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Chemistry IGCSE Textbook

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IGCSE Chemistry Cambridge International Examinations 2016 Syllabus

NES/Chemistry/IGCSE Cambridge International Examinations IGCSE Chemistry Syllabus Code 0620 2016 Syllabus Update Paper Type of Paper Duration Marks Weighting 2 Multiple Choice 45 Minutes 40 30% 4 Written Paper 1 Hour and 15 Minutes 80 50% 20% 6 Alternative to 1 Hour 40 Practical A Gru and minion production. 1

NES/Chemistry/IGCSE Table of Contents Topic Topic Name Page Number The particulate nature of matter 4 Experimental techniques 13 1 Measurement 14 2 Criteria of purity 16 2.1 Methods of purification 20 2.2.1 28 2.2.2 Atoms, elements and compounds 31 3 Atomic structure and the Periodic Table 36 3.1 38 3.2.1 Bonding: the structure of matter 49 3.2.2 Ions and ionic bonds 54 3.2.3 57 3.2.4 Molecules and covalent bonds 59 3.2.5 Macromolecules 73 4.1 Metallic bonding 89 4.2 Stoichiometry 108 5 109 6 The mole concept 114 6.1 Electricity and chemistry 117 6.2 119 7 Chemical energetics 122 7.1 Energetics of a reaction 137 7.2 146 7.3 Energy transfer 151 7.4 Chemical reactions 153 8 Physical and chemical changes 162 8.1 Rate (speed) of reaction 164 8.2 Reversible reactions 172 8.3 8.4 Redox Acids, bases and salts The characteristic properties of acids and bases Types of oxides Preparation of salts Identification of ions and gases 2

NES/Chemistry/IGCSE Table of Contents Topic Topic Name Page Number The Periodic Table 176 The Periodic Table 178 9 178 9.1 Periodic trends 179 9.2 Group properties 183 9.3 Transition elements 184 9.4 185 9.5 Noble gases 187 10 Metals 189 10.1 194 10.2 Properties of metals 199 10.3 Reactivity series 200 10.4 202 11 Extraction of metals 203 11.1 Uses of metals 210 11.2 Air and water 213 11.3 Water 215 11.4 Air 218 12 221 13 Nitrogen and fertilisers 225 14 Carbon dioxide and methane 231 14.1 233 14.2 Sulfur 235 14.3 Carbonates 237 14.4 Organic chemistry 244 14.5 Names of compounds 249 14.6 253 14.7 Fuels 256 14.8.1 Homologous series 258 14.8.2 14.8.3 Alkanes Alkenes Alcohols Carboxylic acids Polymers Synthetic polymers Natural polymers 3

NES/Chemistry/IGCSE Topic 1 - The Particulate Nature of Matter  State the distinguishing properties of solids, liquids and gases  Describe the structure of solids, liquids and gases in terms of particle separation, arrangement and types of motion  Describe changes of state in terms of melting, boiling, evaporation, freezing, condensation and sublimation  Explain changes of state in terms of the kinetic theory  Describe qualitatively the pressure and temperature of a gas in terms of the motion of its particles  Show an understanding of the random motion of particles in a suspension (sometimes known as Brownian motion) as evidence for the kinetic particle (atoms, molecules or ions) model of matter  Describe and explain Brownian motion in terms of random molecular bombardment  State evidence for Brownian motion  Describe and explain diffusion  Describe and explain dependence of rate of diffusion on molecular mass Diffusion is the spreading of one substance through another from a region of high concentration to a region of low concentration due to the continuous random motion of particles 4

NES/Chemistry/IGCSE States of Matter All matter is made of tiny particles and these particles have energy, which causes them to vibrate and/or move. This type of energy is called kinetic energy. There are 3 states of matter: solid, liquid and gas Property Solid Liquid Gas Shape definite shape takes the shape of takes the shape of the bottom of the the whole container container Volume definite definite changes to fill the whole container Arrangement of particles are packed particles are still large distance particles close together in a close together, but between particles regular arrangement Movement of with a random and have a particles they vibrate about a arrangement random fixed point Density particles can slip and arrangement slide over each other particles move in a in a continuous continuous rapid random motion random motion high due to the close less than solids low packed arrangement of particles Forces of very strong weak virtually non attraction existent Diffusion little diffusion slow diffusion rapid diffusion Compressible no no, because the easily compressed particles are packed due to the large very close together distance between particles which can be reduced 5

NES/Chemistry/IGCSE Changing State The 3 states of matter are interchangeable which means a solid can be changed to a liquid and this to a gas. sublimation melting boiling / evaporation solid liquid gas freezing / solidifying condensing sublimation Melting point is the temperature at which a pure solid changes to a liquid without a change in temperature. (water = 0oC, ethanol = -117oC) Boiling point is the temperature at which a pure liquid changes to vapour without a change in temperature (water = 100oC, ethanol = 78oC) Evaporation is when liquids change into gases over a range of temperature below the boiling point. Solidifying point/freezing point is the temperature at which a pure liquid changes to a solid without a change in temperature (water = 0oC, ethanol = -117oC) Condensation point is the temperature at which vapour changes to a liquid without a change in temperature (water = 100oC, ethanol = 78oC) 6

NES/Chemistry/IGCSE Diagrams of Changing State solid liquid gas heat heat Sublimation This is where a solid changes straight to a gas, without changing to a liquid, or a gas changes to a solid without changing to a liquid. The process for both directions is called sublimation.  Example 1: iodine  Example 2: carbon dioxide  Example 3: graphite  Example 4: ammonium chloride Heat Energy Process Change in Heat Energy Exothermic / Endothermic Melting gain endothermic Boiling gain endothermic Condensing loss exothermic Freezing loss exothermic Sublimation loss or gain (depends on process) exothermic or endothermic 7

Kinetic Particle Theory NES/Chemistry/IGCSE When a solid is heated, the particles vibrate faster force of about a fixed point. This causes the particles to attraction move further apart and so the solid expands. When the particles gain sufficient energy to overcome the strong forces of attraction holding them together, they can move out of their fixed positions. They can slip and slide over each other in a continuous random motion. When this happens the solid melts. The particles in the liquid are still close to each other. They have enough kinetic energy to move around each other closely, but do not have enough energy to overcome the forces that hold them close to each other. increasing temperature If more heat energy is supplied to the particles, they move faster until they have enough energy to overcome the forces holding them together. The particles then escape from the liquid surface and move around in a continuous rapid random motion. The liquid now boils. In the vapour formed, the particles move in a rapid random motion. The movement is random due to the collision of the vapour particles with the air particles. An increase in temperature will cause the particles to move faster and an increase in pressure will cause gas particles to move closer together. 8

NES/Chemistry/IGCSE Brownian Motion Gases and Liquids are made up of particles that are always moving. This can been seen using a microscope in a \"smoke cell\". The smoke particles move because they are being constantly bombarded (hit) by gaseous air particles. The smoke particles can be seen moving in a zig-zag pattern. Diffusion Remember diffusion is when the substance moves by itself - not being mixed, or stirred by anyone; or the wind. Most liquids and gases are colourless, so it is hard to see them move. In experiments to show diffusion a coloured substance is usually placed in water, or air so we can watch it move and spread out. Diffusion in the different states of matter: 1. Solids 2. Liquids 3. Gases 1. Solids Solids diffuse so slowly that we say they do not diffuse. 9

NES/Chemistry/IGCSE 2. Liquids Liquids diffuse slowly. It can take many hours, or days for a liquid to diffuse.  Example 5: If a crystal of purple potassium manganate(VII) is placed in a beaker of water, the purple colour spreads throughout the whole beaker due to the continuous random motion of particles. After 24 hours the purple colour would be uniformly spread throughout the whole beaker. at the start later after 24 hours water uniform colour purple potassium manganate(VII) crystal 10

NES/Chemistry/IGCSE 3. Gases Gases diffuse the fastest out of the three states of matter. Different gases diffuse at different speeds - it depends on the mass, or density of the gas.  Example 6: Red brown volatile bromine liquid is placed in the bottom gas jar. This vaporises to fill the whole of the gas jar with orange/brown vapour. A gas jar of air is placed on top. When the glass lid is removed from between the jars the gases diffuse due to the constant rapid random motion of the particles. After about 20 minutes the colour would be uniformly orange/brown in both gas jars. gas jar containing air after 20 uniform colour throughout lid minutes the two gas jars due to the diffusion of gases orange brown bromine vapour Both of these investigations demonstrate that one substance can move through another due to the continuous random motion of particles (diffusion). The purple manganate(VII) ions spread between the water molecules. The molecules of orange brown bromine vapour mix with the gases in air. Gases diffuse faster than liquids. This is because the particles in a gas are moving in a continuous rapid random motion whereas particles in a liquid move much more slowly. A volatile liquid easily turns into a gas as it has a low boiling point (close to room temperature). 11

NES/Chemistry/IGCSE How the Mass of Gas affects Rate of Diffusion  The rate of diffusion is faster for low density/low Mr or Ar gases and slower for higher density/high Mr or Ar gases. (see Topic 4.1 for details on Ar and Mr)  The rate of diffusion will increase as the temperature increases. Long-Tube Experiment to Demonstrate Rate of Diffusion Two volatile liquids (concentrated ammonia and concentrated hydrochloric acid) are put at opposite ends of a long glass tube. The gases will diffuse in the tube and mix together to form a white solid precipitate of ammonium chloride. cotton wool soaked ammonium chloride (NH4Cl) cotton wool soaked in concentrated aqueous in concentrated ammonia hydrochloric acid Where the ammonia gas and hydrogen chloride gas meet, a white precipitate is produced. This is solid ammonium chloride (NH4Cl). To find which gas moves faster we need to calculate the relative molecular mass of both gases, using the mass numbers from the Periodic Table.  ammonia gas (NH3) = 14 + (3 x 1) = 17  hydrogen chloride gas (HCl) = 1 + 35.5 = 36.5 Particles with a lower relative molecular mass/lower density, diffuse faster than particles with a higher relative molecular mass/higher density, with the same energy (at the same temperature). As ammonia gas has a lower relative molecular mass/lower density compared with hydrogen chloride gas, so the ammonia diffuses faster than the hydrogen chloride gas. This means that the white ammonium chloride forms closer to the hydrochloric acid side as it is the slower diffusing gas. 12

NES/Chemistry/IGCSE Topic 2 - Experimental Techniques 2.1 Measurement  Name appropriate apparatus for the measurement of time, temperature, mass and volume, including burettes, pipettes and measuring cylinders 2.2.1 Criteria of purity  Demonstrate knowledge and understanding of paper chromatography  Interpret simple chromatograms  Identify substances and assess their purity from melting point and boiling point information  Understand the importance of purity in substances in everyday life, e.g. foodstuffs and drugs  Interpret simple chromatograms, including the use of Rf values  Outline how chromatography techniques can be applied to colourless substances by exposing chromatograms to substances called locating agents (Knowledge of specific locating agents is not required.) 2.2.2 Methods of purification  Describe and explain methods of purification by the use of a suitable solvent, filtration, crystallisation and distillation (including use of fractionating column). (Refer to the fractional distillation of petroleum in section 14.2 and products of fermentation in section 14.6.)  Suggest suitable purification techniques, given information about the substances involved 13

NES/Chemistry/IGCSE 2.1 - Measurement Time Apparatus Units Temperature stop clock hours, minutes, seconds thermometer Mass electronic balance oC measuring cylinder kg, g, milligram, *tonne Volume *burette (liquids) *pipette cm3, dm3 Volume gas syringe cm3, dm3 (gas) *A pipette measures fixed liquid volumes very accurately e.g. 25.00 cm3 *A burette measures liquid volumes between 0.00 – 50.00 cm3. It is very accurate having 2 decimal places. * A tonne is 1000 kg 14

NES/Chemistry/IGCSE Common Laboratory Apparatus beaker separating funnel conical flask crucible evaporating basin (graduated) pipette burette funnel gas jar gas syringe measuring cylinder pestle + mortar 15

NES/Chemistry/IGCSE 2.2.1 - Criteria of Purity Chromatography Chromatography is used to separate and identify a mixture of substances due to their different solubilities and their attraction to the chromatography paper. In chromatography, a piece of paper is placed in a beaker of solvent. The solvent travels up the paper, and carries the dissolved substances (solutes) with it. The substances will only move if they are soluble. Insoluble substances remain on the pencil line. The substance, which is the most soluble and with the least attraction to the paper, travels furthest up the paper. The distance moved up the paper by the solvent is called the solvent front. The distance moved up the paper by each solute in a mixture is called a 'dot'. Some substances, like amino acids and sugars (see Topic 14.8.3) are colourless and cannot been seen on the chromatograph. They must be sprayed with a locating agent, like ninhydrin, which adds colour and makes the 'dots' visible. We identify unknown mixtures of solutes by comparing them to known chromatograms, or by comparing Rf values. start of experiment glass results rod solvent front chromatography paper datum line/ xx x x base line suitable solvent 16

NES/Chemistry/IGCSE Chromatography Method  Draw a pencil line 2cm from the bottom of the chromatography paper called the datum line. (Pencil is used as it is insoluble in the solvent)  Place concentrated spots of different samples on the base line or datum line.  Suspend the paper in a suitable solvent, with the solvent level below the datum line.  When the solvent front reaches near the top of the paper remove the paper from the solvent and observe the spots if coloured. The paper is now referred to as a chromatogram.  If the spots are colourless then spray the chromatogram with a 'locating agent' to make the colourless spots become visible.  If the chromatogram shows one spot for a sample (sample A), then that sample is pure, but if there are a number of spots per sample (sample B, C, D), then the sample is impure.  A very soluble component in a sample travels high up the paper. Rf values The Rf value of a sample is calculated by:  measuring the distance travelled by the sample from the datum line  measuring the distance from the datum line to the solvent front Now substitute into the equation: Rf valueof a sample  distancetravelledby sample distancetravelledby solventfront 17

NES/Chemistry/IGCSE  Example 1: Identifying Proteins A and B solvent front 10 cm 4 cm 7 cm XX sample A sample B initial level of solvent Rf valueof a sampleA  distancetravelledby sampleA  4 cm  0.4 distancetravelledby solventfront 10 cm Rf valueof a sampleB  distancetravelledby sampleB  7 cm  0.7 distancetravelledby solventfront 10 cm Some Rf values for amino acids are: glutamic acid = 0.4 glycine = 0.5 alanine = 0.7 leucine = 0.9 So, sample A is glutamic acid and sample B is alanine. 18

NES/Chemistry/IGCSE Test for Purity using Melting Point and Boiling Point Pure substances melt and boil at a temperature that is fixed for a particular substance.  Example 2: Pure ice always melts at 0 oC and pure water boils at 100oC.  Example 3: Frozen salt-water melts at a temperature lower than 0oC and over a temperature range, and boils at a temperature above 100oC and over a temperature range A pure and an impure substance was heated from below its melting point to a temperature above its boiling point. Heating curve for a pure substance Heating curve for an impure substance boiling point boiling range temperature (oC) melting range temperature (oC) time (seconds) melting point time (seconds) A pure substance melts and boils at a temperature that is fixed for that substance. An impure substance melts over a temperature range and below that of the pure substance and boils over a temperature range above the boiling point of the pure substance. 19

NES/Chemistry/IGCSE 2.2.2 - Methods of Purification 1. Decanting Decanting is the process of removing a liquid from a solid, which has settled, or from an immiscible heavier liquid by pouring. 2. Filtration Filtration is used to separate a soluble solid from an insoluble solid, or separate an insoluble solid from a liquid. The insoluble material cannot pass through the tiny holes in the filter paper and so remains on the filter paper. It is insoluble solid (rock) now called the residue. filter paper The soluble material consists of very filter funnel small particles, which can pass through the filter paper with the solvent. The liquid containing the dissolved material, which passed through the filter paper, is called the filtrate. filtrate (salt solution) Solute is a soluble solid, which dissolves in a solvent. Solvent is a liquid, which dissolves a solute. Solution is solvent which has a solute dissolved in it. 20

NES/Chemistry/IGCSE 3. Evaporation Evaporation is used to separate a soluble salt from solution by boiling with a Bunsen burner. This method is not used if the salt crystals contain water of crystallisation The solution in the evaporating dish is heated until all the water has evaporated. A steam bath is used to provide a gentle even heating of the evaporating dish and prevents the salt from being lost from the dish by “spitting out”.  Example 4: Separating salt from rock salt (Filtration and Evaporation) Step 1 - Crushing: pestle The material is crushed into small pieces mortar to increase the surface area of the rock salt allowing the salt to dissolve faster. Step 2 - Dissolving: The salt dissolves as it is soluble, but the glass stirring rod rock being insoluble does not dissolve. water The mixture of rock salt and water are wire gauze heated and stirred to increase the rate of dissolving and the amount of salt tripod dissolved. 21

Step 3 - Filtration: NES/Chemistry/IGCSE filter paper Filtration separates the insoluble rock (residue) material (rock) from the soluble material (salt). The dissolved salt filter funnel passes through the filter paper because salt solution its particles are much smaller than the tiny holes in the filter paper, whereas the (filtrate) particles of rock are too big to fit through the holes. Step 4 - Evaporation: The rock remaining in the filter paper is salt solution called the residue. The filtered salt solution is called the steam bath filtrate. This separates the soluble material (salt) from the solvent (water). A steam bath is used to provide a gentle, even heating of the evaporating basin. This reduces the loss of salt from the evaporating basin by “spitting out”. 22

NES/Chemistry/IGCSE 4. Crystallisation Crystallisation is the process of forming crystals (hydrated salts) in a solution of a crystalline substance. Crystallisation is important when producing salt crystals, which contain water of crystallisation. If these salts were heated to dryness, then the salt would not contain the water of crystallisation. The product would not be crystals, but would be an anhydrous powder. Method of Crystallisation 1. Heat the solution on a steam bath until the crystallisation point is reached. 2. Cool the saturated solution to crystallise. 3. Any crystals formed are filtered and then washed with minimal amounts of cold distilled water to obtain pure crystals. 4. Dry the crystals between two wads of filter paper or in a desiccator. A desiccator is an airtight glass container, which contains a chemical, which absorbs water. Crystals can be dried in this apparatus without losing any water of crystallisation. A desiccator will dry the crystals, but an oven would dry and dehydrate the crystals. Finding the Point of Crystallisation Dip a glass rod into the solution being heated, and then remove the glass rod. If crystals form on the rod, then the solution is saturated and at its crystallisation point. A saturated solution is a solution, which contains the maximum amount of dissolved material at that particular temperature. Slow evaporation produces large crystals, and fast evaporation produces small crystals.  Example 5: Blue hydrated copper(II) sulfate crystals (CuSO4.5H2O) contain water of crystallisation. If they were heated to dryness then anhydrous CuSO4 (white) powder would be produced. So the hydrated salt has to be made by crystallisation. 23

NES/Chemistry/IGCSE 5. Sublimation Two solids can be separated if one of the solids sublimes. An example would be a mixture of ammonium chloride (NH4Cl) and sodium chloride (NaCl). Some ammonium compounds sublime. When the mixture is heated, the NH4Cl sublimes to form vapour which then sublimes on the cold surface of the funnel forming the solid NH4Cl. The sodium chloride solid should remain in the evaporating basin. Observations  white fumes  white solid forms on the inner surface of the funnel 6. Simple Distillation Simple distillation separates a solvent from a solution by boiling and condensing. thermometer Distillation is a 2-step process of boiling and then condensing. water out Liebig Residue will be left in the round bottom condenser flask and distillate will be produced in the conical flask. round Anti-bumping granules are added to the round bottom flask in order to make bottom flask water in the boiling smoother. HEAT receiver flask 24

NES/Chemistry/IGCSE 7. Fractional Distillation Fractional distillation is used to separate a mixture of miscible liquids with different boiling points. thermometer To separate a mixture of ethanol and water, heat the mixture up. At 78oC the water out ethanol boils, enters the fractionating Liebig column where it condenses and condenser evaporates repeatedly until it finally enters the condenser and condenses back to fractionating liquid and trickles into the receiving flask. column water in The thermometer at the top of the glass fractionating column tells us the boiling beads point of the vapour entering the electrical receiver flask condenser and thus allows us to identify heating mantle the distillate and possibly determine the purity of the distillate. The longer the fractionating column, the better the separation of liquids will be and the purer the distillate. The water enters from the bottom of the condenser because there is a better flow of water which completely surrounds the inner tube, keeping it cooler and therefore improving the amount of condensing.  Example 6: Separating an ethanol and water mixture. As the ethanol vapour is entering the condenser, the thermometer will read 78oC until all the ethanol has distilled over. Then the thermometer reading will steadily increase until 100oC when water will be entering the condenser. 25

NES/Chemistry/IGCSE Solubility The solubility of a solute in water at a given temperature is the number of grams of that solute which can be dissolved in 100g of water to produce a saturated solution at that temperature. Solubility has the units g of substance / 100g of water.  Example 7: Finding the solubility of sodium chloride at 30oC. Measure *exactly 100 cm3 of distilled water into a beaker. Place this beaker into a water bath at 30oC. Add a known mass (excess) of sodium chloride and stir. When no more solute dissolves filter the mixture. Evaporate the solution in a pre-weighed evaporating basin until dry. Measure the mass of the evaporating basin and solute. Subtract the mass of the empty evaporating basin to obtain the mass of dissolved solute, in g solute / 100g solvent. *If a different volume of water is used then the mass of substance dissolved has to be changed using the following formula: Solubility = 100 x massdissolved,g volumeof water, cm3 Solubility Curves These curves are obtained by finding the maximum amount of solute that can dissolve at various different temperatures. The solubility of most solids increases with temperature. The solubility of most gases decreases with temperature. The composition by mass of a saturated solution at any temperature can be read off a solubility curve. When a saturated solution is cooled, some of the solute crystallises out of solution. By using the solubility curve it is possible to determine the amount of solute which crystallises out of solution at different temperatures. 26

solubility (g substance per 100 g of water) NES/Chemistry/IGCSE  Example 8: Using the graph find the mass of solid that would precipitate if a saturated solution of potassium nitrate was cooled from 50oC to 10oC. sketch of solubility curves 100 90 80 potassium nitrate 70 60 50 40 sodium chloride 30 20 calcium ethanoate 10 0 10 20 30 40 50 60 70 80 90 100 temperature (oC) As the solution of potassium nitrate is saturated, it has the maximum amount of solid dissolved in water. So using the graph, we can see that at 50oC there would be 85g of potassium nitrate dissolved in a saturated solution. The maximum mass of potassium nitrate that can dissolve at 10oC is found from the graph as being only 19g. So as the solution cools it can hold less solid in solution. The mass of potassium nitrate that would precipitate is: 85g - 19g = 66g 27

NES/Chemistry/IGCSE Topic 3 - Atoms, Elements and Compounds 3.1 Atomic structure and the Periodic Table  State the relative charges and approximate relative masses of protons, neutrons and electrons  Define proton number (atomic number) as the number of protons in the nucleus of an atom  Define nucleon number (mass number) as the total number of protons and neutrons in the nucleus of an atom  Use proton number and the simple structure of atoms to explain the basis of the Periodic Table (see section 9), with special reference to the elements of proton number 1 to 20  Define isotopes as atoms of the same element which have the same proton number but a different nucleon number  State the two types of isotopes as being radioactive and non-radioactive  State one medical and one industrial use of radioactive isotopes  Understand that isotopes have the same properties because they have the same number of electrons in their outer shell  Describe the build-up of electrons in ‘shells’ and understand the significance of the noble gas electronic structures and of the outer shell electrons Note: a copy of the Periodic Table, as shown in the Appendix, will be available in Papers 2 and 4 (not Paper 6). 3.2.1 Bonding: the structure of matter  Describe the differences between elements, mixtures and compounds, and between metals and non-metals  Describe an alloy, such as brass, as a mixture of a metal with other elements 28

NES/Chemistry/IGCSE 3.2.2 Ions and ionic bonds  Describe the formation of ions by electron loss or gain  Describe the formation of ionic bonds between elements from Groups I and VII  Describe the formation of ionic bonds between metallic and non-metallic elements  Describe the lattice structure of ionic compounds as a regular arrangement of alternating positive and negative ions 3.2.3 Molecules and covalent bonds  Describe the formation of single covalent bonds in H2, Cl2, H2O, CH4, NH3 and HCl as the sharing of pairs of electrons leading to the noble gas configuration  Describe the differences in volatility, solubility and electrical conductivity between ionic and covalent compounds  Describe the electron arrangement in more complex covalent molecules such as N2, C2H4, CH3OH and CO2  Explain the differences in melting point and boiling point of ionic and covalent compounds in terms of attractive forces 3.2.4 Macromolecules  Describe the giant covalent structures of graphite and diamond  Relate their structures to their uses, e.g. graphite as a lubricant and a conductor, and diamond in cutting tools  Describe the macromolecular structure of silicon(IV) oxide (silicon dioxide)  Describe the similarity in properties between diamond and silicon(IV) oxide, related to their structures 3.2.5 Metallic bonding  Describe metallic bonding as a lattice of positive ions in a ‘sea of electrons’ and use this to describe the electrical conductivity and malleability of metals 29

NES/Chemistry/IGCSE Definitions An atom is the smallest part of an element that can exist as a stable entity An ion is a species with a positive or negative charge An element is a substance that contains only one type of atom A compound is a pure substance, which contains two or more different elements, chemically bonded together A molecule is the smallest part of an element or a compound, which can exist alone under ordinary conditions A mixture is two or more substances that can be separated by physical means Proton (atomic) number is equal to the number of protons in the nucleus of an atom Nucleon (mass) number is equal to the number of particles (protons and neutrons) in the nucleus of an atom Isotopes are atoms of the same element, with the same number of protons but different number of neutrons Radioactivity is a nuclear change where energy is released as radiation and a new nucleus is formed with a different proton number Elements are made up of the same type of atom. They cannot be broken down into simpler chemicals A macromolecule consists of a large number of atoms bonded together to form one molecule with a very high relative molecular mass (Mr) Allotropes are different crystalline forms of the same element. They have the same chemical properties, but different physical properties Metallic bonding is the strong non-directional electrostatic force of attraction between the lattice of positive ions and the mobile ‘sea’ of valence electrons in a 3- dimensional structure 30

NES/Chemistry/IGCSE 3.1 Atomic Structure and the Periodic Table All materials are made from atoms. Atoms are made from 3 types of particles:  protons (found in the nucleus)  neutrons (found in the nucleus)  electrons (found in shells) Relative Charge and Mass of the Particles in an Atom Particle Relative charge Relative mass Proton 1+ Neutron 0 1 Electron 1– 1 1 1837 In an atom, the number of the electrons equals the number of protons in the nucleus. Electrons fill the shells around the nucleus. The electrons start filling the nucleus inner electron shell closest to the nucleus first. Only once first shell holds up to 2 electrons this shell is full can electrons second shell holds up to 8 electrons then start filling up the next third shell holds up to 8 electrons shell, which is further away from the nucleus. Symbols of Elements All elements are denoted by a symbol, which consists of 1 or 2 letters. The first letter is always a capital/upper case, the second (if there is one) is a lower case.  Example 1: sodium = Na chlorine = Cl hydrogen = H 31

NES/Chemistry/IGCSE Proton Number and Nucleon Number Two numbers are associated with the elements in the Periodic Table. 23Na The larger top number is called the nucleon (mass) number (A) The smaller bottom number is called the proton (atomic) number (Z) 11 These two numbers allow us to work out the number of protons, neutrons and electrons in the atom. Proton Number This is how many protons are in the nucleus of an atom. It is also equal to the number of electrons in the shells of the atom. Atoms contain the same number of protons and electrons. Each positively charged proton cancels out each negatively charged electron, so there is no net charge - all atoms are neutral. Nucleon Number This is how many nucleons are in the nucleus of an atom. A nucleon is a particle in the nucleus - either proton, or neutron. Number of neutrons = nucleon number – proton number  Example 2: hydrogen and helium 1 H protons =1 1 =0 neutrons =1 electrons 4 He protons =2 2 =2 neutrons =2 electrons 32

NES/Chemistry/IGCSE Isotopes Isotopes of an element have the same chemical reactions, because isotopes have the same number of valence electrons and so lose, gain or share the same number of electrons to have a full outer shell, and it is these electrons which determine the chemical properties of an element, so therefore the same chemical reactions. The only difference is the density of each isotope.  Example 3: chlorine 35 Cl and 37 Cl 17 17  Example 4: carbon 12 C and 13 C and 14 C 6 6 6 Isotopes may be radioactive, or non-radioactive.  Example 5: isotopes of hydrogen: 1 H non-radioactive 1 non-radioactive radioactive 2 H - also called deuterium and can also have the symbol 2 D 1 1 3 H - also called tritium and can have the symbol 3 T 1 1 Due to radioactivity, a new atom is formed, this is not a chemical change, it is a nuclear change. Some of these radioactive isotopes occur naturally, but others are made in nuclear reactors. Medical Uses of Radioactive Isotopes  cobalt-60 ( 60 Co) is used in treating inaccessible cancerous cells 27  strontium-90 ( 90 Sr) is used in treating skin cancer 38  iodine-131 ( 15331I) is used to monitor the function of the thyroid gland  gamma rays can be used to sterilise surgical equipment 33

NES/Chemistry/IGCSE Industrial Uses of Radioactive Isotopes  strontium-90 ( 90 Sr) is used to monitor the thickness of paper or metal on a 38 production line to make sure it is not too thick or too thin  uranium-235 ( 235 U) is used to generate electricity in a nuclear power station 92 Electronic Configuration of Atoms This is simply the number of electrons in each electron shell of an atom starting from the electron shell closest to the nucleus.  Example 6: Element Proton Number Number of Electrons Electronic Configuration Hydrogen 1 1 1 Oxygen 8 8 2.6 Potassium 19 19 34

hydrogen helium protons = 1 protons = 2 electrons = 1 electrons = 2 neutrons = 0 neutrons = 2 1 2 NES/Chemistry/IGCSElithiumberylliumboroncarbonnitrogenoxygenfluorineneon protons = 3 protons = 4 protons = 5 protons = 6 protons = 7 protons = 8 protons = 9 protons = 10 Electronic Configurations for the First 20 Elementselectrons = 3electrons = 4electrons = 5electrons = 6electrons = 7electrons = 8electrons = 9electrons = 10 neutrons = 4 neutrons = 5 neutrons = 6 neutrons = 6 neutrons = 7 neutrons = 8 neutrons = 10 neutrons = 10 35 2.1 2.2 2.3 2.4 2.5 2.6 2.7 2.8 sodium magnesium aluminium silicon phosphorus sulfur chlorine argon protons = 11 protons = 12 protons = 13 protons = 14 protons = 15 protons = 16 protons = 17 protons = 18 electrons = 11 electrons = 12 electrons = 13 electrons = 14 electrons = 15 electrons = 16 electrons = 17 electrons = 18 neutrons = 12 neutrons = 12 neutrons = 14 neutrons = 14 neutrons = 16 neutrons = 16 neutrons = 18 neutrons = 22 2.8.1 2.8.2 2.8.3 2.8.4 2.8.5 2.8.6 2.8.7 2.8.8 potassium calcium protons = 19 protons = 20 electrons = 19 electrons = 20 neutrons = 20 neutrons = 20

NES/Chemistry/IGCSE 3.2.1 Bonding: The Structure of Matter Elements are made of atoms of the same type. For example, oxygen is made up of oxygen atoms. Properties of Elements A physical property is something about the appearance or state of a material. The most common properties are:  melting and boiling points (either high or low)  electrical conductivity (either good or poor)  density (either high or low) A chemical property is something about the chemistry of a material. The most common properties are:  form positive, or negative ions  form basic, or acidic oxides The Periodic Table arranges elements in order of ascending proton number and can be split into metals and non-metals. There is a change from metallic character to non-metal character going from left to right across the periodic table. I II Group number (number of valence electrons) III IV V VI VII 0 1 H He 2 Li Ne Period number 3 Na Be B CNO F Ar 4K Mg Al Si P S Cl Kr (number of shells) 5 Rb Ca Ga Ge As Se Br Xe 6 Cs Sr In Sn Sb Te I Rn 7 Fr Ba Tl Pb Bi Po At Ra alkali metals transition metals halogens noble gases 36

NES/Chemistry/IGCSE Physical Properties Metals Non-Metals usually high density (except group I) usually low density usually silvery/shiny and can be polished dull malleable (can be hammered into shapes usually soft or brittle when solid without breaking) usually soft or brittle when solid Ductile (can be drawn into thin wires) usually have high melting and boiling points usually have low melting and boiling points good conductor of heat and electricity poor conductors of heat and electricity Chemical Properties Metals Non-Metals usually form basic oxides usually form acidic oxides form positive ions usually form negative ions react with non-metals forming ionic bonds react with metals forming ionic bonds react with non-metals forming covalent bonds Compounds are made up of different elements. They have their elements chemically bonded together and can only be separated by chemical reactions. Mixtures do not have their elements bonded together, so they can be separated by physical methods, such as filtration, decanting, distillation, evaporation, crystallisation and chromatography (see Topic 2). An alloy is a mixture of two or more metals which have been made molten and then mixed together. It is then cooled and allowed to solidify. Alloys are used because they have the properties of each of the different metals mixed together. They are always stronger than the metals they are made from. Steel is an alloy made from iron and carbon, so it only contains one metal (carbon is a non-metal) but it is still an alloy. 37

NES/Chemistry/IGCSE 3.2.2 Ions and Ionic Bonds Ions Ions can be formed two different ways: 1. Cations are positive ions that are formed when metal atoms give away their valence (outer shell) electrons. 2. Anions are negative ions that are formed when a non-metal atom takes in electrons to fill their valence shell. They do this to obtain the noble gas electronic configuration by having a full, or empty, valence (outer) shell. Metal Ions Metal ions have the same electronic configuration as the noble gas, which is at the end of the period above where the metal was.  Example 7: sodium Sodium is in period 3, so it forms a cation with the noble gas electronic structure of neon, which is in period 2. electron X XX + XX X X Na Na + e– XX XX XX XX X X XX XX nucleus sodium atom sodium ion electronic configuration 2.8.1 2.8 Valency is the number of electrons that have been given away, or taken in by an atom. The valency for metals is the same number as the group number in the Periodic table. 38

Element Number of Valence Valency NES/Chemistry/IGCSE Electrons sodium 1 1 Charge on the Ion magnesium 2 2 1+ aluminium 3 3 2+ 3+ Non-Metal Ions When non-metals atoms form ions they have the noble gas electronic configuration of a noble gas in the same period.  Example 8: chlorine Chlorine in period 3 forms an anion with the same electronic configuration as argon in period 3. electron – Cl + e– X Cl nucleus chlorine atom chloride ion 2.8.7 2.8.8 The valency for non-metals = 8 - group number Element Number of Valency Name of Ion Charge on Valence the Ion chlorine Electrons 1 chloride oxygen 7 2 oxide 1– nitrogen 6 3 nitride 2– 5 3– 39

NES/Chemistry/IGCSE Electronic configuration of ions Metal or Electronic Electronic Loss, or Gain Element Non-Metal Configuration of Configuration of Ion of Electrons Atom potassium metal 2.8.8 loss chlorine non-metal 2.8.7 2.8.8 gain oxygen non-metal 2.8 gain 2.6 2.8 loss aluminium metal 2.8.3 2.8 gain nitrogen non-metal 2.5 Ions Period Electronic Noble Gas with the Same Electronic Configuration Configuration Li+, Be2+ 2 He Na+, Mg2+, Al3+ 3 2 Ne K+, Ca2+, Ga3+ 4 2.8 Ar 2 2.8.8 Ne F-, O2-, N3- 2.8 40

NES/Chemistry/IGCSE Ionic Bonds An ionic bond is a strong non-directional electrostatic force of attraction between cations and anions formed due to the transfer of electrons. An ionic bond is formed by a metal atom losing its valence electrons to a non-metal atom or atoms. The metal ion formed has a positive charge and the non-metal ion formed has a negative charge. Both ions formed have a noble gas electronic configuration / full outer shell. In naming ionic compounds, always put the metal ion first.  Example 9: sodium chloride: Formation of sodium chloride showing full electron structures X XX – XX X X+ Na Cl X XX XNa Cl XX X XXX XX 2.8.1 2.8.7 2.8 or [2.8]+ 2.8.8 or [2.8.8]– XX XX Formation of sodium chloride showing valence electrons only – X Na Cl [Na]+ Cl 2.8.1 2.8.7 2.8 or [2.8]+ 2.8.8 or [2.8.8]– Key: X = electron from Na 41 = electrons from Cl

NES/Chemistry/IGCSE Why do Cl– ions have a larger ionic radius than Na+ ions? The Cl– ions have a larger ionic radius than the Na+ ions as chloride ions have three complete shells of electrons, whereas sodium ions have only two complete electron shells.  Example 10: magnesium chloride: Element Symbol Number of Protons Neutrons Electrons 12 magnesium Mg 17 12 12 chlorine Cl 18/20 17 X Cl – Mg X2.8.7 X Cl 2.8.2 [Mg]2+ 2 Cl 2.8 or [2.8]2+ 2.8.8 or [2.8.8]– Formula = MgCl2 One magnesium atom reacts with two chlorine atoms. Magnesium has two valence electrons, which it loses. One chlorine atom can only accept one of the electrons, so a second chlorine atom is required. For every one magnesium atom we need two chlorine atoms. 42

NES/Chemistry/IGCSE  Example 11: magnesium oxide: Element Symbol Protons Number of 12 Neutrons Electrons magnesium Mg 8 oxygen O 12 12 88 X 2– Mg XO X 2.8.2 [Mg]2+ O x 2.6 2.8 or [2.8]2+ 2.8 or [2.8]2– Formula = MgO One magnesium atom reacts with one oxygen atom. Magnesium has two valence electrons, and one oxygen atom requires two electrons to have the noble gas electronic configuration. So one atom of magnesium reacts with one atom of oxygen. 43

NES/Chemistry/IGCSE  Example 12: aluminium fluoride: Element Symbol Protons Number of 13 Neutrons Electrons aluminium Al 9 fluorine F 14 13 10 9 X F – F Al xF F 2.7 X X 2.8.3 [Al]3+ 3 Formula = AlF3 2.8 or [2.8]3+ 2.8 or [2.8]– 44

NES/Chemistry/IGCSE  Example 13: sodium oxide: Element Symbol Protons Number of 11 Neutrons Electrons sodium Na 8 oxygen O 12 11 88 X 2– Na 2[Na]+ X O x O X Na 2.8.1 2.6 2.8 or [2.8]+ 2.8 or [2.8]2– Formula = Na2O 45

NES/Chemistry/IGCSE Ionic Compounds When ions form due to the transfer of electrons, they do not exist in pairs, but form giant 3-dimensional crystal lattice structures with millions of ions arranged in regular rows with opposite charges next to each other. The lattice structure of ionic compounds is a regular arrangement of alternating positive and negative ions in a 3-dimensional structure. Structure In sodium chloride crystals, each sodium ion is surrounded by 6 chloride ions. Each chloride ion is surrounded by 6 sodium ions forming a giant ionic crystal. . Cl– surrounded by 6 Na+ ions Na+ surrounded by 6 Cl– ions The alternating positive and negative ions in the crystal lattice of ionic compounds are held together by strong non-directional electrostatic forces of attraction. 46

NES/Chemistry/IGCSE This diagram below is a 2-dimensional representation of the ionic compound sodium chloride. Na+ Cl Na+ Cl Na+ Cl Na+ Cl Cl Na+ Cl Na+ Cl Na+ Cl Na+ Na+ Cl Na+ Cl Na+ Cl Na+ Cl Cl Na+ Cl Na+ Cl Na+ Cl Na+ 47

NES/Chemistry/IGCSE Physical Properties Property Reason Conduct electricity when molten or dissolved in There are mobile ions solution High melting points and There are strong non-directional electrostatic forces of boiling points attraction holding the ions together in the giant lattice Soluble in water Water molecules are able to bond with the positive and Brittle negative ions, which breaks up the lattice and keeps the ions apart If a force is applied to the bottom rows of ions so they move by 1 ion length, then we see that positive ions are lined up with positive ions, and negative ions are lined up with negative ions. We get repulsion between layers and the crystal breaks Brittle describes a solid material that easily breaks into small pieces under force and does not bend. Na+ Cl Na+ Cl Na+ Cl Na+ Cl Cl Na+ Cl Na+ Cl Na+ Cl Na+ force Na+ Cl Na+ Cl Na+ Cl repulsion Na+ Cl Cl Na+ Cl Na+ Cl Na+ Cl Na+ 48

NES/Chemistry/IGCSE 3.2.3 Molecules and Covalent Bonds Covalent Bonds A covalent bond is a directional bond formed between two non-metal atoms by sharing one or more pairs of electrons in the overlap in order to have a full outer shell. A full valence shell is a stable arrangement of electrons, and this is why the noble gases do not react and have chemical stability.  Example 14: hydrogen gas: Hydrogen gas does not consist of individual atoms (monatomic like the noble gases), but two atoms of hydrogen which are bonded by a covalent bond to form a hydrogen molecule H2 (diatomic) This pairing allows the atoms to share electrons (one from each atom) so that each hydrogen atom has two electrons in its valence shell. The first shell is full with two electrons, and a full shell is a more stable arrangement of electrons. valence electrons covalent bond H H HH + The H2 molecule can be represented as this stick diagram H – H. The dash in between the two H‟s represents two electrons (one from each atom) Group VII (halogens) have seven valence electrons. Two chlorine atoms share a pair of electrons (one from each atom) in the overlap so both atoms have a full outer shell. 49

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