GR 10 Physical Science Chemistry Paper 2 PDF

GR 10 PHYSICAL SCIENCE CHEMISTRY PAPER 2 Compiled by: Miss C. Wilson (HS Birchleigh) and Miss S. Wilson (Edenglen HS) Index Page HOW TO USE THIS DOCUMENT 6 Table of Physical Qua...

GR 10 PHYSICAL SCIENCE CHEMISTRY PAPER 2 Compiled by: Miss C. Wilson (HS Birchleigh) and Miss S. Wilson (Edenglen HS) Index Page HOW TO USE THIS DOCUMENT 6 Table of Physical Quantities 7 Data sheet for Paper 2 7 MODULE 1: MATTER AND MATERIALS Unit 1: Matter and classification 1.1. Exam guidelines 9 1.2. Notes 11 1.2.1. Properties of matter 1.2.2. Mixtures 1.2.3. Element 1.2.4. Compound 1.2.5. Pure substance 1.2.6. Writing of chemical formulae 1.2.7. Metals, non-metals and metalloids 1.2.8. Electrical conductors, semiconductors and insulators 1.2.9. Thermal conductors and insulators 1.2.10. Magnetic and non-magnetic materials 1.3. Exercise 1 Unit 2: States of matter and the Kinetic Molecular Theory. 2.1. Exam guidelines 2.2. Notes 2.2.1. States of matter 2.2.2. Phase change 2.2.3. Heating and cooling curve 2.2.3.1. Explanation of heating curve 2.2.4. Kinetic Molecular Theory 2.3. Exercise 2 i Unit 3: Atomic structure. 3.1. Exam guidelines 3.2. Notes 3.2.1. Atomic number (Z) 3.2.2. Atomic mass (A) 3.2.3. Atom 3.2.4. Ions 3.2.5. Isotopes 3.2.5.1. Relative atomic mass 3.2.6. Electron configuration 3.2.6.1. Orbitals 3.2.6.2. Aufbau 3.2.6.3. sp-notation 3.2.6.4. Abbreviated electron configuration 3.3. Exercise 3 Unit 4: The Periodic Table. 4.1. Exam guidelines 4.2. Notes 4.2.1. Periodic Table as on gr 10 -12 data sheets 4.2.2. Electron configuration and the periodic table 4.2.3. Periodicity 4.2.3.1. Atomic radius 4.2.3.2. Ionisation energy 4.2.3.3. Electron affinity 4.2.3.4. Electron negativity 4.3. Exercise 4 ii Unit 5: Chemical Bonds. 5.1. Exam guidelines 5.2. Notes 5.2.1. Chemical bonding 5.2.2. Lewis dot diagram 5.2.3. Covalent bond 5.2.3.1. Polar and Non-polar 5.2.3.2. Properties of covalent bonds 5.2.4. Ionic bond 5.2.4.1. Crystal lattice 5.2.4.2. Properties of ionic compounds 5.2.5. Metallic bond 5.2.5.1. Properties of metals 5.2.6. Calculate relative molecular masses for covalent molecules 5.2.7. Calculate relative formula masses for ionic compounds 5.3. Exercise 5 MODULE 2: CHEMICAL CHANGE Unit 1: Physical and Chemical change. 1.1. Exam guidelines 1.2. Notes 1.2.1. Separation of particles in physical and chemical change 1.2.1.1. Physical change 1.2.1.2. Chemical change 1.2.2. Conservation of atoms and mass 1.2.3. Representing chemical change 1.3. Exercise 6 iii Unit 2: Reactions in Aqueous solutions. 2.1. Exam guidelines 2.2. Notes 2.2.1. Ions in aqueous solutions: their interaction and effects 2.2.1.1. Ionic solutions 2.2.1.2. Covalent solutions 2.2.2. Electrolytes and extent of ionization as measured by conductivity 2.2.3. Precipitation reactions 2.2.4. Other chemical reaction types in water solution 2.3. Exercise 7 Unit 3: Quantitative Aspects of Chemical change 3.1. Exam guidelines 3.2. Notes 3.2.1. Atomic mass and the mole concept 3.2.2. Molecular and formula masses 3.2.2.1. Relationship between molar mass, mass, and mole 3.2.2.2. Empirical formula 3.2.2.3. Water crystallisation 3.2.3. Determining the composition of substances 3.2.3.1. Determine percentage composition of substances 3.2.3.2. Concentration 3.2.4. Molar volume of gases 3.2.5. Basic stoichiometric calculations 3.2.5.1. Mass calculations 3.2.5.2. Volume calculations 3.2.5.3. Theoretical and Actual yield 3.3. Exercise 8 iv  Organised to fit with the Exam Guidelines for Physical Sciences. HOW TO USE THIS DOCUMENT 1. This document was compiled as an extra resource to help you perform in Physical Science. 2. Firstly, you must make sure that you study the TERMS and DEFINITIONS provided for each Topic. THEORY always forms part of any test or examination, and you should ensure that you obtain FULL MARKS for ALL THEORY questions. REVISE terms and definitions of topics already completed frequently so that you know by the time you are sitting for a test or an examination. 3. Read through the Exam Guidelines for each topic so that you know what can be expected from you in tests and exams. 4. Use these NOTES to study the THEORY for each topic. 5. Your teacher will be using some of the questions to round of the work for a specific topic. The answers for each topic will be supplied after the topic has been completed. 6. Use the questions on a certain topic to prepare for test and examinations. DO NOT look at the answers before attempting the questions. First try it yourself. Compare your answers with the answers. Mark your work with a pencil and do corrections for your incorrect answers. If you do not know how to answer a question, the answers are there to guide you. Use questions that you answered incorrectly to ask your teacher for help. Acquaint yourself with the way in which a particular type of question should be answered. Answers supplied are from memoranda used to mark the questions in previous years. 1  Table with Physical Quantities Preferred Alternative Unit Quantity Unit name symbol symbol symbol Mass m Gram g Mole n Mole mol Molecular mass Mr Grams per mole g∙mol-1 Mole per cubic Concentration c mol∙dm-3 decimetre Volume V Decimetre cube dm-3 DATA FOR PHYSICAL SCIENCES GR 10 PAPER 2 (CHEMISTRY) TABLE 1: PHYSICAL CONSTANTS NAME SYMBOL VALUE Standard pressure p 1,013 x 105 Pa Molar gas volume at STP Vm 22,4 dm3∙mol-1 Standard temperature T 273 K Charge on electron e -1,6 × 10-19 C Avogadro’s constant NA 6,02 × 1023 mol-1 TABLE 2: FORMULAE m N n= n= M NA n m V c= or c = n= V MV Vm 2 z 3  MODULE 1 MATTER AND MATERIAL UNIT 1: MATTER AND CLASSIFICATION 1.1. EXAM GUIDELINES The material(s) of which an object is composed. Describe matter as being made up of particles whose properties determine the observable characteristics of matter and its reactivity. Define properties of materials: o Strength o Brittle: Hard but likely to break easy. Malleable: Ability to be hammered or pressed into shape without breaking or cracking. Ductile: Ability to be stretched into a wire. o Density: The mass per unit volume of a substance. o Melting points and boiling points Boiling point: The temperature of a liquid at which its vapour pressure equals the external (atmospheric) pressure. Melting point: The temperature at which a solid, given sufficient heat, becomes a liquid. Mixtures: heterogeneous and homogeneous Define a homogeneous mixture as a mixture of uniform composition and in which all components are in the same phase, e.g., a solution of salt and water. Define a heterogeneous mixture as a mixture of non-uniform composition and of which the components can be easily identified, e.g., sand and water. Give examples of heterogeneous and homogeneous mixtures. Classify given mixtures as homogenous and heterogeneous. Pure substances: elements and compounds Use symbols to represent elements and compounds. Define an element as a pure substance consisting of one type of atom. Define a compound as a pure substance consisting of two or more different elements. Define a pure substance as a substance that cannot be separated into simpler components by physical methods. 4  Classify given substances as pure or impure and as compounds or elements. Names and formulae of substances Write names of compounds from given formulae or write down formulae of compounds from given names. Write names of ions from given formulae or formulae from given names. Write names of substances or ions ending on -ide, -ite and –ate. Write names of substances using the prefixes di-, tri-, etc. Metals, metalloids and non-metals Classify substances as metals, metalloids and non-metals using their properties. Identify the metals, their positions on the periodic table and their numbers in comparison with the number of non-metals. Identify the non-metals and their positions on the periodic table. Describe metalloids as having properties of metals and non-metals. Describe the characteristic property of metalloids that show increasing conductivity with increasing temperature (the reverse of metals), e.g., silicon and graphite. Identify the metalloids and their position on the periodic table. Electrical conductors, semiconductors, and insulators Define the terms electrical conductor, semiconductor, and electrical insulator: o Electrical conductor: A material that allows the flow of charge. o Semiconductor: A substance that can conduct electricity under some conditions, but not others, making it a good medium for the control of electrical current. o Electrical insulator: A material that prevents the flow of charge. Classify materials as electrical conductors, semiconductors, and insulators. Give examples of electrical conductors, semiconductors, and insulators. Thermal conductors and insulators Define the terms thermal conductor and thermal insulator. A thermal conductor is a material that allows heat to pass through easily, whilst a thermal insulator does not allow heat to pass through it. Describe a test to classify materials as thermal conductors and insulators. Give examples of materials that are thermal conductors and insulators. Magnetic and nonmagnetic materials Describe how to test and classify materials as magnetic and non- magnetic. 5  Give examples of materials that are magnetic and non-magnetic. Give examples of how we use magnets in daily life (in speakers, telephones, electric motors and as compasses). 1.2. NOTES Chemistry is the science of matter and the changes it undergoes. Matter is made up of particles whose properties determine the observable characteristics of matter and its reactivity. Matter can be seen as anything that occupies space and possesses mass. There is many different types of matter, it can be hard but brittle, soft but strong, etc. Air, water, soil, people, and Earth itself are all made up of matter. 1.2.1. Properties of matter. Strength Strong - when it can support a heavy load without breaking, tearing, or changing form. E.g., cement and steel. Weak - break or bend easily. E.g., paper and fabric. Brittle Hard but likely to break easily. How easily it breaks (shatters) e.g., pottery. Malleable Ability to be hammered or pressed into shape without breaking or cracking. e.g., Aluminium foil used mostly in the food industry. 6 Ductile Ability to be stretched into a wire. e.g., Copper to make electric wire. Density The mass per unit volume of a substance. An example of a dense material is concrete, and a less dense material is polystyrene. Lead is much denser than Aluminium. Melting point The temperature at which a solid, given sufficient heat, becomes a liquid. The melting point of water (ice) is 0ºC, that’s why water a liquid at room temperature (25ºC), 25ºC > 0ºC (already melted). Boiling point The temperature of a liquid at which its vapour pressure equals the external (atmospheric) pressure. The boiling point of water is 100ºC, that’s why water is a liquid at room temperature (25ºC), 25ºC < 100ºC (not boiling yet). Melting and boiling points of materials is specific to each material, which helps in the classification of materials as gases, liquids, or solids at specific temperatures. E.g., Iron melts at 1538 ºC and boils at 2750 ºC, which makes it a solid at room temperature (25 ºC). 7 1.2.2. Mixtures A mixture is a combination of two or more substances retain their own properties. A mixture’s composition can vary. Mixtures can be separated by physical means. E.g., air, fizzy drinks, alloys. Homogeneous mixtures. HOMOGENEOUS MIXTURE: Is a mixture of uniform composition and which all components are in the same phase. The different components cannot be distinguished from each other. The components are all in the same phase (state). E.g. Fruit juice (different fruit juices), Air (Nitrogen, Oxygen, carbon dioxide, water vapour), Cool drink concentrates in water (consist of two liquids that are mixed), Alloys (brass – Cu and Zn, Steel – Fe and C, Bronze – Cu and Sn, Stainless steel – Fe, Ni and Cr.) Methods to separate homogeneous mixtures o Evaporation, liquid evaporate and solid stays behind e.g. sugar and water. o Distillation, evaporate the liquid and then let the liquid condense e.g. water from a solution o Fractional distillation, the same as distillation but at specific temperatures so it is possible to separate two liquids with different boiling points, e.g. alcohol and water o Chromatography, paper or gas chromatography, smallest particles move through the furthest and fastest where biggest particles shorter distance and the slowest e.g. mixture of gasses, liquids or dissolved substances, inks are mixed to get a specific colour. 8 Heterogeneous mixtures. HETEROGENEOUS MIXTURE: Is a mixture of non-uniform composition and of which the components can easily be identified. Particles are not in the same phase (state) Can see the different particles clearly. E.g. Mud (a mixture of different soil particles and water). Stew (the meat, vegetables and gravy is clearly distinguishable) Salad dressing (a mixture of vinegar, oil and herbs) Blood (red blood cells, white blood cells, platelets, plasma, etc.) Methods to separate heterogeneous mixtures. o Filtration, use filter paper with specific density that allows specific size particles through only, e.g. sand and water. o Separating funnel, a funnel with a small tap where the substances have different densities when you open the tap the bottom part can be funnelled off and tap closed again to stop top substances e.g. oil and water. o Decantation, throwing off the liquid at the top and leaving the solid parts behind e.g., sand and water. o Centrifuge, spinning fast different size particles forms layers e.g. blood, spinning water out of washing. 1.2.3. Element An element is a pure substance consisting of one type of atom. It cannot be broken down into simpler substances, with either chemical or physical means. The smallest unit of an element is the atom or a diatomic element or a molecule consisting out of the same element. e.g., H, H2, He, Li, Cl2, O2, O3, S8. 9 1.2.4. Compound A compound is a pure substance consisting of two or more different elements. When two or more different elements combine in a fixed ratio. The particles can be broken down by chemical means. The smallest particle is a molecule in a covalent bond and an ion in an ionic bond. e.g., H2O, FeS, NaCl. Each compound has its own formula that tells us: The type of elements in the compound. The number of atoms of each element in the compound. SiH4 (Siliane) 1 Si-atom and 4 H-atoms. Cl2 (Chlorine) 2 Cl-atoms. HCl (Hydrogen chloride) 1 H-atom and 1 Cl-atom. 10 NO2 (Nitrogen dioxide) 1 N-atom and 2 O-atoms. 1.2.5. Pure Substance A pure substance is a substance that cannot be separated into simpler components by physical methods. Pure substance can be elements or compounds. They are substances with one type of constituent particle (same ingredient). ∴ not a mixture. Distilled water consist only out of H2O which makes it a pure substance. Whereas still mineral water (Valprѐ, Bon aqua, Nestlé, etc) have minerals like calcium, magnesium, etc which makes it not pure. Name and formulae of substances. Each element on the periodic table has a specific name. When elements combine compounds are formed. The compound’s chemical name will always include the names of the elements that combined to form it. o A compound of hydrogen (H) and chlorine (Cl) is hydrogen chloride (HCl). o A compound of hydrogen (H) and oxygen (O) is hydrogen oxide (H2O), commonly known as water. In a compound the element on the left of the periodic table is usually used first in naming. 11  e.g., KCl, potassium (K) is in group 1 (left) and chlorine (Cl) in group 7 (right). Potassium therefore comes first in the compound name, potassium chloride. A compound may contain ions (an atom that have either lost an electron, positive ion, or gained an electron, negative ion). The ions can be single (containing one element) or compound (containing two or more different elements). Table 1: Cations (positive ions) Hydrogen ion 𝐻+ Lead (II) ion 𝑃𝑏2+ Lithium ion 𝐿𝑖 + Chromium (II) ion 𝐶𝑟 2+ Sodium ion 𝑁𝑎+ Manganese (II) ion 𝑀𝑛2+ Potassium ion 𝐾+ Iron (II) ion 𝐹𝑒 2+ Silver ion 𝐴𝑔+ Cobalt (II) ion 𝐶𝑜 2+ Mercury(I) ion 𝐻𝑔+ Nickel (II) ion 𝑁𝑖 2+ Copper(I) ion 𝐶𝑢+ Copper (II) ion 𝐶𝑢2+ Ammonium ion 𝑁𝐻4 + Zinc ion 𝑍𝑛2+ Oxonium ion H3O+ Aluminium ion 𝐴𝑙 3+ Beryllium ion 𝐵𝑒 2+ Chromium (III) ion 𝐶𝑟 3+ Magnesium ion 𝑀𝑔2+ Iron (III) ion 𝐹𝑒 3+ Calcium ion 𝐶𝑎2+ Cobalt (III) ion 𝐶𝑜 3+ Barium ion 𝐵𝑎2+ Chromium (VI) ion 𝐶𝑟 6+ Tin(II) ion 𝑆𝑛2+ Manganese (VII) ion 𝑀𝑛7+ 12 Table 2: Anions (negative ions) Fluoride ion 𝑭− Oxide ion 𝑶𝟐− Chloride ion 𝑪𝒍− Peroxide ion 𝑶𝟐 𝟐− Bromide ion 𝑩𝒓− Carbonate ion 𝑪𝑶𝟑 𝟐− Iodide ion 𝑰− Sulphide ion 𝑺𝟐− Hydroxide ion 𝑶𝑯− Sulphite ion 𝑺𝑶𝟑 𝟐− Nitrite ion 𝑵𝑶𝟐 − Sulphate ion 𝑺𝑶𝟒 𝟐− Nitrate ion 𝑵𝑶𝟑 − Thiosulphate ion 𝑺𝟐 𝑶𝟑 𝟐− Hydrogen carbonate ion 𝑯𝑪𝑶𝟑 − Chromate ion 𝑪𝒓𝑶𝟒 𝟐− Hydrogen sulphite ion 𝑯𝑺𝑶𝟑 − Dichromate ion 𝑪𝒓𝟐 𝑶𝟕 𝟐− Hydrogen sulphate ion 𝑯𝑺𝑶𝟒 − Manganate ion 𝑴𝒏𝑶𝟒 𝟐− Dihydrogen phosphate ion 𝑯𝟐 𝑷𝑶𝟒 − Oxalate ion 𝑪𝟐 𝑶𝟒 𝟐− Hydrogen Hypochlorite ion 𝑪𝒍𝑶− 𝑯𝑷𝑶𝟒 𝟐− phosphate ion Chlorate ion 𝑪𝒍𝑶𝟑 − Nitride ion 𝑵𝟑− Permangate ion 𝑴𝒏𝑶𝟒 − Phosphate ion 𝑷𝑶𝟒 𝟑− Acetate/ethanoate ion 𝑪𝑯𝟑 𝑪𝑶𝑶− Phosphide ion 𝑷𝟑− Pay special attention to the ending of the names -ide (no O-atom), - ite (contain at least 1 O-atom) and -ate (containing the most O-atoms of similar compound ions). 13  Prefixes, mono- (one), di- (two) and tri- (three) are also used for the same element. e.g., SO2 – sulphur dioxide – 2 O-atoms SO3 – sulphur trioxide – 3 O-atoms. Table 3: Formulae, Chemical name and Common (household) name. Formula Chemical name Common name H 2O Hydrogen oxide Water HCl Hydrogen chloride Hydrochloric acid HNO3 Hydrogen nitrate Nitric acid HNO2 Hydrogen nitrite Nitrous acid H2SO4 Hydrogen sulphate Sulphuric acid H2SO3 Hydrogen sulphite Sulphurous acid H2CO3 Hydrogen carbonate Carbonic acid (COOH)2 Hydrogen oxalate Oxalic acid NaOH Sodium hydroxide Caustic soda NH3 Hydrogen nitride Ammonia NaCl Sodium chloride Table salt CaO Calcium oxide Quicklime Ca(OH)2 Calcium hydroxide Slaked lime Mg(OH)2 Magnesium hydroxide Milk of magnesia Fe3O4 Magnetic iron oxide Magnetite K2CO3 Potassium carbonate Potash CaCO3 Calcium carbonate Marble, limestone, chalk Na2CO3 Sodium carbonate Washing soda, soda ash Sodium hydrogen Baking soda, bicarbonate NaHCO3 carbonate (Sodium of soda bicarbonate) KNO3 Potassium nitrate Saltpetre NaNO3 Sodium nitrate Chile saltpetre MgSO4 Magnesium sulphate Epsom salt CaSO4 Calcium sulphate Gypsum, plaster CuSO4 Copper sulphate Blue vitriol H 2S Hydrogen sulphide Sulphureted hydrogen 14 1.2.6. Writing of chemical formulae. 1. Sodium oxide 𝑁𝑎 𝑂 𝑁𝑎+ 𝑂2− d Two sodium’s needed to balance the -2 charge of oxygen. Na2O 2. Iron(III) chloride 𝐹𝑒 𝐶𝑙 𝐹𝑒 3+ 𝐶𝑙 − Three chlorides needed to balance the +3 charge of iron. FeCl3 15 1.2.7. Metals, Non-metals, and Metalloids. The periodic table can be divided into 3 major groups. On the left side the metals except for hydrogen (H) and the right side the non-metals, the division between the two, forms a basic zig-zag line which makes out the metalloids. NON-METALS H B METALS Si Ge As Sb Te Po At METALLOIDS Figure 1: Diagram showing part of the periodic table. Metals Solids, except for mercury (Hg). High melting and boiling points. Usually, shiny. Malleable and ductile. Good conductors of electricity (ability decrease as it becomes hot). Good conductor of heat. Usually strong and can hold large weights. Grey (silver) in colour, except for gold and copper. 16 Non-metals Solids or gases at room temperature, except bromine (Br) which is a liquid. Low melting and boiling points. Are softer than metals except diamonds which is very hard and have a very high melting point (consist out of C-atoms). Density is often low. Poor conductors of electricity, except graphite (consists out of C- atoms). Tend to be insulators. Poor thermal (heat) conductors. Brittle, with a dull surface when it’s a solid. Varies in colour (have different colours). Metalloids Have metal as well as non-metal properties. They often look like metals, but they are brittle like non-metals. They are neither conductors nor insulators but make excellent semiconductors. The more it is heated the better its electrical conductivity becomes. Boron (B); Silicon (Si); Germanium (Ge); Arsenic (As); Antimony (Sb); Tellurium (Te); Polonium (Po); and Astatine (At). 17 1.2.8. Electrical conductors, semiconductors, and insulators. Electrical conductors. A material that allows the flow of charge. Conductors let electricity pass through them easily. Metals are generally good conductors of electric current. Metal’s ability to conduct electricity decrease with an increase in temperature. Semiconductor. A substance that can conduct electricity under some conditions, but not others, making it a good medium for the control of electrical current. Metalloids are weak conductors and are called semiconductors. Metalloid’s ability to conduct electricity increase with the increase in temperature. Electrical insulator. A material that prevents the flow of charge. Non-metals are poor conductors of electricity and are called insulators (except graphite). e.g., Electrical conductor Semiconductor Insulator Gold Silicon Sulphur Graphite Antimony Oxygen 18 1.2.9. Thermal conductors and insulators. Thermal conductor. Is a material that allows heat to pass through easily. Metals are generally good conductors of heat. The way in which metals are arranged is the reason why they are good conductors. Thermal insulator. Does not allow heat to pass through it. Non-metals are poor conductors of heat. A test to classify materials as thermal conductors and insulators. o Place a piece of substance in a beaker that is filled with boiled water and note how fast it heats up. Be careful not to burn. o Experiment 2 pg. 30 Docscientia Chemistry Gr 10. ▪ Drip a little candle wax on one end of each material being investigated. ▪ Heat the other end of each material over a flame or dip it into boiling water. ▪ Observe how quickly the candle wax melts, and which candle wax melts first. 19 1.2.10. Magnetic and Non-magnetic materials. Magnetism is a force that magnetic objects, can exert on each other without physically touching. A magnetic object is surrounded by a magnetic field that gets weaker as one moves further away from the object. Earth behaves like a giant magnet. Earth’s core consists of molten iron (Fe) or nickel (Ni). The elements exhibit strong magnetic properties, namely Fe, Ni, and Co (cobalt), which are called ferromagnetic materials. Some alloys (mixtures of metals) are also magnetic. o Ferromagnetic materials are made up of magnetically aligned regions called domains. o Each domain behaves like a tiny magnet, with a north and south pole. To test if a substance is magnetic use a magnet. Magnets are used in our everyday life’s, e.g., compasses, motors, generators, speakers etc. 20 1.3. EXERCISE 1 1. Copy and complete the following table. Classify the following as a mixture or a pure substance, for mixtures indicate the type of mixture: Homogeneous or Mixture or Pure Substance Heterogeneous substance mixture 1.1. Fizzy cold drink 1.2. Steel Iron fillings and 1.3. sulphur powder 1.4. Smoke Limestone 1.5. (CaCO3) 1.6. Blood 1.7. Bottled water 1.8. Distilled water 1.9. Table salt 1.10. Air 1.11. Milk 1.12. Muesli 1.13. Bronze 1.14. Cup of coffee Sugar 1.15. (C12H22O11) 2. Copy and complete the following table by giving the formulae of the compounds formed: OH- NO3- SO42- CO32- MnO4- PO43- Na+ Ca2+ K+ Mg2+ Al3+ NH4+ 21 3. Copy and complete the following: Common name Formula Chemical name (Household name) HCl Hydrogen chloride Hydrochloric acid 3.1. 3.2. Sulphureted hydrogen H 2O 3.3. 3.4. 3.5. Copper sulphate 3.6. CaSO4 Calcium sulphate 3.7. 3.8. Magnesium sulphate 3.9. NaNO3 3.10. 3.11. 3.12. 3.13. Saltpetre NaHCO3 3.14. 3.15. 3.16. Potassium carbonate Potash 3.17. Sodium hydroxide 3.18. NH3 3.19. 3.20. 3.21. Hydrogen sulphate 3.22. 3.23. 3.24. Nitrous acid 4. Copy and complete the following table: Metal / Thermal Electrical Magnetic / Substance Non- conductor conductor Non- metal / insulator / insulator magnetic Iron Copper Graphite (C) Nickel Silicon Sulphur Hydrogen Chlorine Argon 5. Grade 10 learners were given the substances in the table below: Bronze Salad dressing Mercury Simple syrup Epsom salt Silver oxide Carbon dioxide Air Salsa Chlorine gas Iron pipe Mud From the table above write down: 22 5.1. An element. 5.2. Homogeneous mixture. 5.3. A solid compound at room temperature. 5.4. A heterogeneous mixture. 5.5. A solid mixture. 5.6. A diatomic gas. 5.7. A liquid metal. 6. Define the following terms: 6.1. Strength 6.2. Brittle 6.3. Malleable 6.4. Ductile 6.5. Density 6.6. Boiling point 6.7. Melting point 6.8. Element 6.9. Compound 6.10. Pure substance 6.11. Electrical conductor 6.12. Semiconductor 6.13. Electrical insulator 6.14. Thermal conductor 6.15. Thermal insulator 7. Choose from the following list the description that fits with each of the pictures below and write it down next to the pictures number. Element Compound Mixture of elements Mixture of compounds Mixtures of elements and compounds Each circle represents an atom. Each different colour represents a different kind of atom. When two atoms touch they are bonded together. 23 7.1. 7.7. 7.2. 7.8. 7.3. 7.9. 7.4. 7.10. 7.5. 7.11. 7.6. 7.12. 24  UNIT 2: STATES OF MATTER AND THE KINETIC MOLECULAR THEORY 2.1. EXAM GUIDELINES Three states of matter Describe the particle nature of matter by referring to diffusion and Brownian motion. Diffusion: The movement of atoms or molecules from an area of higher concentration to an area of lower concentration. Brownian motion: The random movement of microscopic particles suspended in a liquid or gas, caused by collisions between these particles and the molecules of the liquid or gas. List and characterise the three states of matter. Define freezing point, melting point, and boiling point. Boiling point: The temperature of a liquid at which its vapour pressure equals the external (atmospheric) pressure. Melting point: The temperature at which a solid, given sufficient heat, becomes a liquid. Freezing point: The temperature at which a liquid change to a solid by the removal of heat. Interpret/Draw heating and cooling curves and interpret data given on heating and cooling curves. Identify the physical state of a substance at a specific temperature, given the melting point and the boiling point of the substance. Define melting, evaporation, freezing, sublimation, and condensation as changes in state. Melting: The process during which a solid change to a liquid by the application of heat. Evaporation: The change of a liquid into a vapour at any temperature below the boiling point. (Note: Evaporation takes place at the surface of a liquid, where molecules with the highest kinetic energy are able to escape. When this happens, the average kinetic energy of the liquid is lowered, and its temperature decreases. Freezing: The process during which a liquid change to a solid by the removal of heat. Sublimation: The process during which a solid change directly into a gas without passing through an intermediate liquid phase. 25  Condensation: The process during which a gas or vapour changes to a liquid, either by cooling or by being subjected to increased pressure. Kinetic Molecular Theory Describe a solid, a liquid, and a gas according to the Kinetic Molecular Theory in terms of particles of matter. According to the Kinetic Molecular Theory: o Matter consists of small particles. o The particles are in constant motion. o There are forces of attraction between the particles. o Particles collide (with the sides of the container and each other) and exert pressure. o The temperature of a substance is a measure of the average kinetic energy of the particles. o A phase change may occur when the energy of particles changes. 2.2. NOTES The physical state (phase) of substances is one way to classify matter. The Kinetic Molecular Theory with Intermolecular forces forms the basis for solid (s), liquid (ℓ), and gas (g). Robert Brown studied pollen grains that was suspended in liquid. He noted that the pollen grains had jerky random motion, this later was called the Brownian motion. Microscopic view of movement of a small particle suspended in a liquid. Diffusion: The movement of atoms or molecules from an area of higher concentration to an area of lower concentration. 26 Brownian motion: The random movement of microscopic particles suspended in a liquid or gas, caused by collisions between these particles and the molecules of the liquid or gas. 2.2.1. States of matter Table 1: States of matter (phases). Solid (s) Liquid (ℓ) Gas (g) Packed close together Arranged irregularly in a regular Packed together in an and spread very far arrangement or irregular arrangement. apart. lattice. Does not move freely, Move about freely in a Have enough kinetic but vibrate about their confined space, some energy to enable them fixed positions, almost kinetic energy. to move random. no kinetic energy. Fixed volume but take The matter expands Fixed volume and on the shape of the to occupy whatever shape. container. volume is available. Have attractive Have almost no Have strong attractive (intermolecular) forces attractive (intermolecular) forces between the particles, (intermolecular) forces between the particles. but the forces are between the particles. weaker than in solids. Liquids exert pressure Gas exerts pressure Cannot be in all directions and in all directions and compressed. can be slightly can be compressed. compressed. Specific boiling point Specific melting point Specific condensation and freezing point under standard points under standard under standard conditions. conditions. conditions. 27 2.2.2. Phase change During a phase change: The physical properties as well as the potential energy of the substance change. The chemical composition (formula) stays the same. Substances undergoes phase change because of specific boiling points; melting points; and freezing points of a specific substance. Boiling point: The temperature of a liquid at which its vapour pressure equals the external (atmospheric) pressure. Melting point: The temperature at which a solid given sufficient heat, becomes a liquid. Freezing point: The temperature at which a liquid change to a solid by the removal of heat. Table 2: Melting and Boiling points. Substance Melting point (ºC) Boiling point (ºC) Metals Copper 1 083 2 600 Mercury -39 357 Sodium 98 892 Magnesium 650 1 107 Lead 327 1 750 Iron 1 536 3 000 Metalloids Silicon 1 420 3 280 Germanium 940 2 830 Salts Table salt (NaCl) 801 1 413 Rust (Fe2O3) 1565 decomposes 28 Melting: The process during which a solid change to a liquid by the application of heat. Evaporation: The change of a liquid into a vapour at any temperature below the boiling point. Evaporation takes place at the surface of a liquid, where molecules with the highest kinetic energy are able to escape. When this happens the average kinetic energy of the liquid is lowered, and its temperature decreases. Table 3: Difference between boiling and evaporation. Boiling Evaporation Occurs at a specific temperature Takes place at any temperature (boiling point). below boiling point. Takes place at the surface of the Takes place throughout the liquid. liquid only. Takes place when there are more Takes place when the pressure of the high energy particles, which can vapour (gas) in the bubbles is equal overcome the attraction forces and to the pressure of the atmosphere. escape from the liquid. Sublimation is the process where a solid changes directly to a gas without going through an intermediate liquid state, by adding heat. 29 Deposition is the process where a gas changes directly to a solid without going through an intermediate liquid state, by the removal of heat. Condensation is a process where a gas changes to a liquid by cooling or increasing the pressure. Freezing: The process during which a liquid change to a solid by the removal of heat. When the melting and boiling point of a substance is known the state (s; ℓ; or g) of the substance can be determined at any temperature. e.g., the melting point for HF is -83ºC and the boiling point is 20ºC, HF will be a solid at -90ºC and a gas at 25ºC. 2.2.3. Heating and Cooling curve A heating curve of a substance gives the changes in temperature as it moves from a solid to a liquid to a gas. A cooling curve gives the changes in temperature as it moves from a gas to a liquid to a solid. An important observation is that as a substance melts or boils, the temperature remains constant until the substance has change state completely. This is because all the heat energy goes into breaking or forming the bonds between the molecules. Phase change: o Temperature and kinetic energy remain the same. o Potential energy changes. 30 2.2.3.1. Explanation of heating curve AB Solid absorb energy without melting. The temperature increases. The particles will vibrate more. The kinetic energy increases. BC Substance begins to melt. The temperature remains the same till all the solid particles has melted. Movement of the particles remains the same. The potential energy increases. CD The liquid absorb energy without boiling. The temperature increases. The particles move faster. The average kinetic energy increases. DE The substance boils. The temperature remains the same while its boiling until all the particles becomes a gas. The movement of the particles remains the same. The potential energy increases. EF The vapour particles absorb energy. The temperature increases. The particles move faster. The average kinetic energy increases. 2.2.4. The Kinetic Molecular Theory The kinetic theory of matter attempts to explain the behaviour of matter in different phases. The kinetic theory of matter is composed of particles which have a certain amount of energy which allows them to move at different speeds depending on the temperature (Energy). There are spaces between the particles and attractive forces between particles when they come closer together. 31 Table 4: Kinetic Molecular Theory for the three states of matter, solid, liquid, and gas. Property of Solid Liquid Gas matter Diagram Atoms or Atoms or Atoms or Particles molecules. molecules. molecules. More energy High energy than in a solid, and are Low energy, Energy and but less energy constantly particles vibrate movement of than a liquid. moving. around a fixed particles. Collisions Collisions point. causing causing pressure. pressure. Very little space between Bigger spaces Spaces Large spaces particles. than solids, but between because of high Particles are smaller than particles. energy. tightly packed gases. together. Weak forces of Very strong Weaker forces attraction Attractive forces forces of of attraction because of the between attraction. than in solids, large empty particles. Solids have a but stronger spaces between fixed volume. than in gases. particles. A gas becomes a liquid or solid when it is Becomes a gas cooled. Solids become if temperature is Particles have liquids or gases increased. less energy, Changes in if their Becomes a move closer phase. temperature is solid if together so increased. temperature is attractive forces decreased. increase, and the gas becomes a solid or a liquid. 32 2.3. EXERCISE 2 1. The following table gives the melting and boiling points of various substances under specific circumstances. Boiling Melting Boiling Melting Substance point Substance point point point (ºC) (ºC) (ºC) (ºC) Water Bromine 0 100 -7 59 (H2O) (Br2) Ethanol -114 78.4 Iodine (I2) 114 184 (C2H6O) Hydrogen Barium (Ba) 704 1 700 chloride -115 -85 (HCl) Ammonia Brass 900 1 100 -78 -33 (NH3) Hydrogen Calcium 850 1 439 fluoride -83 20 (Ca) (HF) Hydrogen Paraffin 52 300 telluride -49 -2 (H2Te) Phosphorus Octane 44 280 -57 126 (P) (C8H18) Potassium Acetic acid 63 762 17 118 (K) (C2H4O2) Chlorine Propanol -101 -35 -127 98 (Cl2) (C3H8O) Use the table of melting and boiling points to determine the state (phase) in which you will find the following: 1.1. Octane at room temperature (25ºC). 1.2. Hydrogen oxide at 110ºC. 1.3. Hydrogen nitride at -50ºC. 1.4. Brass at 1000ºC. 1.5. Phosphorus at 40ºC. 1.6. Chlorine at 20ºC. 1.7. Hydrogen fluoride at 19º. 1.8. Potassium at 765ºC. 1.9. Propanol at -130ºC. 1.10. Iodine at 120ºC. 33 2. In which phase (state) would you find a substance if it has the following property / properties: 2.1. Packed close together in a regular arrangement or lattice. 2.2. Fixed volume but take on the shape of the container. 2.3. Arranged irregularly and spread very far apart. 2.4. Cannot be compressed. 2.5. Have weaker intermolecular forces between the particles. 2.6. Have enough kinetic energy to enable them to move randomly. 2.7. Specific condensation points under standard conditions. 2.8. Exerts pressure in all directions and can be compressed. 3. Study the following diagram and answer the questions that follow: 3.1. Name the phase of A, B and, C. 3.2. Name the processes 1 to 6. 34 4. Study the heating curve of an unknown substance below and answer the questions that follow: 4.1. Name the state(s) in which you would find the substance between the following points: 4.1.1. AB 4.1.2. BC 4.1.3. CD 4.1.4. DE 4.1.5. EF 4.2. Write down the alphabet letter(s) where the following processes will take place. 4.2.1. Melting 4.2.2. Boiling 4.3. What is the: 4.3.1. Boiling point value? 4.3.2. Melting point value? 4.4. Give the chemical formula of the substance. 4.5. How will the average kinetic energy of the particles of the substance be affected between: 4.5.1. AB 4.5.2. BC 4.5.3. EF 35 5. Define the following: 5.1. Diffusion 5.2. Brownian 5.3. Freezing point 5.4. Melting point 5.5. Boiling point 5.6. Melting 5.7. Evaporation 5.8. Freezing 5.9. Sublimation 5.10. Condensation 6. Study the cooling curve of an unknown substance below and answer the questions that follow: 6.1. Name the state(s) in which you would find the substance between the following points: 6.1.1. AB 6.1.2. BC 6.1.3. CD 6.1.4. DE 6.1.5. EF 36 6.2. Name the processes that takes place between: 6.2.1. BC 6.2.2. DE 6.3. What is the temperature at which condensation takes place? 6.4. What is the temperature at which freezing takes place? 6.5. Give the name of the unknown substance (use the table of melting and boiling points in question 1). 6.6. Explain why the temperature between points BC stays the same. 37 UNIT 3: ATOMIC STRUCTURE 3.1. EXAM GUIDELINES Models of the atom Describe the major contributions (Dalton, Thomson, Rutherford, Bohr, and Chadwick) to the atomic model used today. Structure of the atom: protons, neutrons, electrons Define the atomic number as the number of protons in an atom of an element. Given a periodic table or suitable data, determine for an atom/ion the: o Atomic number o Number of protons o Number of electrons o Number of neutrons o Mass number Show that by removing electrons from an atom the neutrality of the atom is changed. Determine the charge on an ion after removing electrons from or adding electrons to an atom. Isotope Define isotopes as atoms of the same element having the same number of protons, but different numbers of neutrons. Define relative atomic mass as the mass of a particle on a scale where an atom of carbon-12 has a mass of 12. Calculate the relative atomic mass of naturally occurring elements from the percentage of each isotope in a sample of the naturally occurring element and the relative atomic mass of each of the isotopes. Represent atoms using the notation AZ E where E is the symbol of the element, Z is the atomic number and A is the mass number. Electron configuration Use Aufbau diagrams (orbital box diagrams) and the electron configuration notation (sp notation) to give electronic arrangements of atoms up to Z = 20. Know that every orbital corresponds to a specific energy value that electrons have when occupying it. Describe atomic orbitals as the 38  most probable regions in space where electrons that have the specific energy corresponding to the orbital are found. Describe the shape of s-orbitals as spherical and that of p-orbitals as pairs of dumb-bells aligned along the x-, y- and z-axes at 90° to each other. State Hund's rule: No pairing in p orbitals before there is not at least one electron in each of them. State Pauli's Exclusion Principle: Maximum of two electrons per orbital provided that they spin in opposite directions. 3.2. NOTES Atoms are the basic building block of all matter. An atom consists out of a nucleus (centre) which contains protons (p+) and neutrons (no). The protons are positively charged. The neutrons are neutral. Protons and neutrons have the same mass (very small), 1,67 × 10−27 kg, called the atomic mass unit, symbol u. The mass of a proton and a neutron is 1u. Protons and neutrons together are called nucleons. There are electrons (e-) that moves constantly around the nucleus in spaces called orbitals. 39 3.2.1. Atomic number (Z) The atomic number is the number of protons in an atom. Indicated by the symbol Z. Which indicates the number of protons in the nucleus of the atom, also indicates the position you will find the atom on the periodic table. 3.2.2. Mass number (A) The mass of an atom is determined by its nucleus, which contains the number of protons and neutrons. The symbol for mass number is A. A = Z + no 3.2.3. Atom Electrons is much smaller than protons and neutrons. 1 The mass of an electron is u. 1840 Electrons is negatively charged. Electrons are found in a space around the nucleus. The positive charge on a proton is equal to the negative charge on an electron. A neutral atom has equal number of protons and electrons. That will make the number of electrons is also equal to the atomic number (Z). e.g., 168O O is the symbol for the oxygen atom. The mass number (A) of the oxygen atom is 16. The atomic number of (Z) oxygen is 8. There are 8 protons in the nucleus of the oxygen atom. The oxygen atom is neutral. There are 8 electrons in the space around the nucleus. There are 8 neutrons in the nucleus of the atom. (no = A – Z) 40 IONS Ions are small, charged particles. A neutral atom either lose an electron (positively charged) or gain an electron (negatively charged). A positive ion is known a cation has less electrons than protons. A negative ion known as an anion has more electrons than protons. The number of protons and neutrons remains the same in the nucleus. e.g., p+ =8 e- = 10 (8 + 2; gained 2 electrons to have a 2- charge) no = 16 – 8 = 8 e.g., p+ = 56 e- = 54(56 - 2; lost 2 electrons to have a 2+ charge) no = 137 – 56 = 81 41 3.2.4. Isotopes Isotopes is atoms of the same elements having the same number of protons, but different numbers of neutrons. 35 e.g., 17𝐶𝑙 17p+ and 18no = 35 u 37 17𝐶𝑙 17p+ and 20no = 37 u ∴ Both have seventeen protons, but the number of neutrons is different. If there are different isotopes, the mass number is different which means the atom will have a relative atomic mass. 3.2.4.1. Relative atomic mass The Relative atomic Mass is the mass of a particle where an atom of C-12 has a mass of 12. C-12 atom was chosen as the standard and one C-12 has a mass of 12 atomic mass units (u). In a sample of C, there will be atoms with different masses. 1 out of a 100 C-atoms might be a C-13. 1 out of millions of C-atoms might be a C-14. The rest will be C-12 atoms which makes the average atomic mass of C-atoms just over 12. e.g., In a sample of a 100 Cl-atoms 75 is Cl-35 atoms and 25 Cl-37 atoms. (75×35)+(25×37) 𝐴𝑟 = 100 𝐴𝑟 = 35,5 𝑢 Or 𝐴𝑟 = 35,5 gmol-1 17 ∴ 35,5Cl 42  124 e.g., Element X is found naturally in three forms, namely 25% of 248X; 25% of 124 124 250X and 50% of 249X. What is the relative atomic mass of element X? (25×248)+(25×250)+(50×249) 𝐴𝑟 = 100 𝐴𝑟 = 249 𝑢 Or 𝐴𝑟 = 249 gmol-1 ∴ 124 249X 3.2.5. Electron configuration Electrons are negatively charged particles. Electron’s mass is 9,11×10-31 kg, which is very small. Electron’s move in an area around the nucleus of the atom called orbitals. Electrons don’t move randomly around the nucleus. The way in which electrons are arranged around the nucleus is called the atom’s electron configuration. Only electrons are involved in the changes that happen during chemical reactions. 3.2.5.1. Orbitals Electrons are found in different energy levels that contains sub- shells called orbitals. Each energy level contains sub-shells called orbitals. Four types: s; p; d; and f. 43 44  s-orbitals has a spherical shape. p-orbitals has a dumbbell shape along the 𝑥-; 𝑦- and 𝑧-axes at 90o to each other. Each orbital can only contain 2 electrons. Each energy level has a specific number of orbitals. Each orbital can only take 2 electrons. We only do the following (first 20 elements) 1st Energy level - 1 orbital - 2 electrons nd 2 Energy level - 4 orbitals - 8 electrons 3rd Energy level - 4 orbitals - 8 electrons th 4 Energy level - 1 orbital - 2 electrons 45 3.2.5.2. Aufbau Orbitals are represented by blocks (or circles) in an energy level. Electrons in the orbitals are represented by arrows. s-orbital fill up first before p-orbitals get 1 electron each. The electrons in p-orbitals are found in p𝑥 -, p𝑦 – and p𝑧 – orbitals with a maximum of 2 electrons in each. HUND’S RULE: No pairing in p-orbitals before there is at least5 one electron in each of them. 46 PAULI’S EXCLUSION PRINCIPLE: Maximum of two electrons per orbital provided that they spin in opposite directions. 3.2.5.3. sp-notation (electron configuration notation) e.g., for Phosphorus atom 𝟏𝟓 𝟑𝟏𝐏 Z = 15 15p+ = 15 e- A = 31 1s2;2s22p6;3s23p3 The sum of the numbers 2+2+6+2+3 = 15, corresponds with the atomic number of phosphorus. e.g., for Oxygen atom 𝟏𝟔𝟖𝐎 Z=8 8p+ = 8 e- A = 16 1s2;2s22p6;3s23p3 The sum of the numbers 2+2+4 = 8, corresponds with the atomic number of oxygen. 47 3.2.5.4. Abbreviated electron configuration. Because electron configuration for Lithium is the same as helium plus one electron in the 2s-orbital, the electron configuration can be written as follows: 𝟐 𝟒𝐇𝐞 : 1s2 𝟑 𝟕𝐋𝐢 : 1s2; 2s1 ∴ 𝟑𝟕𝐋𝐢: [He] 2s1 𝟏𝟎 2 2 6 𝟐𝟎𝐍𝐞 : 1s ; 2s 2p 𝟏𝟐 2 2 6 𝟐𝟒𝐌𝐠 : 1s ; 2s 2p ; 3s2 ∴ 𝟏𝟐 𝟐𝟒𝐌𝐠: [Ne] 3s 2 3.3. EXERCISE 3 1. Define the following terms: 1.1. Atomic number. 1.2. Isotopes. 1.3. Relative atomic mass. 2. Explain what is meant with the following terms: 2.1. Protons. 2.2. Neutrons. 2.3. Electrons. 2.4. Nucleons. 2.5. Nucleus. 2.6. Orbital. 48 3. Copy and complete the following table: Number Atomic Atomic Number of Number of Symbol of Atom / ion mass number electrons neutrons ( 𝑨𝒁𝑬) protons (A) (Z) (e-) (nº) (p+) 9 4𝐵𝑒 Beryllium 9 4 4 4 5 27 3+ Aluminium 13𝐴𝑙 ion 27 13 13 10 14 24 12𝑀𝑔 3.1. 24 12 3.2. 3.3. 3.4. 37 − 17𝐶𝑙 3.5. 3.6. 3.7. 17 18 3.8. 3.9. Phosphorus 31 15 3.10. 3.11. 3.12. 3.13. Sodium ion 23 3.14. 11 10 3.15. 3.16. 3.17. 3.18. 3.19. 16 16 16 14 3− 7𝑁 3.20. 14 7 3.21. 3.22. 3.23. 59 3+ 26𝐹𝑒 Ferrous ion 3.24. 3.25. 26 3.26. 3.27. 4. Calculate the relative atomic mass for the following: 4.1. Iron if there is in a 100g sample, 5.845% of 54 26𝐹𝑒 , 91.754% of 56 57 58 26𝐹𝑒 , 2.119% of 26𝐹𝑒 , and 0.286% of 26𝐹𝑒. 4.2. Neon if there is in a 100g sample, 90.48% of 20 10𝑁𝑒, 0.27% of 21 22 10𝑁𝑒, and 9.25% of 10𝑁𝑒. 5. State the following: 5.1. Hund’s rule. 5.2. Pauli’s exclusion principle. 6. Draw the Aufbau diagrams for the first 20 elements. 7. Give the sp-notation for the first 20 elements. 8. Give the abbreviated sp-notation of atoms nr 11, 12, 16, 17, 19, and 20. 49 9. Answer the following questions on an element represented by Y 82 with the symbol 207 𝑌. 9.1. How many electrons does this atom have? 9.2. How many neutrons does this atom have? 9.3. How many protons does this atom have? 9.4. What is the atomic mass of this atom? 9.5. What i9s the atomic number of this atom. 9.6. How many electrons will an ion of this atom have with a charge of 2+? 9.7. Give the name of this atom. 9.8. Give the correct symbol of this atom. 10. Draw the Aufbau diagrams of the ions of the following atoms: 10.1. Lithium. 10.2. Magnesium. 10.3. Oxygen. 10.4. Chlorine. 11. Give the abbreviated sp-notion of the ions of the following atoms: 11.1. Sodium. 11.2. Beryllium. 11.3. Fluorine. 11.4. Sulphur. 50 UNIT 4: PERIODIC TABLE 4.1. EXAM GUIDELINES The positions of the elements in the periodic table related to their electronic arrangements. Describe the periodic table as displaying the elements in order of increasing atomic number and showing how periodicity of the physical and chemical properties of the elements relates to atomic structure. Define the group number and the period number of an element in the periodic table. Groups are the vertical columns in the periodic table. Some groups have names, e.g, alkali metals (group I), earth-alkaline metals (group II), halogens (group 17 or VII) and noble gases (group18 or VIII). Periods are the horizontal rows in the periodic table. Relate the position of an element in the periodic table to its electronic structure and vice versa. Describe periodicity from Li to Ar in terms of atomic radius, ionisation energy, electron-affinity, and electronegativity. Describe the changes in terms of change in charge of the nucleus and distance between the nucleus and the electron. Periodicity is the repetition of similar properties in chemical elements, as indicated by their positioning in the periodic table. Define atomic radius, ionisation energy, electron-affinity, and electronegativity. Atomic radius: Radius of an atom, i.e., the mean distance from the nucleus to the border of the outer orbital. Ionisation energy: Energy needed per mole to remove an electron(s) from an atom in the gaseous phase. First ionisation energy: Energy needed per mole to remove the first electron from an atom in the gaseous phase. Electron affinity: The energy released when an electron is attached to an atom or molecule to form a negative ion. Electronegativity: A measure of the tendency of an atom in a molecule to attract bonding electrons. 51 Similarities in chemical properties among elements in Groups 1, 2, 17 and 18 Relate the electronic arrangements to the chemical properties of group 1, 2, 17 and 18 elements. Describe the trend in reactivity of elements in groups 1, 2 and 17. Groups 1 and 2: Chemical reactivity increases from top to bottom. Group 17: Chemical reactivity decreases from top to bottom. Predict chemical properties of unfamiliar elements in groups 1, 2, 17 and 18 of the periodic table. Indicate that metals are found on the left-hand side of the periodic table. Indicate that non-metals are found on the right-hand side of the periodic table. Indicate where transition metals are to be found on the periodic table. 4.2. NOTES Dimitri Mendeleev, a Russian chemist, known as the father of the Periodic Table, was the first to place elements in order of increasing, atomic mass. Other chemist before him have grouped together elements with similar properties. The first Periodic Table was printed in 1869. 52 4.2.1. Periodic table as on gr 10 -12 data sheets. Elements are arranged in seven horizontal rows (periods), in order of increasing atomic number from left to right and from top to bottom. These rows are called periods which also indicates the energy levels where you find the electrons. Elements with similar chemical properties form vertical columns called groups, 1-18 or I – VIII Groups 1 (I), 2(II) and 13 (III) through to 18 (VII) are main group elements. (The main groups are indicated with the roman numbers as 1 to 8). Groups 3 to 12 are in the middle of the Periodic Table and are called the transition metals. Group 1 (I) - alkali metals Group 2 (II) - alkali earth metals Group 17 (VII) - halogens Group 18 (VIII) - noble gases When halogens react with metals, they form ionic solids called halides. The halide names end in –ide, for example fluoride and chloride. 53  Oxygen combines with almost all elements; we can compare the properties of oxides to see how metals differ from metalloids and non-metals. The oxides of metals are ionic compounds that have high melting and boiling points. They dissolve in water to form basic solutions. The oxides of non-metals form small molecules that can dissolve in water to form acids. GROUPS: are the vertical columns in the periodic table. PERIODS: are the horizontal rows in the periodic table. 4.2.2. Electron configuration and the Periodic Table The electron configuration of atoms is linked to their position in the Periodic Table. Each consecutive (following in numbers) atom has one more electron. The number of electrons in its highest (outer most) energy level is the same as the group number that it is in, this called the valence electrons. Group I: only have 1 electron in its highest energy level. Group II: have 2 electrons in its highest energy level. Group III: have 3 electrons in its highest energy level. Group IV: have 4 electrons in its highest energy level. Group V: have 5 electrons in its highest energy level. Group VI: have 6 electrons in its highest energy level. Group VII: have 7 electrons in its highest energy level. Group VIII: have 8 electrons in its highest energy level. E.g., Aluminium is in group III (13): 1s2; 2s22p6; 3s23p1, its highest energy level is 3, which is the period (row) it’s found in, and it only have 3 electrons (3s23p1) in its highest energy level, which is the same as the group number III and the amount of valence electrons. 54  Each group have a full set of inner electrons that was filled in the previous period and that are indicated by the symbol of the preceding noble gas (previous row). Each group’s valence electron structure is the same. The number of valence electrons of an element determines its chemical properties because they are involved in chemical bonding which makes chemical reactions possible. The noble gases highest energy level is completely filled, which means they do not react with other elements, which makes them inert (they do not react). Elements with incomplete energy levels wants to be completely filled, because they are unstable, and seek to change to a noble gas electron configuration so that their outermost orbitals are completely filled. That is why ions form: - negative ion – adding of electrons. - positive ion – loosing of electrons. 55 4.2.3. Periodicity Periodicity is the repetition of similar properties in chemical elements as indicated by their positioning on the Periodic Table. 4.2.3.1. Atomic radius The size of an atom on the Periodic Table decreases from left to right and increase from top to bottom. ATOMIC RADIUS: the radius of an atom, the mean distance from the nucleus to the border of the outer orbital. The nucleus of an atom is positively charged, because of the protons, and attract the electrons in the orbitals around the nucleus. The higher the atomic number the higher the positive charge of the nucleus the more it will attract the electrons which means the atomic radius becomes smaller from left to right in the Periodic Table or the atomic radius increase from right to left. The atomic radius increases from top to bottom on the Periodic Table. 56 4.2.3.2. Ionisation energy IONISATION ENERGY: Energy needed per mole to remove an electron(s) from an atom in the gaseous phase. FIRST IONISATION ENERGY: Energy needed per mole to remove the first electron from an atom in the gaseous phase. When an electron is removed from an atom, the atom gains a positive charge, become a cation. To remove an electron energy is needed which is known as the ionisation energy. The ionisation energy increase from left to right in the Periodic Table. Which means the energy needed to remove an electron from the non-metals is high. That’s why metals will more likely give electrons away and non-metals will gain electrons. Ionisation energy decrease from top to bottom. The outer electrons are further away, which means less energy needed. 57 4.2.3.3. Electron affinity ELECTRON AFFINITY: the energy released when an electron is attracted to an atom or molecule to form a negative ion. The atom’s ability to accept one or more electrons. Electron affinity is the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion. X(g) + e- → X-(g) + electron affinity Electron affinity increases from left to right on the Periodic Table (period). The values vary little within a group. 4.2.3.4. Electron negativity. ELECTRON NEGATIVITY: a measure of the tendency of an atom in a molecule to attract the bonding electrons. Electron negativities increase from left to right across a period on the Periodic Table. Electron negativities decrease from top to bottom in s group on the Periodic Table. 58 4.3. EXERCISE 4 1. Define the following terms: 1.1. Group number. 1.2. Period number. 1.3. Atomic radius. 1.4. Ionisation energy. 1.5. Electron affinity. 1.6. Electronegativity. 2. Fill in the missing words: 2.1. Elements are arranged in seven horizontal rows, in order of ______ atomic number from ___________ and from. 2.2. Elements with _________ chemical properties form _______ columns called groups. 3. Give the name of the following group(s): 3.1. VIII 3.2. VII 3.3. II 3.4. I 3.5. Nr 3 -12 59 4. The electron configuration of atoms is linked to their … 5. The amount of electrons in its highest energy level is called the … 6. How many electrons are found in the highest energy level of group…? 6.1. VIII 6.2. V 6.3. III 6.4. VI 6.5. IV 7. What determines the chemical property of an element? 8. How does a 8.1. Positive ion form? 8.2. Negative ion form? 9. The repetition of similar properties in chemical elements on the periodic table is called? 10. Name the positively charged region of an atom. 11. Explain how atoms are arranged on the periodic table according to their… 11.1. Atomic radius. 11.2. Ionisation energy. 11.3. Electron affinity. 11.4. Electronegativity. 12. 20 Elements have been given imaginary chemical symbols in the form of the letters of the alphabet, A – T. Using a Periodic Table and the clues given below, place these 20 letters in the correct place on the table given below. e.g. U – Ferromagnetic metal also used in nails. 60  I II III IV V VI VII VIII U The following elements belong to the same group: CHJL; DMR; FNP; AB; EQ; IO; KS; GT J - the non-metal in its group. L - the biggest atomic radius. C - found in period two of the periodic table. M - An element used in the light signage of a roadhouse or restaurant to stand out at night. D - The smallest atom in its group. R - an atom with 18 protons in its nucleus. N - A metal that burns with a blinding white light. P - An element necessary for strong bones. F - An element that have 5 neutrons in its nucleus. A - A halogen with an electronegativity of 3,0. B - An element used in toothpaste. E - A metal used in producing of food containers. Q - The element with the highest ionisation energy in its group. O - An element used in Silicon chips. I - An element in its crystal form is the hardest substance on earth. T - An element with an atomic number of 15. G - An element in group 15 and period 2. K - An atom with an atomic mass of 32. S - An element in its compound form is necessary for combustion. 61 UNIT 5: CHEMICAL BONDING 5.1. EXAM GUIDELINES Covalent bonding, ionic bonding, and metallic bonding Define a chemical bond as a mutual attraction between two atoms resulting from the simultaneous attraction between their nuclei and the outer electrons. (The energy of the combined atoms is lower than that of the individual atoms resulting in higher stability.) Draw Lewis dot diagrams of elements. A Lewis dot diagram is a structural formula in which valence electrons are represented by dots or crosses. It is also known as an electron dot formula, a Lewis formula, or an electron diagram. Define a covalent bond as the sharing of electrons between atoms to form molecules. Molecule: A group of two or more atoms that are covalently bonded and that functions as a unit. Draw Lewis dot diagrams of simple covalent molecules containing single, double, and triple covalent bonds: H2; F2, Cℓ2, O2, N2, HF, HCℓ, CH4, NH3, H2O In a Lewis dot diagram two dots between atoms represent a covalent bond. These two electrons are known as a bonding pair, whilst non- binding electron pairs are called lone pairs. Write names and formulae of covalent compounds. Define ionic bonding as the transfer of electrons to form cations (positive ions) and anions (negative ions) that attract each other to form a formula-unit. A formula-unit is the simplest empirical formula that represents the compound. An ion is a charged particle made from an atom by the loss or gain of electrons. An anion (negative ion) is a charged particle made from an atom by the gain of electrons. A cation (positive ion) is a charged particle made from an atom by the loss of electrons. Draw Lewis dot diagrams of cations and anions. Draw Lewis dot diagrams to show the formation of simple ionic compounds such as NaCℓ, KCℓ, KBr, CaCℓ2 and MgBr2. Predict the ions formed by atoms of metals and non-metals by using information in the periodic table. Metals occur on the left-hand side of 62  the periodic table and form positive ions, whilst non-metals occur on the right-hand side of the periodic table and form negative ions. Name ionic compounds based on the component ions. Describe the structure of the sodium chloride crystal. In the crystal each sodium ion is surrounded by six chloride ions to form a cubic structure. Each chloride ion is also surrounded by six sodium ions. A crystal lattice: An orderly three-dimensional arrangement of particles (ions, molecules, or atoms) in a solid structure. Define metallic bonding as the bond between positive ions and delocalised valence electrons in a metal. Valence electrons or outer electrons are the electrons in the highest energy level of an atom in which there are electrons. Calculate relative molecular masses for covalent molecules, e.g., Mr (HCℓ) = 35, 5. Calculate relative formula masses for ionic compounds, e.g., Mr (NaCℓ) = 57, 5. 5.2. NOTES 5.2.1. Chemical bonding Chemical bonding is an important process because it allows different combinations of atoms and molecules. A chemical bond is formed when electrons are shared between atoms or when electrons were transferred between atoms. A CHEMICAL BOND: is a mutual attraction between two atoms resulting from the simultaneous attraction between nuclei and the outer electrons. Chemical bonds between atoms forms new compounds. When the atoms are close enough to each other the attractive forces between the nucleus of the one atom and the valence electrons of the other atom becomes stronger than the repulsive forces between the atom’s orbitals. 63  A simple diagram of a chemical bond that forms. The compound forms a stable structure. (Noble gas structure, outer orbitals filled completely). The compound now has a lower potential energy. There are three different types of chemical bonds. o Covalent bonds – between non-metals. o Ionic bonds – between metals and non-metals. o Metallic bonds – between metals. 5.2.2. Lewis dot diagram The structural formula in which valence electrons are either represented as dots around the atom or as crosses. The symbol of the element is used to represent the nucleus (which contains the protons and neutrons) and the inner (core) electrons of the atom. VALENCE ELECTRONS: the amount of electrons in its highest (outer most) energy level in which there are electrons. The valence electrons are equal to the group number (column) of the atom found on the Periodic Table. The highest energy level is equal to the period (row) on the Periodic Table. 64 E.g. Phosphorus 1531𝑃 is found in period 3 and group 5 (v) on the Periodic Table. A neutral atom of phosphorus will have 15 electrons over 3 energy levels and 5 valence electrons. Valence electrons 1s2;2s22p6;3s23p3 Fluorine 199𝐹 is found in period 2 and group 7 (VII) on the Periodic Table. A neutral atom of fluorine will have 9 electrons over 2 energy levels and 7 valence electrons. Valence electrons 1s2;2s22p5 5.2.3. Covalent bond COVALENT BOND: the sharing of electrons between atoms to form molecules. The outermost (highest energy level) orbitals of the atoms overlap so that unpaired electrons in each of the bonding atoms can be shared. The outermost orbitals overlap in such a way that the unpaired electrons in each of the bonding atoms can be shared, which means the outer most energy level are filled. Normally following the “octet rule” where most the of the atoms now appear as if they all have eight electrons around them, exceptions is Hydrogen which will now appear as if it’s 1s-orbital is now completely filled. MOLECULE: a group of two or more atoms that are covalently bonded and that functions as a unit. 65 Unpaired electrons are involved with chemical bonding. Paired electrons also known as lone pairs. e.g. Overlapping of Single covalent bond unpaired electrons Double covalent bond Overlapping of unpaired electrons Triple covalent bond Overlapping of unpaired electrons Overlapping of unpaired electrons Two single covalent bonds 66 The bonding electron pairs is closer to Sulphur because sulphur has the higher electronegativity (EN). ENH = 2.1 ENS = 2.5 Two double covalent bonds The bonding pairs are closer to the Oxygen atoms because Oxygen has the higher electronegativity (EN). ENC = 2.5 ENO = 3.5 5.2.3.1. Polar or Non-polar Non-polar bonds: Atoms attract the shared pair of electrons equally; they have the same electronegativity. Polar bond: One atom (highest electronegativity) attracts the shared pair of electrons more than the other atom (lowest electronegativity). 5.2.3.2. Properties of Covalent bonds Melting and boiling points are generally lower than that of ionic compounds. Generally, more flexible than ionic compounds. The molecules are able to move around to some extent. Generally, not soluble in water for example, plastics. Do not conduct electricity when dissolved in water. 67 5.2.4. Ionic bond The transfer of electrons to form cations (positive ions) and anions (negative ions) that attract each other to form a formula unit. Instead of sharing electrons between the atoms it is a transfer of electrons between the atoms, so that each atom acquires a completely filled highest energy level. The following three stages are identified: Metal atom(s) give away electron(s) in their highest energy level to become positively charged known as cations. E.g. Non-metal atom(s) receive electron(s) to fill its highest energy level to become negatively charged known as anions. E.g. The positive and negative ions attract each other through electrostatic forces to form an ionic bond. E.g. KF – Potassium fluoride 68  Two fluoride ions formed. CaF2 – Potassium fluoride. Metals occur on the left-hand side of the Periodic table and forms positive ions. Non-metals occur on the right-hand side of the Periodic table and forms negative ions. 5.2.4.1. Crystal lattice An orderly three-dimensional arrangement of particles (ions, molecules, or atoms) in a solid structure. The arrangement of ions in a regular, geometric structure is called a crystal lattice/ Sodium chloride crystal In the crystal each sodium ion is surrounded by six chloride ions is surrounded by six chloride ions to form a cubic structure. Each chloride ion is also surrounded by six sodium ions. 69 5.2.4.2. Properties of Ionic compounds Arranged in a lattice structure. Ionic solids are crystalline at room temperature. Strong electrostatic attraction. Hard and have high melting and boiling points. Are brittle and bonds are broken along planes when the compound is put under pressure. Solid crystals do not conduct electricity, but ionic solutions do. 5.2.5. METALLIC BOND THE BOND BETWEEN POSITIVE IONS AND DELOCALISED VALENCE ELECTRONS IN A METAL. Bonds between atoms of metals. Atoms are closely packed so that the outermost orbitals overlap. The valence electrons are delocalised (removed from their own atoms) and can move in adjacent orbitals. Positive core ions from a compact crystal lattice. 5.2.5.1. Properties of metals Metals are shiny. Conduct electricity because electrons are free to move. Conduct heat because the positive nuclei are packed closely together and can easily transfer the heat. High melting point because the bonds are strong and high density because of the tight packing of the nuclei. 70 5.2.6. Calculate relative molecular masses for covalent molecules. 𝑀𝑟 (𝐻2 𝑂) = (1 × 2) + 16 = 18 𝑔 ∙ 𝑚𝑜𝑙 −1 𝑀𝑟 (𝐶𝐻4 ) = 12 + (1 × 4) = 16 𝑔 ∙ 𝑚𝑜𝑙 −1 𝑀𝑟 (𝐻2 𝑆𝑂4 ) = (1 × 2) + 32 + (16 × 4) = 98 𝑔 ∙ 𝑚𝑜𝑙 −1 5.2.7. Calculate relative formula masses for ionic compounds. 𝑀𝑟 (𝐾𝐶𝑙) = 39 + 35.5 = 74.5 𝑔 ∙ 𝑚𝑜𝑙 −1 𝑀𝑟 (𝑁𝑎𝑂𝐻) = 23 + 16 + 1 = 40 𝑔 ∙ 𝑚𝑜𝑙 −1 5.3. EXERCISE 5 1. Define the following terms: 1.1. Chemical bond. 1.2. Covalent bond. 1.3. Molecule. 1.4. Ionic bonding. 1.5. Formula unit. 1.6. Ion. 1.7. Anion. 1.8. Cation. 1.9. Metallic bonding. 2. When does a chemical bond form? 3. What happens to the attractive forces between the nucleus of one atom and the valence electrons of another atom when the atoms are close enough to each other? 4. Draw Lewis dot diagrams for: 4.1. F2 4.2. Cl2 4.3. O2 4.4. HF 4.5. CH4 71 4.6. NH3 4.7. H 2O 5. Explain what a bonding pair of electrons is. 6. A non-bonding pair of electrons is also known as a … 7. Name four properties of covalent bonds. 8. Draw Lewis dot diagrams to show the formation of ionic compounds. 8.1. NaCl 8.2. KCl 8.3. KBr 8.4. CaCl2 8.5. MgBr2 9. Complete the following table and write the ionic compound names: OH- SO42- CO32- NO3- Sodium Na+ hydroxide Potassium K+ sulphate Mg2+ Ca2+ 10. Describe what is a crystal lattice. 11. Explain what is meant by valence (outer) electrons. 12. Name five properties of Ionic compounds. 13. Name four properties of metals. 72 14. Calculate the relative molecular masses of: 14.1. Cl2 14.2. HF 14.3. CH4 14.4. C3H18 14.5. NH3 15. Calculate the relative formula mass of: 15.1. NaCl 15.2. CaCl2 15.3. Mg(OH)2 15.4. CaCO3 15.5. H2SO4 15.6. NH4OH 15.7. (NH4)2SO3 73  MODULE 2 CHEMICAL CHANGE UNIT 1: PHYSICAL AND CHEMICAL CHANGE 1.1. EXAM GUIDELINES Separation of particles in physical and chemical change Define a physical change as a change in which: o No new substances are formed. o Energy changes are small in relation to chemical changes. o Mass, numbers of atoms and molecules as being conserved. Describe the rearrangement of molecules during physical changes, e.g. o Molecules separate when water evaporates to form water vapour. o When ice melts molecules become disorderly arranged due to breaking of intermolecular forces Define a chemical change as a change in which: o New chemical substances are formed. o Energy changes are much larger than those of the physical change. Endothermic reaction: Energy is absorbed during the reaction. Exothermic reaction: Energy is released during the reaction. o Mass and atoms are conserved, but the number of molecules is not. Describe examples of a chemical change that include the: o Decomposition of hydrogen peroxide to form water and oxygen o Synthesis reaction that occurs when hydrogen burns in oxygen to form water. o Heating of iron and sulphur o Reaction of lead (II) nitrate and potassium iodide (in solid phase and/or as solutions) o Titration of hydrochloric acid with sodium hydroxide to measure the change in temperature. 74 Conservation of atoms and mass Calculate relative molecular masses of reactants and products in balanced equations to illustrate that atoms are conserved during chemical reactions, but not molecules. Representing Chemical Change Balanced chemical equations. Write and balance chemical equations. Use formulae with subscripts to represent phases, viz. (s), (ℓ), (g) and (aq). Interpret balanced reaction equations in terms of: o Conservation of atoms. o Conservation of mass (use relative atomic masses). 1.2. NOTES 1.2.1.Separation of particles in physical and chemical change. Substances (chemicals) undergoes several changes in different situations. These changes can either be physical or chemical. 1.2.1.1. Physical change. No new substances are formed. Energy changes are small in relation to chemical change. Mass, numbers of atoms and molecules are conserved. Physical change is when a substance ‘s formula remains the same throughout the change. E.g., water stays H2O. 75  Image from Istock The formula of water, H2O stays the same throughout the physical (phase) change. When the ice melts the crystal lattice becomes less arranged because the intermolecular forces (forces of attraction that’s between the water molecules to keep them close) are overcome and the solid turns into a liquid. When water liquid turns into water vapour it gained energy, either some molecules because of evaporation or all the molecules because of boiling. 1.2.1.2. Chemical change. New chemical substances are formed. Energy changes are much larger than those of physical changes. A chemical change is a change that involves the transformation of one or more substances into one or more different substances. 76  The change of the substance (particle or molecules) forms a total new substance (particle or molecule). The formula of the substance will be different. We have reactants that change into products a new substance with different properties. The reactants are a substance a reaction start with. The products are the substance a reaction ends with. The energy transfer in chemical reactions are much higher than in physical changes. The strong interatomic force within a molecule has to be broken which needs more energy. e.g., water forms from the combustion of hydrogen gas. Hydrogen + Oxygen → water The substances at the left-hand side of the arrow are the reactants. The substances at the right-hand side of the arrow are the products. 2H2(g) + O2(g) → 2H2O(ℓ) 77  The mass and the atoms are conserved but not the number of molecules. Reactants 2 H2-molecules ∴ 4 H-atoms 1 O2-molecule ∴ 2 O-atoms ∴ 3 molecules reactants ∴ 6 atoms Products 2 H2O-molecules ∴ 4 H-atoms and 2 O-atoms 2 molecules products ∴ 6 atoms The combustion of hydrogen gives heat off during the reaction, which makes it an exothermic reaction. Exothermic reaction: Energy is released during the reaction. e.g., The heating of a mixture of iron fillings and sulphur powder. Fe(s) + S(s) → FeS(s) When the reaction mixture is heated it starts to turn red. The grey iron fillings and the yellow sulphur powder mixture then turns darker grey to form the iron sulphide. The iron that was magnetic is nonmagnetic as part of the iron sulphide. With a chemical change the products have different properties than the reactants. The compound iron sulphide cannot be separated by any physical means, whereas a mixture of iron and sulphur can be separated using a magnet. 78  Drawing form alamy In the reaction of iron and sulphur to form iron sulphide heat is taken in, which makes it an endothermic reaction. Endothermic reaction: Energy is absorbed during the reaction. Decomposition: The breaking down of chemical substances into smaller chemical substances. Usually a big molecule (reactant) breaks up into two or more smaller molecules. e.g., heating