fit Graph of volume vs mass of a substance y Data points Volume (cm3) line of best fit Mass (g) x Note: Computer softwares, such as Microsoft Excel, are now commonly used to analyse data in a more efficient way. CHEMISTRY FOR YEAR 13 STRAND 1 Page 4 Relationship and Trends obtained from Line Graphs 1. Directly proportional (Positive Linear Relationship)  Such graphs show that the two variables are directly related. Example Graph of volume of gas vs temperature y Volume (cm3) Temperature (oC) x The above graph shows that the volume of a gas increases as the temperature is increased. 2. Inversely proportional (Negative Linear Relationship)  Such graphs show that the two variables are inversely related. Example Graph of density of water vs temperature y Density (g cm-3) Temperture (oC) x The above graph shows that the density of water decreases as the temperature is increased. CHEMISTRY FOR YEAR 13 STRAND 1 Page 5 Interpreting Line Graphs Interpreting a graph means to explain in words what the graph shows. The following format is frequently used. A. Describe what happens to the dependent (responding) variable as the independent (manipulated) variable changes. Example Graph of pressure in a container vs temperature y Pressure (atm) Temperature (oC) x The above graph shows that the pressure in the container increases as the temperature is increased. The relationship between the temperature and pressure is positive linear (directly proportional). B. If the relationship between the variables is not linear (if the data points fall on a curved line), then the relationship is described in two parts: 1. The relationship is described until the curve changes direction. 2. The relationship is described for the rest of the curve. Example Graph showing the effect of fertiliser on plant height y Plant height(cm) 10 20 x -1 Amount of Fertiliser (g L ) The above graph shows that the plant height increased as the amount of fertiliser was increased until the amount of fertiliser reached 10 g L-1. Then, the plant height rapidly decreased as more fertiliser was added until at 20 g L-1, all plants died. CHEMISTRY FOR YEAR 13 STRAND 1 Page 6 Making estimates of unmeasured data Graphs are useful in Chemistry because they allow us to make predictions. Predictions can be made about what happens: (a) Between two known points on the graph (interpolation). (b) Before the first known point on the graph (extrapolation). (c) After the last known point on the graph (extrapolation). Interpolation Interpolation is estimating unmeasured data points within the range of data plotted on the graph. This means to insert points between known points on the graph. The value of a point on the graph that occurs between two known values on the same graph is known as interpolation. An assumption of interpolation is that the overall relationship described for the known points is also true between known points. These lines are drawn as solid lines between plotted points. Extrapolation Extrapolation is predicting beyond existing data. This happens by inserting points either before the first known point or after the last known point on the graph. Extrapolating means extending the line of best fit past the last known point. It is important that the known parts of the graph is distinguished from the extrapolated parts of the graph. Therefore, these lines are drawn as dotted lines (or sometimes dashed lines) beyond the known plotted points. Extrapolation also assumes that the overall relationship described for the known points is also true for points before the first known value and points after the last known value. CHEMISTRY FOR YEAR 13 STRAND 1 Page 7 Example 1 In the line segment AB given below, the point C is interpolated; while the points D and E are extrapolated by extending the straight line beyond AB. y Extrapolated, D Measured, B Interpolated, C Measured, A Extrapolated, E x This means that the data of points A and B were measured in an experiment and data for points C, D and E were not measured, but can be predicted using the available data. Example 2 Jacques Charles carried out an experiment to determine the relationship between temperature and volume of a gas. The results obtained is shown below. Temperature Volume of (°C) Gas (cm3) 20 60 40 65 60 70 80 75 100 80 120 85 a. Using the above data, draw a graph to show the relationship between temperature and volume of gas. b. Predict the volume occupied by the gas when the temperature is 11 ℃ and 50 ℃, respectively. CHEMISTRY FOR YEAR 13 STRAND 1 Page 8 Solution a. Graph showing the relationship between volume of a gas and temperature 90 80 70 Volume (cm3) 60 50 40 30 20 10 0 0 20 40 60 80 100 120 140 Temperature (℃) Thus, the above graph shows that the volume of the gas increases as the temperature is increased. This means that temperature and volume is directly related. Graph showing the relationship between volume of a b. gas and temperature Volume (cm3) Temperature (℃) According to the extrapolation and interpolation as shown above, the volume occupied at 11 ℃ and 50 ℃ are as follows. i. 11 ℃ → 58 cm3 ii. 50 ℃ → 68 cm3 CHEMISTRY FOR YEAR 13 STRAND 1 Page 9 Safety in the Laboratory Understanding chemistry requires laboratory work which utilises lots of harmful chemicals. Proper care and precaution must always be taken while working in a laboratory. However, accidental chemical exposures can still occur even with good care and safety precautions. For this reason, it is essential to look beyond the use of safety specs, face masks and other personal protective equipment. This is because exposure to a hazardous substance, especially corrosive substances, are critical. Delaying treatment, even for a few seconds may cause serious injury. Safety showers (also known as emergency showers) and eyewash fountains are a necessary backup to minimise the effects of accidental exposure to chemicals. Safety showers and eyewash fountains provide on-the-spot decontamination and flushes away hazardous substances that can cause injury. Safety showers can also be used effectively in extinguishing clothing fires or for flushing contaminants off clothing. Note: When safety showers and eyewash fountains are not available in the school laboratory you are working in, use the sink (tap water) in the laboratory immediately in case of any accident (such as chemical spillage). Quick Exercise: Pick out some risks in the picture below, how showing some students doing titration experiment in a lab. Source: http://www.snehinternationalschool.com CHEMISTRY FOR YEAR 13 STRAND 1 Page 10 Below is a table describing the Eyewash Fountain and Safety Showers Name Diagram and Explanation Eyewash fountain (Flushes away chemicals that enter the eyes or are splashed on the face) Picture: USP Chemistry Lab Source: www.indiamart.com How to Use 1. Remove the caps on the mouth of the sprayers. 2. Place your eyes directly over the water sprayers and firmly push the lever. 3. Continue this treatment for at least 15 minutes, even if you feel the irritation has subsided. Safety Showers (Flushes away chemicals that are spilled on the body) Picture: USP Chemistry Lab Source:www.arjunfiresafety.com How to Use 1. Stand directly under the shower. 2. Pull the lever immediately to flush the affected area with plenty water for 15 minutes. Protect the eyes from any further damages. 3. Remove all contaminated clothing, jewellery and shoes. CHEMISTRY FOR YEAR 13 STRAND 1 Page 11 Chemical Hazards Being familiarised with the chemical hazard symbols helps in having a better understanding of the safety aspects of any chemical being used. Old Hazard Symbols Source: www.chemistry.iiti.ac.in New International Hazard Symbols Adapted from: www.chemistry.iiti.ac.in CHEMISTRY FOR YEAR 13 STRAND 1 Page 12 Safety Measures of Some Hazardous Substances Name Safety Precaution Examples  Do not place them near open flames and - Ethanol spark-producing equipment.  Eliminate ignition sources (sparks, - Acetone smokes, flames, hot surfaces) when - Benzene working with flammable and combustible liquids. - Cyclohexane  Use the smallest amount of flammable liquid necessary in the work area. - Diethyl ether Flammable  Keep storage areas cool and dry.  Store flammable and combustible liquids away from incompatible materials (e.g. oxidisers).  Store, handle and use flammable and combustible liquids in well-ventilated areas.  Keep containers closed, except when in use.  Store corrosives in suitable labelled -Hydrochloric containers away from incompatible acid materials, in a cool, dry area. - Sulphuric acid  Use corrosives in well-ventilated areas.  Always use hand gloves, aprons and - Nitric acid safety specs. Corrosive  Inspect containers for damage or leaks - Acetic acid before handling. Never use containers that appear to be swollen. - Ammonium  Handle containers safely to avoid hydroxide damaging them. - Potassium  Dispense corrosives carefully and keep hydroxide containers closed when not in use.  Stir corrosives slowly and carefully into - Sodium cold water when the job requires mixing hydroxide corrosives and water.  Store, handle and use toxic materials in - Hydrogen well-ventilated areas away from cyanide combustible and other incompatible Toxic - Hydrogen materials. sulfide  Wear the appropriate personal protective equipment such as safety specs, apron - Mercury and gloves. CHEMISTRY FOR YEAR 13 STRAND 1 Page 13  Keep containers closed when not in use. - Lead  Keep only the smallest amounts possible in the work area. -Formaldehyde  Do not return contaminated or unused - Chloroform toxics back to the original container.  Handle and dispose of toxic wastes safely. - Methanol - Acetonitrile  Explosive chemicals decompose or burn - Ammonium very rapidly when subjected to shock or nitrate ignition. They produce large amounts of heat and gas when triggered and thus are - Nitrogen Explosive extremely dangerous. trichloride  Keep explosive chemicals out of sunlight and away from heat source. - Silver nitride  Be familiar with the chemical you are using.  Prevent any shock or ignition.  Familiarise yourself with oxidising - Bromine materials.  Store oxidising materials in suitable - Peroxides labelled containers, in a cool, dry place. (E.g. Hydrogen peroxide)  Avoid or eliminate ignition sources (sparks, smokes, flames, hot surfaces) - Permanganates when working with oxidising materials. (E.g. Potassium Oxidising  Store, handle and use oxidising materials permanganate) in well-ventilated areas away from combustible and other incompatible - Nitrates materials. - Nitric acid  Handle containers safely to avoid damaging them. - Chromates  Dispense oxidising materials carefully, and using compatible equipment and Dichromates containers. (E.g. Potassium  Keep containers closed when not in use. dichromate)  Use only the smallest amounts possible.  Do not return contaminated or unused oxidisers to the original container.  Handle and dispose of oxidising wastes safely. CHEMISTRY FOR YEAR 13 STRAND 1 Page 14 Disposal of wastes It is the clear responsibility of any person working in the lab to ensure the safe and correct disposal of all wastes produced during the course of their work. Below is a description of some ways in which waste materials can be disposed in the laboratory. 1. Poured in the sink (Washed down the drain)  Some chemicals can be poured in the sink after dilution with lots of water.  Dilute acids and alkalis must be neutralised first before pouring them in the sink.  Harmless soluble inorganic salts (including all drying agents such as CaCl2, MgSO4, NaCl, Na2SO4 and P2O5) can be dissolved and poured in the sink.  Alcohol containing salts can also be dissolved and poured in the sink. Note: These substances should NEVER be washed down a drain.  Compounds of the following elements: arsenic, barium, beryllium, boron, cadmium, chromium, cobalt, copper, lead, mercury, nickel, selenium, silver, tellurium, tin, titanium, uranium, vanadium and zinc.  Organohalogen, organophosphorus or organonitrogen pesticides and herbicides (Note: Organo means containing organic groups; such as bromomethane and chloroethane).  Cyanides (Note: Cyanides contain C≡N group).  Mineral oils and hydrocarbons.  Poisonous organosilicon compounds, metal phosphides and phosphorus elements.  Fluorides and nitrites. 2. Stored in waste bottles  All organic solvents such as cyclohexane and acetone, including water miscible ones such as alcohol.  Soluble organic wastes, including most organic solids.  Paraffin and mineral oils. Note: Halogenated wastes must be kept separate to other organic solvents. Example: Acetone and chloroform should be kept separately since their mixture can explode. CHEMISTRY FOR YEAR 13 STRAND 1 Page 15 3. Controlled waste  Broken laboratory glassware, sharp objects of metal or glass, fine powders (preferably inside a bottle or jar) and dirty sample tubes or other items lightly contaminated with chemicals should be properly and separately disposed in the lab. 4. Waste for special disposal  Poisons (but not cyanides) and other highly toxic chemicals.  Materials contaminated with mercury.  Carcinogenic solids. Note: Cyanide wastes must be placed in an appropriate waste bottle and the solution kept alkaline at all times. 5. Laboratory waste bins  Items in this category include dirty paper, plastic, rubber and wood. Safety Data Sheet (SDS) A Safety Data Sheet (SDS), previously called a Material Safety Data Sheet (MSDS), is a document that provides information on the properties of hazardous chemicals and how they affect health and safety either in a laboratory or any workplace. It is the responsibility of any chemical manufacturer, distributor or importer to provide SDS for each hazardous chemical to users to communicate information on the chemical hazards. An SDS includes information such as the:  identity of the chemical.  properties of the chemical.  physical, health and environmental hazards.  protective, safe handling and storage procedures.  disposal considerations.  safety precautions for handling, storing and transporting the chemical.  other important information for the particular chemical. Important: It is essential to read the SDS of a chemical before using it (especially when used for the first time), to get familiarised with the dangers associated with it. Note: An example of a SDS of sulphuric acid can be found in Appendix 2. This SDS is extracted from: www.teck.com/media/Products-Sulphuric-Acid-SDS-2015.pdf CHEMISTRY FOR YEAR 13 STRAND 1 Page 16 Experimental Techniques 1. Reflux When studying chemistry, especially organic chemistry, the experimental technique of ‘reflux’ is often used. Many organic chemical reactions take very long to complete and in order to speed up these reactions, heat is applied. However, organic compounds with low boiling points are usually volatile and if heated they will evaporate. The solution to this problem is to heat the reaction mixture under reflux. The term ‘reflux’ describes an arrangement in which a reaction is carried out in a boiling solvent with the vapour being condensed and returned to the reaction vessel. Refluxing is carried out when reactions need to be heated to give a reasonable yield of product in a reasonable time. The reactants for reflux experiments can be solid and liquid, or both liquids. The diagram below shows the basic set-up for refluxing Water Out Condenser Water In Heating bath Hot Plate Adapted from: http://www.eplantscience.com The condenser is always completely filled with water to ensure efficient cooling. The vapours, which are given off from the liquid reaction mixture, change from gas phase back to liquid phase due to heat loss. This then causes the liquid mixture to fall back into the round bottom flask. In this way, it is ensured that the chemical reaction involving organic compounds will give a higher yield of product. CHEMISTRY FOR YEAR 13 STRAND 1 Page 17 Difference between Distillation and Reflux Reflux and distillation are two chemistry laboratory techniques which involve boiling and condensing of a solution. Reflux helps complete a reaction and distillation separates components of a mixture. Some differences between reflux and distillation are summarised in the table shown below. Reflux Distillation Procedure  In reflux, the liquid is boiled,  A liquid is heated until it is but the vapour is allowed to boiling. The vapour which is condense and flow back into produced is condensed and the original flask. collected in a separate flask. Separation  When the condensed vapour is  Collecting the condensed allowed to flow back into the vapour in a second flask, original reaction flask, the allows you to separate out the hard to dissolve compounds individual components of a are easily dissolved. mixture.  This continual recycling of the  Separation by distillation solvent helps drive the works because each liquid reaction to completion. has a different boiling point. The liquid with the lowest boiling point will vaporise and condense first. Equipment  A reflux set up includes a  In distillation, the round round bottom flask connected bottom flask is connected to a directly to a condenser. y-adapter. The side arm of the y-adapter connects to the condenser.  The condenser connects to a receiving flask via a vacuum adapter. CHEMISTRY FOR YEAR 13 STRAND 1 Page 18 Set-up of a Distillation Thermometer Condenser Clamp Water outlet Water inlet Heater Adapted from: https://drugs-forum.com Note: Distillation can be used to separate a pure liquid from a mixture of liquids. It works when the liquids have different boiling points. Example Distillation is commonly used to separate ethanol from water. The mixture is heated in a flask. Ethanol has a lower boiling point than water so it evaporates first. The ethanol vapour is then cooled and condensed inside the condenser to form a pure liquid. The thermometer shows the boiling point of the pure ethanol liquid. When all the ethanol has evaporated from the solution, the temperature rises and the water evaporates. CHEMISTRY FOR YEAR 13 STRAND 1 Page 19 2. Determination of Melting Point The melting point of a substance is the temperature at which the material changes from a solid to a liquid state. The melting point is a physical property characteristic of a particular compound. A pure crystalline compound usually possesses a sharp melting point and it melts completely over a narrow temperature range of not more than 0.5-1.0 ℃, provided good technique is followed. The presence of even a small amount of impurity in a substance changes (decreases) its melting point by a few degrees and broaden the melting point temperature range. More impurities increase these effects. Therefore, determining the melting point is a simple and fast method to obtain a first impression of the purity of a substance. Also the melting point of a substance can be obtained and compared with the literature value to confirm its identity. The test is an important technique for gauging purity of organic and pharmaceutical compounds. There are many different types of melting point apparatus available these days. The diagram of a simple melting point apparatus is shown below Thermometer M.P tube with sample Observation window Light source Source: http://www.scottsmithonline.com CHEMISTRY FOR YEAR 13 STRAND 1 Page 20 Exercise 1. During the course of an experiment, a student determined the volume of gas released from an experiment at different temperatures and obtained the following results. Temperature (℃) Volume (cm3) 20 55 40 65 60 75 80 85 100 95 120 105 140 115 160 125 180 135 200 145 a. Identify the independent and the dependent variable in the above experiment. b. Draw a labelled line graph to show the relationship between temperature and volume. c. From your graph, determine the relationship between temperature and volume of a gas. d. Using your graph, predict the volume of gas at 45 ℃ and 220 ℃ respectively. 2. Briefly explain how you would dispose the following chemical wastes in the laboratory. a. Cyclohexane b. Ethanol c. Sodium chloride d. HCl solution 3. A student accidently splashed some chemical solution in her eyes. What would you immediately advise her to do (Note: An eyewash fountain and a safety shower was present in that laboratory). 4. Why is it important to read the SDS of a chemical before using it? 5. Refluxing is a common laboratory practice. What is the purpose of refluxing? 6. Give an importance of determining the melting point of a substance. 7. What happens to the melting point of a substance when some impurities are present in it? CHEMISTRY FOR YEAR 13 STRAND 1 Page 21 1.2 Green Chemistry Achievement Indicators: Upon completion of this sub-strand, students will be able to:  Explain the importance of sustainable development and green chemistry.  Describe the relationship between sustainable development and green chemistry.  Describe atom economy.  Calculate the atom economy of a reaction. Sustainable Development To help deal with increasing world population, the world’s economy needs to grow. However, economic growth is often linked to environmental pollution problems. The challenge is to develop in a way that meet the needs of the present generation without compromising the ability of future generations to meet their own needs. In other words, the economy should grow without causing a lot of environmental damage and wasting limited resources. This type of development is called ‘sustainable development’. One of the ways in which the chemical industry is working towards sustainable development is by using ‘Green Chemistry’. Green Chemistry Green chemistry is the use of chemical products and processes that minimises or eliminates the production of hazardous substances. Wherever possible, green chemistry utilises renewable raw materials. Green chemistry reduces waste materials, hazards, energy, cost and risk. One of the basic ideas of Green Chemistry is to prevent pollution and the production of hazardous materials instead of producing them and then cleaning them up. Green chemistry is safe, conserves raw materials and energy and more cost effective than conventional methods. There are three main ways to make chemical processes ‘greener’:  Redesign production methods to use different, less hazardous starting materials.  Use milder reaction conditions, better catalysts and less hazardous solvents.  Use production methods with fewer steps and higher atom economy. CHEMISTRY FOR YEAR 13 STRAND 1 Page 22 Applying Green Chemistry Example 1 A chemical reaction is such: A + B → C + D If the only desired product is C, an alternative to A or B must be found to avoid production of D. Example 2 Chlorination is used to disinfect water. However, during the process other harmful chlorinated compounds are formed. Therefore, an alternative is to use other oxidising agents which will form less harmful products. Atom Economy Atom economy of a chemical reaction measures the amount of reactants that become useful products. Higher atom economy of a reaction, means ‘greener’ processes involved. For instance, a reaction which has 100 % atom economy means that all the atoms in the reactants have been converted to the desired product. Calculating atom economy Molecular weight of the desired product Atom economy = x 100 Sum of the molecular weight of all substances produced Example 1 Determine the atom economy for making hydrogen gas by reacting zinc with hydrochloric acid. Solution Zn(s) + 2HCl (aq) → ZnCl2(aq) + H2(g) Mr of H2 = 2 Mr of ZnCl2 = 136 2 Atom economy = x 100 (136 + 2) = 1.45 % Note: Industrial processes need as high atom economy as possible, because this:  Reduces the production of unwanted products.  Makes the process more sustainable. CHEMISTRY FOR YEAR 13 STRAND 1 Page 23 Example 2 Determine the atom economy of the reaction that produces hydrogen from methane and steam. Solution Now looking at the atom CH4(g) + 2H2O(l) → CO2(g) + 4H2(g) economy of the reactions of Examples 1 and 2, Mr of H2 = 8 which one will you use to Mr of CO2 = 44 produce hydrogen gas? 8 Why? Atom economy = x 100 (44 + 8) = 15.38 % Exercise 1. Why is it better to prevent pollution and the production of hazardous materials than to produce them and then clean them up? 2. Explain why using a catalyst may make a chemical process ‘greener’? 3. The reaction equation for extracting iron from its ore using carbon is: 2Fe2O3(s) + 3C(s) → 4Fe(s) + 3CO2(g) Calculate the atom economy of this reaction. 4. Titanium (Ti) can be extracted from its ore by two different methods. The reaction equation for each method is shown below. i. TiO2(s) + 2Mg(s) → Ti(s) + 2MgO(s) ii. TiO2(s) → Ti(s) + O2(g) a. Calculate the atom economy for each reaction. b. Which method is ‘greener’? What else might you want to know before making a final decision? c. Using equation (ii), what is the atom economy if oxygen is the required product? 5. Alkanes can be cracked to form alkenes. Decane can be cracked to form two products: C10H22(l) → C2H4(g) + C8H18(l) a. Calculate the atom economy of this process if only the alkene is required and can be sold? b. Calculate the atom economy if both products can be sold? c. Explain why your answers to (a) and (b) are different. 6. The key reaction in the Haber process for making ammonia is: N2(g) + 3H2(g) ⇋ 2NH3(g). Calculate the atom economy of this reaction. 7. Explain why using reactions with high atom economy is important for sustainable development. CHEMISTRY FOR YEAR 13 STRAND 1 Page 24 STRAND 2 INVESTIGATING MATTER STRAND OUTCOME Demonstrate an understanding of the differences in structure, bonding and properties of the different types of matter. SUB-STRANDS 2.1: ATOMIC STRUCTURE AND BONDING 2.2: STATES OF MATTER CHEMISTRY FOR YEAR 13 STRAND 2 Page 25 2.1 Atomic Structure and Bonding Atomic Structure and Bonding looks at the inner structures of atoms and explains why atoms combine to form compounds. To understand about atomic structure and bonding, emphasis will be made on electron configuration, quantum numbers, periodic trends, Lewis structures, molecular geometry and different types of bonds. Achievement indicators Upon completion of this sub-strand, students will be able to:  Calculate the relative atomic mass of isotopes.  Describe electron configuration of an atom.  Write electron configuration using s, p, d notation for the first 30 elements and their ions.  Describe the four quantum numbers.  Determine the four quantum numbers of an electron.  Explain the Aufbau Principle, the Pauli Exclusion Principle, and the Hund’s rule.  Apply the rules for assigning electrons to sub-shells and draw orbital diagrams.  Describe the trends in atomic radii across the period and down the group in a Periodic Table.  Compare the atomic radius with the ionic radius of a respective element.  Describe the trends in 1st ionisation energies across the period and down the group in a Periodic Table.  Compare the 1st, 2nd and 3rd ionisation energies of a respective element.  Draw the Lewis structures of some molecules and polyatomic ions.  Determine shapes of molecules and ions using Lewis structure and VSEPR theory.  Explain the formation of sigma and pi bonds and determine the number of sigma and pi bonds present in a compound.  Determine the polarity of molecules and relate this to their degree of electronegativity.  Define dative bonding and explain with diagrams how dative bonds are formed.  Describe intermolecular attractions between molecules or ions.  Describe the physical properties due to the intermolecular forces. CHEMISTRY FOR YEAR 13 STRAND 2 Page 26 Introduction Atoms are the basic units of matter and the defining structure of elements. All atoms are made up of three sub-atomic particles: protons, neutrons and electrons. Protons and neutrons are present in the centre of the atom, which forms the nucleus. The protons and neutrons together make up most of the mass of the atom. Electrons are extremely lightweight and exist in a cloud orbiting the nucleus as shown below. Source: http://chemistry.tutorcircle.com Protons are positively charged and electrons are negatively charged. Neutrons carry no charge or are electrically neutral. Thus the nucleus is always positively charged. The number of positive charges is always exactly balanced by an equal number of electrons, each of which carries one negative charge. The total number of protons in an atom is called its atomic number. Atoms are arranged in the periodic table in order of increasing atomic number. The total number of protons and neutrons in an atom is called the atomic mass number. Relative atomic mass is the mass of an atom of an element compared with the mass of an atom of carbon – 12. Masses of atoms are expressed as a ratio of their mass to that of carbon – 12 atom. The value of the relative atomic mass is dependent on the composition of each isotope of that element present in the sample. Isotopes are atoms of the same element with same atomic number but different atomic mass number. Relative molecular mass of molecules is calculated by adding the relative atomic mass of the atoms present in the molecule. The term formula weight is used to describe the mass of any chemical, whether it is an atom, molecule or ion. CHEMISTRY FOR YEAR 13 STRAND 2 Page 27 Calculation of Relative Atomic Mass To calculate the relative atomic mass of a sample of element, it is important to know which isotope of that element is present in the sample and in what proportions. The formula to calculate relative atomic mass is: ∑(𝐏𝐞𝐫𝐜𝐞𝐧𝐭𝐚𝐠𝐞 𝐀𝐛𝐮𝐧𝐝𝐚𝐧𝐜𝐞 × 𝐚𝐭𝐨𝐦𝐢𝐜 𝐦𝐚𝐬𝐬) Relative atomic mass = 𝟏𝟎𝟎 Example 1 Copper has two isotopes: 63 65 29Cu and 29Cu in the ratio of 69 % and 31 % respectively. Calculate the relative atomic mass of copper. Solution ∑(Percentage Abundance × atomic mass) Relative atomic mass = 100 [ (69 ×63)+(31 ×65)] = 100 = 63.62 Example 2 Thallium has two isotopes of atomic mass 203 amu and 205 amu. If the relative atomic mass of Thallium is 204.38, calculate the percentage abundance of each isotope. Solution Assume that 204.38 is = 100 % Let: x = % of 203Tl 100 - x = % of 205Tl ∑(Percentage Abundance × atomic mass) Relative atomic mass = 100 203x + 205 (100 − x) 204.38 = 100 20438 = 203x + 20500 – 205x 20438 = -2x + 20500 20438 − 20500 x= −2 x = 31 % Therefore: % of 203Tl = 31 % % of 205Tl = 100 – x = 100 – 31 % = 69 % CHEMISTRY FOR YEAR 13 STRAND 2 Page 28 Exercise 1. Lithium consists of 7.4 % 6Li and 92.6 % 7Li. Calculate the relative atomic mass of Lithium. 2. The element Rhenium consists of two isotopes: 185Re and 187Re in a ratio of 40 % and 60 %, respectively. Calculate the relative atomic mass of Rhenium. 3. Silicon has three isotopes: 28Si, 29Si and 30Si with a percentage abundance of 92.2 %, 4.7 % and 3.1 % respectively. Using the given information, calculate the relative atomic mass of Silicon. 4. Bromine has two isotopes, 79Br and 81Br. Both exist in equal amounts. Calculate the relative atomic mass of Bromine. 5. Rubidium has two naturally occurring isotopes, 85Rb and 87Rb. If Rubidium has a relative atomic mass of 85.47, calculate the percentage abundance of each isotope. 6. Two common isotopes of naturally occurring Neon are 20Ne and 22Ne. Calculate the percentage abundance of each isotope if the relative atomic mass of naturally occurring Neon is 20.18. Electron Configuration Electrons are arranged in energy levels or shells, and different energy levels can hold different number of electrons. Each energy level can also have sub- shells. The electron structure of an atom is a description of how the electrons are arranged, which can be shown in a diagram or by numbers. There is a link between the position of an element in the periodic table and its electronic structure. Elements in the same Group have the same number of valence electrons. Elements in the same Period have the same number of electron shells. The chemical properties of each element reflects its electron configuration. The electron configuration of an atom is the representation of the arrangement of electrons distributed among the shells and subshells. CHEMISTRY FOR YEAR 13 STRAND 2 Page 29 Orbitals are regions of space around the nucleus of an atom where an electron is likely to be found. Each orbital has the capacity to hold two electrons.  Electron Shell (Main Energy Level) – is a group of atomic orbitals with the same value of the principal quantum number. It can also be thought of as an orbit followed by electrons around an atoms nucleus. Electron shells have one or more electron subshells, or sub-levels.  Electron Subshells (Sub-Levels) – is a sub-division of an electron shell separated by electron orbitals. Subshells are labelled as s, p, d, and f in an electron configuration. A ‘s’ subshell has one orbital, a ‘p’ subshell has three orbitals, a ‘d’ subshell has five orbitals and a ‘f’ subshell has 7 orbitals. The table below shows how the shells split into subshells and the number of orbitals and electrons each subshell can hold. Shell Subshell Number of Maximum (Main Energy (Sub-Level) Orbital(s) Number of Level) Electrons Occupied 1 s 1 2 2 s 1 2 p 3 6 s 1 2 3 p 3 6 d 5 10 s 1 2 4 p 3 6 d 5 10 f 7 14 To correctly write the electron configuration of elements, it is very important to know the subshells which make up the shells and the number of electrons that can be placed in each subshell. Before attempting to write electron configuration using s, p, d notation, it is important to have knowledge on Quantum Numbers, Aufbau Principle, Pauli Exclusion Principle and Hund’s Rule. CHEMISTRY FOR YEAR 13 STRAND 2 Page 30 Quantum Numbers Quantum numbers describe the location and characteristics of an electron in an atom. A total of four quantum numbers are used to describe completely the movement and position of each electron in an atom. Each electron in an atom has a unique set of four quantum numbers according to the Pauli Exclusion Principle. Pauli Exclusion Principle states that no two electrons in an atom can have the same combination of four quantum numbers. Quantum numbers are important because they can be used to determine the electron configuration of an atom and the probable location of the atom's electrons. Quantum numbers are also used to determine other characteristics of atoms, such as ionisation energy and the atomic radius. The Four Quantum Numbers There are a total of four quantum numbers: the principal quantum number (n), the secondary quantum number (ℓ ), the magnetic quantum number (𝑚ℓ ) and the spin quantum number (𝑚𝑠 ). 1. The Principal Quantum Number (n) The principal quantum number, n, designates the electron shell or the main energy level. It describes the energy of an electron and the most probable distance of the electron from the nucleus. The larger the principal quantum number, the further the electron is from the nucleus therefore the larger the size of the orbital. All orbitals that have the same value of n are said to be in the same shell. The principal quantum number can be any positive integer starting at 1, as n = 1 designates the first principal shell or the innermost shell. The first principal shell is also called the ground state, or lowest energy state. The principal quantum number does not start from zero or a negative value since no electron exists with zero or negative amount of energy. 2. The Secondary Quantum Number (ℓ ) The secondary quantum number, ℓ, determines the shape of an orbital or the subshell in which electrons are placed. It divides the shells into smaller groups of subshells called orbitals. Each value of ℓ indicates a specific s, p, d subshell which are all unique in shape. The secondary quantum number is dependent on the principal quantum number. The secondary quantum number is also referred as the Azimuthal Quantum Number. The secondary quantum number starts from a value of zero to a positive integer one less than the principal quantum number (n – 1) ℓ = 0 to (n – 1) can range from 0 to (n – 1) CHEMISTRY FOR YEAR 13 STRAND 2 Page 31 Each subshell is denoted by a specific ℓ value. The table below shows the secondary quantum number (ℓ) value assigned to each of the subshell. Value of ℓ 0 1 2 Subshell s p d Example Determine all the possible values of the secondary quantum number (ℓ), if n = 3 and identify the type of subshells present. Note: the number of Solution subshells in a given shell equals the value ℓ = 0 to (n – 1) (all positive integers from 0 to n-1) of n for that shell (n = ℓ = 0 to (3 – 1) 3, so three subshells; ℓ = 0 to 2 s, p, d). ℓ = 0, 1, 2 (from 0 to 2) The third shell has three subshells: s (ℓ = 0), p (ℓ = 1), d (ℓ = 2) 3. The Magnetic Quantum Number (𝒎𝓵 ) The magnetic quantum number, 𝑚ℓ , specifies the orientation of orbital in space. It divides the subshell into individual orbitals which hold electrons. The value of mℓ depends on the secondary quantum number, ℓ, and the values range from -ℓ to +ℓ. 𝒎𝓵 = -ℓ to +ℓ Note: For a particular subshell, there are (2ℓ + 1) integral values of 𝑚ℓ. So if ℓ = 1, there are 2(1) + 1 = 3 𝑚ℓ values. These values are -1, 0, +1. So for the third shell (n = 3), the calculated ℓ values from the above example are 0, 1, 2. Using the above formula the 𝑚ℓ values can be assigned to the orbitals as follows: Principal Secondary Sub- Magnetic Quantum Number (𝒎𝓵 ) = -ℓ to +ℓ Quantum Quantum shell (ℓ) Number Number (ℓ ) (Number of mℓ values = 2ℓ + 1) (n) 2(0) + 1 = 1 0 s=0  1 value 𝑚ℓ = 0 mℓ = 0 2(1) + 1 = 3 3 1 p=1  3 values 𝑚ℓ = -1, 0, +1 mℓ = -1 0 +1 2(2) + 1 = 5 2 d=2  5 values 𝑚ℓ =-2,-1,0,+1,+2 mℓ =-2 -1 0 +1 +2 CHEMISTRY FOR YEAR 13 STRAND 2 Page 32 4. The Spin Quantum Number (𝒎𝒔 ) The spin quantum number, 𝑚𝑠 , determines the direction of the electron spin. The spin quantum number value of +½ indicates the electron is spinning upwards (spin up) and is represented by an upward half arrow (↿). The spin quantum number value of -½ indicates the electron is spinning downwards (spin down) and is represented by a downward half arrow (⇂). The spin quantum number indicates that electrons are spinning and as a result have potential to generate magnetic field. Appendix 3 shows all the possible quantum number values for the electrons in the 1st, 2nd and 3rd shell. Filling of Orbitals Each orbital has the capacity to hold a maximum of two electrons, each of which will have opposite spin. The s subshell has one orbital, so can accommodate two electrons and the p subshell has three orbitals, so it can hold six electrons. The d subshell has five orbitals, so can accommodate ten electrons and the f subshell has seven orbitals, so it can hold fourteen electrons (refer to the Table on page 30). While filling orbitals, it is very important to consider the energy level. Electrons are filled from the lowest energy level first. This is the Aufbau principle. The lower the principal quantum number, the lower the energy level. Figure 1 below shows the energy profile diagram for some orbitals and Figure 2 shows the order in which electrons are filled in orbitals. Figure 1 Figure 2 Energy Source: https://www.quora.com  According to Figure 1, 4s orbital is slightly lower in energy than the 3d orbitals.  Therefore, the order of filling orbitals (Figure 2) is 1s 2s 2p 3s 3p 4s 3d 4p …..  Total number of electrons in a shell = 2(n2), where n is the principal quantum number. CHEMISTRY FOR YEAR 13 STRAND 2 Page 33 IMPORTANT PRINCIPLES AND RULES  Aufbau Principle Aufbau Principle states that electrons always fill orbitals with lowest energy level first. This means that the orbitals closer to the nucleus are filled first.  Hund’s Rule Hund’s Rule states that when orbitals of equal energy are available, electrons are filled singularly first before any pairing can occur. This rule ensures that there are maximum unpaired electrons with minimum repulsions.  Pauli Exclusion Principle Pauli Exclusion Principle states that no two electrons in an atom can have the same combination of four quantum numbers. Although the first three quantum numbers (n, ℓ, 𝑚ℓ ) may be same, the spin quantum number (𝑚𝑠 ) would be different. Example 1 (a) Write the electron configuration for carbon. Solution Carbon has 6 electrons since atomic number is 6. Electron configuration: 1s2 2s2 2p2 (b) Draw the orbital diagram for carbon using the electron configuration. Solution CORRECT INCORRECT Note: When orbitals of equal energy are available, electrons are filled singularly first before any pairing occurs (Hund’s Rule). CHEMISTRY FOR YEAR 13 STRAND 2 Page 34 Example 2 Write the set of four quantum numbers for the 5th electron of carbon. Solution To find out the four quantum numbers, let us consider the orbital diagram of carbon. 5th electron  Principal Quantum Number (n) = 2 [5th electron is in the 2nd shell]  Secondary Quantum Number (ℓ ) = 0 to (n – 1) ℓ = 0 to (n – 1) = 0 to (2 – 1) = 0 to 1 = 0, 1 [s = 0, p = 1] ℓ = 1 [since 5th electron is present in p sub-shell]  Magnetic Quantum Number (𝒎𝓵 ) = -ℓ to +ℓ Number of 𝑚ℓ values = 2ℓ + 1 = 2(1) + 1 = 3 [from -ℓ to +ℓ] 𝑚ℓ = -1, 0, +1 [There are three values for mℓ] [The first p-orbital is assigned -1 0 +1 the lowest 𝑚ℓ value]. 5th electron Therefore, for the 5th electron, 𝑚ℓ = -1 𝟏  Spin Quantum Number (𝒎𝒔 ) = +𝟐 [Electron is spinning upwards] So for the 5th electron of carbon, the four quantum numbers are: n ℓ 𝒎𝓵 𝒎𝒔 (Appendix 3 can be referred 2 1 -1 𝟏 +𝟐 to verify the values.) CHEMISTRY FOR YEAR 13 STRAND 2 Page 35 Filling of 3d and 4s Orbitals Transition metals have d orbitals, and the 3d orbitals are filled after filling the 4s orbital since the empty 4s orbital is slightly lower in energy. Although 4s orbital is filled before 3d orbitals, the electron configuration is written showing all the orbitals in a given energy level together, in a sequence. Example Write the electron configuration of Iron and show how the electrons are filled in orbitals. Solution Iron (Atomic Number = 26) 1s2 2s2 2p6 3s2 3p6 3d6 4s2 and not 1s2 2s2 2p6 3s2 3p6 3d8 (CORRECT) (INCORRECT) 1s2 2s2 2p6 3s2 3p6 3d6 4s2 Exercise 1. Which of the following orbitals does not exist? A. 1s B. 2p C. 2d D. 3d 2. Which of the following can be an electron configuration of a Group II element? A. 1s2 2s2 2p6 3s2 3p2 C. 1s2 2s2 2p6 3s2 3p6 4s2 B. 1s2 2s2 2p6 3s2 3p6 4s1 D. 1s2 2s2 2p6 3s2 3p6 3d6 4s2 3. Consider elements Oxygen and Nickel: (a) Write their electron configuration using s, p, d notation. (b) Draw their orbital diagrams. (c) Determine the four quantum numbers for the: I. 8th electron of Oxygen II. 18th electron of Nickel CHEMISTRY FOR YEAR 13 STRAND 2 Page 36 Note:  Fully-filled orbitals (eg. 3d10) are more stable than half-filled orbitals (eg. 3d5).  Fully-filled orbitals (eg. 3d10) and half-filled orbitals (eg. 3d5) are more stable than partially filled orbitals (eg. 3d6). Electron configuration of Chromium is [1s2 2s2 2p6 3s2 3p6 3d5 4s1] and not [1s2 2s2 2p6 3s2 3p6 3d4 4s2]. The first configuration [1s2 2s2 2p6 3s2 3p6 3d5 4s1] has 3d and 4s orbitals half-filled whereas the second configuration [1s2 2s2 2p6 3s2 3p6 3d4 4s2] has 4s orbital fully filled and 3d orbital partially filled. Having 3d and 4s orbitals half-filled is more stable than having 4s orbital fully filled and 3d orbital partially filled. Electron configuration of Copper is [1s2 2s2 2p6 3s2 3p6 3d10 4s1] and not [1s2 2s2 2p6 3s2 3p6 3d9 4s2]. The first configuration [1s2 2s2 2p6 3s2 3p6 3d10 4s1] has 3d orbital fully- filled and 4s orbital half-filled whereas the second configuration [1s2 2s2 2p6 3s2 3p6 3d9 4s2] has 4s orbital fully-filled and 3d orbital partially filled. Extra stability is gained by having a fully filled and half- filled (3d10 4s1) combination. Electron Configuration of Ions Cations Ions which are positively charged are called cations. Cations form when a neutral atom loses electrons. Electrons are removed from the outermost orbitals first since these orbitals are furthest away from the nucleus. The electrons in these orbitals are weakly held by the nucleus. Thus, the 4s electrons are lost first before the 3d electrons as the 4s orbital is further away from the nucleus and becomes higher in energy once filled. The electrons in the 4s orbital are weakly held by the nucleus compared to 3d orbital electrons. Example Write the electron configuration of Zn2+. Solution Zn (atomic number = 30): 1s2 2s2 2p6 3s2 3p6 3d10 4s2 (30 electrons) Loss of 2 electrons from 4s subshell. Zn2+ (28 electrons): 1s2 2s2 2p6 3s2 3p6 3d10 CHEMISTRY FOR YEAR 13 STRAND 2 Page 37 Anions Ions which are negatively charged are called anions. Anions form when a neutral atom gains electrons. Electrons are added to outermost empty orbital or subshell. Example Write the electron configuration of S2-. Solution S (atomic number = 16): 1s2 2s2 2p6 3s2 3p4 (16 electrons) 2 electrons added to the 3p subshell. The resultant is a fully filled 3p subshell. S2- (18 electrons):1s2 2s2 2p6 3s2 3p6 1. Write the electron configuration for the following ions: (a) O2- (b) Cl- (c) Fe2+ (d) Cr3+ Abbreviated Electron Configuration (Noble Gas Notation) As chemists, we are primarily interested in the valence shell electrons. These electrons are involved in bonding and are responsible for the chemical properties of an atom. The inner shell electrons are called the core electrons and in an abbreviated configuration, it is represented by a noble gas. It is written in brackets and its configuration is same as the core electron configuration. This is followed by the configuration of the valence shell electrons for the particular element. The noble gas used in the abbreviated configuration is the noble gas that occurs at the end of the period preceding (before) the period containing the element. CHEMISTRY FOR YEAR 13 STRAND 2 Page 38 For transition metals, the d and the outer s subshell are written after the noble gas. The electrons below the outer s and d subshells of transition metals are unimportant. Example Write the abbreviated electron configuration for the following elements: (a) Sulphur (b) Iron Solution (a) Sulphur (atomic number = 16): 1s2 2s2 2p6 3s2 3p4 Ne (atomic number = 10) Sulphur - [Ne] 3s2 3p4 (b) Iron (atomic number = 26): 1s2 2s2 2p6 3s2 3p6 3d6 4s2 Ar (atomic number = 18) Iron - [Ar] 3d6 4s2 1. Write the abbreviated electron configuration for the following elements: (a) Si (b) Mn 2. Identify the element with the following abbreviated electron configuration: (a) [He] 2s2 2p1 (b) [Ne] 3s2 3p4 (c) [Ar] 3d8 4s2 CHEMISTRY FOR YEAR 13 STRAND 2 Page 39 Trends in the Periodic Table Periodic trend – is a pattern of change in the properties of elements in the periodic table with increasing atomic number. Important trends that would be studied include atomic radii, ionisation energy and electronegativity. 1. Atomic Radii Atomic radius describes the size of an atom. There is a variation in the atomic radii across the period and down the group of the periodic table. Across The Period Trend The atomic radii decreases across the period. Atoms become smaller moving from left to right of the periodic table. Reason Moving across the period, the effective nuclear charge increases as electrons are added to the same shell. As the nuclear charge increases, electrons are pulled closer to the nucleus, therefore, radius decreases. Down The Group Trend The atomic radius increases down the group. Atoms become larger moving down the group of the periodic table. Reason Moving down the group, the nuclear charge increases, however, more electron shells are added and valence electrons become further away from the nucleus creating a shielding effect. This masks the effect of increased nuclear charge. Due to the increased shielding effect the effective nuclear charge decreases, therefore, radius increases. Nuclear charge – is the total charge of all the protons in the nucleus. It has the same value as the atomic number. Effective nuclear charge – is the net positive charge experienced by valence electrons. When more electron shells are added, the effective nuclear charge decreases (even though nuclear charge increases) due to shielding effect. Shielding effect - describes the decrease in attraction between an electron and the nucleus in any atom with more than one electron shell. The greater the number of electron shells, the greater the shielding effect experienced by the valence electrons. CHEMISTRY FOR YEAR 13 STRAND 2 Page 40 2. Ionic Radii Ionic radius describes the size of an ion. Cations Cations (positively charged ions) are always smaller than the atoms from which they are formed. As electrons are removed from an atom, the electron- electron repulsion decreases. This allows the electrons to be pulled closer to the nucleus. Another reason for the decreased size of cations is that when electrons are removed, there are now fewer electrons to be pulled towards the nucleus by the same number of protons. This results in stronger attraction, therefore size decreases. Anions Anions (negatively charged ions) are always larger than the atoms from which they are formed. When electrons are added to an atom, there is an increase in the electron-electron repulsion. This causes the electrons to push apart and occupy a larger size. Adding electrons also results in weaker nuclear pull as more electrons are now available with the same number of protons. This results in weaker attraction, therefore size increases. Comparison of Ionic Radius with its Respective Neutral Atom Cations are smaller than Anions are larger than their respective atom. their respective atom. 1. Using the periodic table, arrange the elements from smallest to the largest in the set: C, F, Br, Ga 2. Explain why: (a) Ca2+ ion has a smaller radius than calcium atom. (b) O2- ion has a larger radius than oxygen atom. CHEMISTRY FOR YEAR 13 STRAND 2 Page 41 3. The radii of two atoms and an ion are given in the table below. Atom/ion Radius (picometres) Sodium atom 154 Chlorine atom 99 Chloride ion 181 Based on the above information, explain why the: (i) sodium atom is larger than the chlorine atom. (ii) chloride ion is larger than the chlorine atom. 3. Ionisation Energy Ionisation energy is the energy required to remove an electron from an atom in its gaseous state. It reflects how tightly an electron is held by the nucleus. X(g) X + (g) + e- Across The Period Trend The ionisation energy increases across the period. Reason Moving across the period, the effective nuclear charge increases, therefore the electrons are held strongly by the nucleus. The small size of the atom also results in the electron being held tightly. Down The Group Trend The ionisation energy decreases down the group. Reason Moving down the group, even though the nuclear charge increases, electrons are added to new energy levels. In addition, electrons are also more shielded, resulting in weak attraction. The large size of the atom also results in lower ionisation energy. The removal of the first electron from an atom is termed as the 1st ionisation energy. The successive removal of electrons from the ion can be termed as 2nd, 3rd, … ionisation energy. CHEMISTRY FOR YEAR 13 STRAND 2 Page 42 IMPORTANT NOTE  To remove electrons from shells, subshells or orbitals which are fully or half-filled require more energy as they are very stable.  Comparison of the successive ionisation energy values will show a large increase if electrons are removed from higher energy levels. 1. (a) In which

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