Chemistry IGCSE Textbook

NES/Chemistry/IGCSE Strength of Oxidising Agents Very reactive non-metals are strong oxidising agents, because they readily gain electrons to obtain a full valence shell. The more reactive a non-metal is, the stronger an oxidising agent it will be. Fluorine is the most reactive halogen. F2 can oxidise any of the halide anions to produce fluoride anions and a diatomic halogen molecule. Chlorine can oxidise both bromide and iodide anions because chlorine is more reactive that bromine or iodine. Bromine can oxidise iodide anions to iodine molecules.  Example 7: halogen displacement reaction Symbol equation Cl2 + 2KI(aq)  2KCl(aq) + I2 Ionic equation Cl2 + 2I–  2Cl– + I2 Ionic half-equations Cl2 + 2e–  2Cl– reduction 2I–  I2 + 2e– oxidation 150

NES/Chemistry/IGCSE Topic 8 - Acids, Bases and Salts 8.1 The characteristic properties of acids and bases • Describe the characteristic properties of acids as reactions with metals, bases, carbonates and effect on litmus and methyl orange • Describe the characteristic properties of bases as reactions with acids and with ammonium salts and effect on litmus and methyl orange • Describe neutrality and relative acidity and alkalinity in terms of pH measured using Universal Indicator paper • Describe and explain the importance of controlling acidity in soil • Define acids and bases in terms of proton transfer, limited to aqueous solutions • Describe the meaning of weak and strong acids and bases 8.2 Types of oxides • Classify oxides as either acidic or basic, related to metallic and non-metallic character • Further classify other oxides as neutral or amphoteric 8.3 Preparation of salts • Demonstrate knowledge and understanding of preparation, separation and purification of salts as examples of some of the techniques specified in section 2.2.2 and the reactions specified in section 8.1 • Demonstrating knowledge and understanding of the preparation of insoluble salts by precipitation • Suggest a method of making a given salt from a suitable starting material, given appropriate information 151

NES/Chemistry/IGCSE 8.4 Identification of ions and gases • Describe the following tests to identify: aqueous cations: aluminium ammonium calcium chromium(III) copper(II) iron(II) and iron(III) zinc (using aqueous sodium hydroxide and aqueous ammonia as appropriate) cations: use of the flame test to identify lithium, sodium, potassium and copper(II) anions: carbonate (by reaction with dilute acid and then limewater) chloride, bromide and iodide (by reaction under acidic conditions with aqueous silver nitrate) nitrate (by reduction with aluminium) sulfate (by reaction under acidic conditions with aqueous barium ions) sulfite (by reaction with dilute acids and then aqueous potassium manganate(VII) ) gases: ammonia (using damp red litmus paper) carbon dioxide (using limewater) chlorine (using damp litmus paper) hydrogen (using lighted splint) oxygen (using a glowing splint) sulfur dioxide (using aqueous potassium manganate(VII) ) 152

NES/Chemistry/IGCSE 8.1 The Characteristic Properties of Acids and Bases An acid is a proton donor A base is a proton acceptor Acids When an acid reacts, it gives away at least one of its hydrogen ions, H+. So all acids must contain hydrogen (H). Strong Acid Weak Acid Name Formula Name Formula sulfuric acid H2SO4 ethanoic acid CH3COOH hydrochloric acid HCl propanoic acid C2H5COOH nitric acid HNO3 butanoic acid C3H7COOH Characteristic Properties of Acids  pH = 0 - 6  Turn blue litmus red  Taste sour  Corrosive  All acids are solutions  Strong acids completely ionise when dissolved in water  Weak acids only partly ionise when dissolved in water 153

NES/Chemistry/IGCSE Bases When a base reacts, it takes in hydrogen ions, H+ to form water and a salt.  A base is a metal oxide - base is short for basic oxide  An alkali is a soluble base, a metal hydroxide. Strong Alkali Weak Alkali Name Formula Name Formula sodium hydroxide NaOH aqueous ammonia NH4OH potassium hydroxide KOH Characteristic Properties of Bases  pH = 8 - 14  Turn red litmus blue  Feel soapy  Corrosive  All alkalis are solutions  Strong alkalis completely ionise when dissolved in water  Weak alkalis only partly ionise when dissolved in water Acid Reactions Acid + Metal  Salt + Hydrogen Acid + Base  Salt + Water Acid + Alkali  Salt + Water Acid + Metal Carbonate  Salt + Water + Carbon Dioxide 154

NES/Chemistry/IGCSE 1. Acid - Metal Reactions A metal will react with an acid provided the metal is more reactive than hydrogen. Because the metal is more reactive than hydrogen the metal atoms lose their valence electrons to form their positive ions and the hydrogen ions gain electrons and so are reduced to hydrogen atoms, so the hydrogen ions are displaced forming hydrogen gas. Copper, silver, gold, or platinum will not react with dilute acid as they are below hydrogen in the reactivity series.  Example 1: magnesium and sulphuric acid balanced equation Mg(s) + H2SO4(aq)  MgSO4(aq) + H2(g) ionic equation Mg(s) + 2H+(aq)  Mg2+(aq) + H2(g) Never add potassium, or sodium to acid as it is explosive. Observations:  metal solid decreases in size  effervescence and bubbles of a colourless gas are evolved  temperature of the reaction mixture increases  In the case of magnesium an extra observation is that even though magnesium has a higher density than acid, the Mg floats at the surface. This is because the numerous bubbles of H2 gas forming on the surface keep the Mg afloat. 2. Acid - Base Reactions A base is an insoluble alkali. This reaction is slow so we heat up the acid/base mixture to increase the rate of reaction.  Example 2: magnesium oxide and sulphuric acid balanced equation MgO(s) + H2SO4(aq)  MgSO4(aq) + H2O(l) ionic equation MgO(s) + 2H+(aq)  Mg2+(aq) + H2O(l) Observations:  solid decreases in size 155

NES/Chemistry/IGCSE 3. Acid - Alkali Reactions An alkali is a soluble base. Bases dissolve in water producing OH– ions. Base + Water  Alkali  Example 3: sodium hydroxide and sulphuric acid balanced equation 2NaOH(aq) + H2SO4(aq) Na2SO4(aq) + 2H2O(l) ionic equation OH–(aq) + H+(aq)  H2O(l) Observations:  temperature of the acid/alkali mixture increases 4. Acid - Metal Carbonate Reactions Metal carbonates react very rapidly with acids. A metal carbonate has to be added slowly and in small amounts to an acid otherwise the froth formed will overflow the beaker.  Example 4: magnesium carbonate and sulphuric acid balanced equation MgCO3(s) + H2SO4(aq)  MgSO4(aq) + H2O(l) + CO2(g) ionic equation MgCO3(s) + 2H+(aq)  Mg2+(aq) + H2O(l) + CO2(g) Observations:  solid rapidly decreases in size  effervescence and bubbles of a colourless gas are evolved 156

NES/Chemistry/IGCSE Test for acids 1 Dip a pH meter into the suspected acid sample. If the solution is acidic then the pH meter gives a reading less than pH 7. 2 Add Universal or blue litmus indicator paper or solution to the suspected sample and note the colour change. Dry Universal indicator paper changes from yellow to red/orange Blue litmus paper changes from blue to red 3 Add a small spatula of solid sodium carbonate to the suspected acid sample. If an acid is present then:  the solid decreases in size  there is effervescence and bubbles of a colourless gas are evolved which turns limewater from colourless to milky. 4 Add a piece of magnesium ribbon to the suspected acid sample. If an acid is present then:  the metal decreases in size  there is effervescence and bubbles of a colourless gas are evolved which produces a ‘squeaky pop’ with a lighted splint. 157

NES/Chemistry/IGCSE Base Reactions Base + Acid  Salt + Water Base + Ammonium Salt  Salt + Water + Ammonia 1. Base - Acid Reactions A base is an insoluble alkali. This reaction is slow so we heat up the acid/base mixture to increase the rate of reaction.  Example 5: calcium oxide and sulphuric acid balanced equation CaO(s) + H2SO4(aq)  CaSO4(aq) + H2O(l) ionic equation CaO(s) + 2H+(aq)  Ca2+(aq) + H2O(l) Observations:  solid decreases in size This is the same reaction as the Acid - Base reaction above. 2. Base - Ammonium Salt Reactions The mixture has to be warmed gently in order to produce ammonia gas. This reaction is used to test for the ammonium cation (NH4+) which is part of the analysis topic.  Example 6: sodium hydroxide and ammonium chloride balanced equation NaOH(aq) + NH4Cl(aq)  NaCl(aq) + H2O(l) + NH3(g) ionic equation OH–(aq) + NH4+(aq)  H2O(l) + NH3(g) Observations:  effervescence and bubbles of a colourless gas are evolved 158

NES/Chemistry/IGCSE Concentration of Acids / Alkalis Concentration of an acid is the number of moles of acid molecules per unit volume. The higher the number of moles of acid molecules per unit volume, the higher the concentration. A 1 mol/dm3 solution of HCl(aq) contains 1 mole of HCl in 1dm3 (see Topic 4.2). Concentration of an alkali is the number of moles of alkali molecules per unit volume. The higher the number of moles of alkali molecules per unit volume, the higher the concentration. A 1 mol/dm3 solution of NaOH(aq) contains 1 mole of NaOH in 1dm3. Strength of Acids Acidity is caused by the presence of H+ ions (hydrogen ions) in a solution. Strength is a measure of the degree of ionisation of the acid molecules. The stronger the acid, the greater the degree of ionisation, producing a higher concentration of H+ ions and the lower the pH number. Strong Acids Strong acids completely ionise in water forming H+ ions. Hydrochloric acid is formed by dissolving hydrogen chloride gas in water. All the HCl molecules fully ionise forming the ions H+(aq) and Cl–(aq). HCl(g) + H2O(l)  H+(aq) + Cl–(aq) Weak Acids Weak acids only partially ionise in water forming H+ ions. A weak acid forms fewer H+(aq) ions in comparison to a strong acid of the same concentration. Ethanoic acid is an example of a weak acid. This is an example of a reversible process. CH3COOH ⇌ H+ + CH3COO– 159

NES/Chemistry/IGCSE Concentration and pH A strong acid and a weak acid of the same concentration can have different concentrations of H+ ions and therefore a different pH. This is because a strong acid fully ionised in water producing more H+ ions per unit volume compared to a weak acid, which only partially ionises in water producing fewer H+ ions per unit volume. This means a strong acid is a better electrical conductor than a weak acid due to a higher concentration of ions. Strong and Weak Alkalis Alkalinity is caused by the presence of OH– ions (hydroxide ions) in a solution. When sodium hydroxide is dissolved in water all the NaOH molecules ionise forming the ions OH–(aq) and Na+(aq), which will have a high pH number (12-14). NaOH(s) + H2O(l)  Na+(aq) + OH–(aq) When ammonia dissolves, it only partly ionises, forming less OH- ions. The pH will be lower than that of a strong alkali, from 8-11. NH3(g) + H2O(l) ⇌ NH4+(aq) + OH–(aq) Monoprotic and Diprotic Acids Monoprotic acids (eg HCl) produce 1 hydrogen ion per acid molecule, whereas diprotic acids (eg H2SO4) produce two hydrogen ions per acid molecule. HCl(g) + H2O(l)  H+(aq) + Cl–(aq) H2SO4(aq) + H2O(l)  2H+(aq) + SO42–(aq) 160

NES/Chemistry/IGCSE pH Scale pH is a measure of the concentration of H+ ions per dm3. This is a scale, which indicates the strength of an acid, or alkali and ranges from 0 to 14. Acidic Neutral Alkaline neutral strong acid weak acid weak alkali strong alkali 7 0123456 8 9 10 11 12 13 14 Indicators Most indicators simply show us whether a substance is acidic or alkali, but not the strength. Indicator Colour in acidic solution Colour in alkaline solution phenolphthalein colourless pink methyl orange red yellow methyl red red yellow red litmus red blue blue litmus red blue Universal Indicator Universal Indicator paper or solution is the only indicator which can be used to give an indication of strength of an acid or alkali. The indicator paper when dry is yellow. If it is dipped in water the colour will be green. Universal indicator cannot be used for titration experiments. red orange yellow green blue purple 012 8 9 10 11 12 13 14 strong acid 345 6 7 strong alkali weak alkali weak acid neutral 161

NES/Chemistry/IGCSE 8.2 Types of Oxides Metal oxides and non-metal oxides are made by burning the elements in oxygen. Metal oxides are solids so they do not change the colour of dry universal indicator paper. Damp universal indicator paper, or solution is used so the material can dissolve in water. Metal Oxides Metal oxides are usually basic, which means that they will react with acids. Some metal oxides are amphoteric, which means that they will react with both acids and bases. Amphoteric oxides are found near the metal / non-metal border on the Periodic Table. Non-Metal Oxides Non-metal oxides are usually acidic, which means that they will react with bases. Some non-metal oxides are neutral, which means that they will not react with acids, or bases.  Examples: Acidic Oxides Neutral Oxides Amphoteric Oxides sulfur dioxide water aluminium oxide Basic Oxides carbon dioxide lead(II) oxide iron(III) oxide carbon monoxide zinc oxide copper(II) oxide calcium oxide magnesium oxide 162

NES/Chemistry/IGCSE Reactions of Amphoteric Oxides Amphoteric oxides can react with acids and bases.  Example 1: Aluminium oxide reacting with an acid Al2O3(s) + 6HCl(aq)  2AlCl3(aq) + 3H2O(l) In this reaction, Al2O3 is reacting as a basic oxide.  Example 2: Aluminium oxide reacting with an alkali Al2O3(s) + 2NaOH(aq) + 3H2O(l)  2NaAl(OH)4(aq) (sodium aluminate) In this reaction, Al2O3 is reacting as an acidic oxide. 163

NES/Chemistry/IGCSE 8.3 Preparation of Salts A salt is a substance formed when all the replaceable hydrogen ions of an acid are completely replaced by metal ions, or the ammonium ion (NH4+) A precipitate is an insoluble salt formed when two salt solutions are mixed Salts can be either soluble, or insoluble. Soluble Compounds Insoluble Compounds all sodium, potassium and ammonium salts except silver and lead except barium, lead and calcium all nitrates all halides (Cl–, Br–, I–) all other carbonates most other metal oxides and hydroxides all sulfates sodium carbonate potassium carbonate ammonium carbonate most oxides and hydroxides of Group I + II  Example 1: sodium carbonate is soluble  Example 2: lead chloride is insoluble  Example 3: barium sulphate is insoluble  Example 4: calcium bromide is soluble Making Soluble Salts Titration - using an acid (solution) and an alkali (solution) Method 1: acid + alkali  soluble salt + water Excess solid - using a metal, or a base with an acid Method 2: acid + metal  soluble salt + hydrogen Method 3: acid + base  soluble salt + water Method 4: acid + metal carbonate  soluble salt + water + carbon dioxide 164

NES/Chemistry/IGCSE Naming Salts The name of the salt will depend on the metal/metal compound and the acid used. The metal remains unchanged and is the first part of the name. The acid used gives the second part of the salt's name: Acid used Salt made sulfuric acid sulfate nitrate nitric acid chloride hydrochloric acid ethanoate ethanoic acid  Example 5: Using calcium oxide and nitric acid will make the salt calcium nitrate, as well as water. Making Insoluble Salts Precipitation - using 2 soluble salts Method 5: soluble salt + soluble salt  insoluble salt + soluble salt 165

NES/Chemistry/IGCSE Method 1 - Titration Titration is used to make soluble salts. The reaction is done the first time with an indicator (but not universal indicator) to get the volumes required of acid and alkali for exact neutralisation. Then the experiment is repeated using these volumes, but without the indicator, which would be an impurity in the salt made.  Example 6: Preparation of the soluble salt sodium chloride by titration using hydrochloric acid and sodium hydroxide. pipette clamp  Pipette 25.0 cm3 of HCl(aq) into a HCl(aq) conical flask burette containing  Add 3 drops of phenolphthalein NaOH(aq) indicator  Add NaOH(aq) drop-wise from the burette until the colour just changes from colourless to pink  Record the volume of NaOH(aq) added and repeat the experiment using the same volumes but without the indicator salt salt  Evaporate the salt solution on a solution crystals steam bath until the point of crystallisation steam bath  Cool to crystallise then filter to obtain the crystals HEAT  Wash the crystals with minimal amounts of cold distilled water  Dry the crystals in a desiccator 166

NES/Chemistry/IGCSE Method 2 - Excess Solid (metal) Excess solid method is used to make soluble salts. The solid is added to the acid until no more will react. This excess solid has to be removed by filtration. The filtrate is used to make the salt by evaporation.  Example 7: Preparation of the soluble salt magnesium sulphate by excess solid method using sulphuric acid and magnesium. H2(g)  Use a measuring cylinder to measure out about 25 cm3 of sulfuric acid into a beaker magnesium metal  Add magnesium ribbon to the sulfuric acid  Add the magnesium ribbon until the metal stops decreasing in size, excess magnesium remains and there is no more effervescence  Filter to remove the excess magnesium ribbon excess magnesium ribbon magnesium sulfate solution salt  Evaporate the filtrate on a steam bath until the solution point of crystallisation steam bath  Cool to crystallise  Filter to remove any crystals formed HEAT  Wash the crystals with minimal amounts of cold distilled water.  Dry the crystals in a desiccator 167

NES/Chemistry/IGCSE Method 3 - Excess Solid (insoluble metal oxide) Excess solid method is used to make soluble salts. The solid is added to the acid until no more will react. This reaction is slow and has to be heated. The excess solid has to be removed by filtration. The filtrate is used to make the salt by evaporation.  Example 8: Preparation of the soluble salt copper (II) sulphate by excess solid method using sulphuric acid and copper (II) oxide. glass  Use a measuring cylinder to measure out about 25 cm3 stirring of sulfuric acid into a beaker and heat rod H2SO4(aq)  Add black copper(II) oxide one spatula at a time to the CuO(s)  acid and stir Add CuO until the black solid stops decreasing in size and excess CuO remains HEAT excess  Filter to remove the excess CuO CuO(s) CuSO4(aq) salt  Evaporate the filtrate on a steam bath until the point of solution crystallisation steam  Cool to crystallise bath  Filter to remove any crystals formed  Wash the crystals with minimal amounts of cold HEAT distilled water  Dry the crystals in a desiccator 168

NES/Chemistry/IGCSE Method 4 - Excess Solid (insoluble metal carbonate) Excess solid method is used to make soluble salts. The solid is added to the acid until no more will react. This excess solid has to be removed by filtration. The filtrate is used to make the salt by evaporation.  Example 9: Preparation of the soluble salt copper (II) sulphate by excess solid method using sulphuric acid and copper (II) carbonate. glass  Use a measuring cylinder to measure out about 25 stirring cm3 of sulfuric acid into a beaker. rod  Add green copper(II) carbonate one spatula at a time H2SO4(aq)  to the acid. Stir the mixture until the effervescence stops and CuCO3(S) excess CuCO3 remains  Filter to remove the excess CuCO3 excess CuCO3(s) CuSO4(aq) salt  Evaporate the filtrate on a steam bath until the point solution of crystallisation steam  Cool to crystallise bath  Filter to remove any crystals formed  Wash the crystals with minimal amounts of cold HEAT distilled water.  Dry the crystals in a desiccator 169

NES/Chemistry/IGCSE Method 5 - Precipitation Precipitation is used to make an insoluble salt. No acid is used, just soluble salts.  Example 10: Preparation of the insoluble salt lead (II) iodide using soluble potassium iodide and lead (II) nitrate. potassium distilled  Pour a slight excess of water into iodide solid water each beaker containing the soluble solids lead(II)nitrate and potassium lead(II) iodide iodide. precipitate  Stir to dissolve the soluble solid salts to obtain two solutions  Mix the two solutions and stir to obtain the precipitate lead(II) iodide solid lead(II) nitrate distilled water wash  Filter the mixture to separate the precipitate from the solution  Wash the precipitate with distilled water  Dry the precipitate in a desiccator 170

NES/Chemistry/IGCSE Summary of Salt Preparation Method Type of Reagents Heating Soluble or reaction acid + soluble base (alkali) required insoluble salt 1.Titration neutralisation acid + metal  soluble  soluble 2.Excess redox acid + insoluble base solid acid + insoluble metal  soluble 3.Excess neutralisation carbonate  soluble solid two soluble salt solution neutralisation  insoluble 4.Excess double solid decomposition 5.Precipitation 171

NES/Chemistry/IGCSE 8.4 Identification of Ions and Gases Test for Gases Gas Test Result ammonia (NH3) Damp red litmus The paper changes from red carbon dioxide Bubble the gas through to blue (CO2) limewater The limewater will turn from colourless to white chlorine (Cl2) Damp red, or blue litmus paper precipitate Paper will be bleached (turn hydrogen gas (H2) Lighted splint white) oxygen (O2) Glowing splint ‘squeaky pop’ relight sulfur dioxide (SO2) Acidified potassium manganate (VII) paper Paper changes from purple to colourless. Note - Ammonia is the only basic gas. 172

NES/Chemistry/IGCSE Tests for Aqueous Cations Cation Effect of Aqueous Sodium Effect of Aqueous Ammonia aluminium (Al3+) Hydroxide white precipitate forms, which is zinc (Zn2+) white precipitate forms, which is insoluble in excess calcium (Ca2+) soluble in excess giving a chromium (Cr3+) colourless solution white precipitate forms, which is copper (Cu2+) white precipitate forms, which is soluble in excess giving a soluble in excess giving a colourless solution iron(II) (Fe2+) colourless solution no precipitate or very slight white iron(III) (Fe3+) precipitate. If a slight precipitate ammonium (NH4+) white precipitate forms, which is forms it is insoluble in excess insoluble in excess grey-green precipitate forms, which is insoluble in excess green precipitate forms, which is light blue precipitate forms, soluble in excess which is soluble in excess giving a dark blue solution light blue precipitate forms, green precipitate forms, which is which is insoluble in excess insoluble in excess. On standing for 5 minutes the surface turns green precipitate forms, which is red-brown as Fe3+ is formed insoluble in excess. On standing red-brown precipitate forms, for 5 minutes the surface turns which is insoluble in excess red-brown red-brown precipitate forms, which is insoluble in excess ammonia gas produced on warming which turns damp red litmus blue Metal Ion Flame Tests Metal Ion Flame Colour Lithium Red Sodium Yellow Potassium Lilac Copper (II) Blue-green 173

NES/Chemistry/IGCSE Tests for Anions Anion Test Result Bubble the gas evolved through carbonate (CO32–) add dilute nitric acid (gas limewater. The limewater [solid, or solution] produced) changes from colourless to white precipitate (carbon chloride (Cl–) acidify with dilute nitric acid, dioxide) [in solution] then add aqueous silver nitrate White precipitate forms bromide (Br–) (silver(I) chloride) [in solution] add aqueous sodium hydroxide then aluminium foil; warm Cream precipitate forms iodide (I–) carefully (gas produced) (silver(I) bromide) [in solution] acidify with dilute nitric acid, nitrate (NO3–) then add aqueous barium nitrate Yellow precipitate forms [in solution] or aqueous barium chloride (silver(I) iodide) sulfate (SO42–) Add dilute hydrochloric acid, Gas produced which turns damp [in solution] warm gently (gas produced) red litmus paper blue (ammonia) sulfite (SO32-) White precipitate forms (barium sulfate) Gas produced which will turn acidified potassium manganate (VII) from purple to colourless (sulfur dioxide) Analysis of Halide Ions using Silver and Lead Cations Metal Cation Added Chloride Bromide Iodide silver(I) ions cream precipitate of white precipitate of yellow precipitate of silver(I) chloride silver(I) bromide silver(I) iodide bright yellow lead(II) ions white precipitate of white precipitate of lead(II) chloride lead(II) bromide precipitate of lead(II) iodide 174

NES/Chemistry/IGCSE Other Notes on Analysis  Transition metal compounds are colourful. Most other chemicals are white as solids and colourless as solutions.  Metal oxides are usually basic and will form aqueous salts when reacted with acids.  Aluminium oxide, zinc oxide and lead (II) oxide are amphoteric and will react with both acids and bases.  Hydrated compounds contain water of crystallisation which will evaporate as steam when the solid chemical is heated. This steam will condense at the top of the test tube as water. The solid, anhydrous powder remaining is often a different colour to the hydrated compound.  Hydrogen peroxide (colourless liquid) is an oxidising agent (can turn Fe2+ to Fe3+)  Hydrogen peroxide decomposes slowly to give off oxygen gas. This can be made decompose faster by adding a catalyst, such as Fe2+.  Ammonium chloride (white solid) decomposes when heated to ammonia gas (basic) and hydrogen chloride gas (acidic)  Organic compounds (not carboxylic acids) are flammable.  Carboxylic acids have a smell like vinegar  Esters have fruity smells 175

NES/Chemistry/IGCSE Topic 9 - The Periodic Table 9.1 The Periodic Table  Describe the Periodic Table as a method of classifying elements and its use to predict properties of elements 9.2 Periodic trends  Describe the change from metallic to non-metallic character across a period  Describe and explain the relationship between Group number, number of outer shell electrons and metallic/nonmetallic character 9.3 Group properties  Describe lithium, sodium and potassium in Group I as a collection of relatively soft metals showing a trend in melting point, density and reaction with water  Predict the properties of other elements in Group I, given data, where appropriate  Describe the halogens, chlorine, bromine and iodine in Group VII, as a collection of diatomic non-metals showing a trend in colour and density and state their reaction with other halide ions  Predict the properties of other elements in Group VII, given data where appropriate  Identify trends in Groups, given information about the elements concerned 9.4 Transition elements  Describe the transition elements as a collection of metals having high densities, high melting points and forming coloured compounds, and which, as elements and compounds, often act as catalysts  Know that transition elements have variable oxidation states 176

NES/Chemistry/IGCSE 9.5 Noble gases  Describe the noble gases, in Group VIII or 0, as being unreactive, monatomic gases and explain this in terms of electronic structure  State the uses of the noble gases in providing an inert atmosphere, i.e. argon in lamps, helium for filling balloons 177

NES/Chemistry/IGCSE 9.1 - The Periodic Table The periodic table arranges elements in order of ascending proton number and can be used to predict properties of elements. Across a period (from left to right) the atoms of the elements have the same number of electron shells, but the number of valence electrons increases by 1. On the Periodic Table the metallic elements are on the left hand side, and the non-metallic elements are on the right hand side. The properties of the different elements depends on their position on the Periodic Table. 9.2 - Periodic Trends The elements on the left (and in the middle) of the Periodic Table are metals and the elements on the right of the Periodic Table are non-metals. As we move across the Periodic Table (from left to right), the elements change in character from metallic to non- metallic. So the most metallic metals are in Group I. There is a diagonal line that can be drawn on the Periodic Table dividing it into metals on the left and non-metals on the right The Periodic Table is also made up of rows called periods and columns called groups Periods Each period on the Periodic Table is a row of elements with the same number of electron shells. So moving down the groups the atoms get bigger as they have more shells of electrons. Period Number Maximum Number of Number of Elements of Shells Electrons in the Valence Shell in the Period 1 2 2 1 2 8 3 2 8 8 3 8 Groups The group number tells us the number of valence electrons in an atom. Elements in the same group have similar chemical properties and therefore react similarly. Going down a group, the number of shells increases by 1, but the number of valence electrons remains constant. So all Group I metals have 1 electron in their valency shell, which is why they all react in a similar way. 178

NES/Chemistry/IGCSE 9.3 - Group Properties Group Number Name of Group Structure of Element I Alkali metals II Alkaline Earth metals giant metallic VII giant metallic 0 Halogens simple molecular (no number) Noble gases monatomic hydrogen Transition metals giant metallic (no Group) none simple molecular Group I - Alkali Metals All the group members react with water to form strong alkaline solutions (fully dissociated) with a pH in the range of 12-14. The alkali metals are stored in oil to prevent them from reacting with oxygen and water vapour. Physical Properties of Group I Metals  Good conductors of electricity as there are mobile valence electrons  Are very soft and can be cut with a knife  Li, Na and K are less dense than water and so float Trends Down Group I Metals  Melting point and boiling point decrease down the group  Density increases down the group  Softness increases down the group 179

NES/Chemistry/IGCSE Decreasing Melting and Boiling Point Down Group I The atomic radius increases down the group, and the valence electrons are in higher shells. The more shells there are, the more ‘shielding’ there is of the valence electrons from the nucleus. The electrostatic force of attraction between the nucleus and valence electrons decreases therefore less energy is needed to overcome the forces of attraction resulting in a lower melting point and boiling point. Chemical Reactions of Group I Metals with Water Element Reaction with Water  Observations Reacts readily Li  effervescence and bubbles of a Na 2Li + 2H2O  2LiOH + H2 colourless gas evolved  floats and moves around the K Reacts vigorously  surface of the water 2Na + 2H2O  2NaOH + H2 metal decreases in size  effervescence and bubbles of a Reacts violently colourless gas evolved 2K + 2H2O  2KOH + H2  floats and moves quickly around  the surface of the water metal melts and decreases in size  effervescence and bubbles of a colourless gas evolved  floats and moves very quickly around the surface of the water metal burns with a lilac flame and rapidly decreases in size In order for a metal to form its ion it must lose its 1 valence electron. Energy must be supplied (endothermic) to do this. The stronger the electrostatic force of attraction between the nucleus and the valence electrons, the more energy has to be supplied and so the less reactive the metal is. The factors that affect the strength of the electrostatic force of attraction are:  Distance of electron from nucleus  Shielding by other electrons in closer shells  Number of protons in the nucleus (nuclear charge) The reactivity increases down the group (lose their 1 valence electron more easily) and therefore the reducing power increases down the group. 180

NES/Chemistry/IGCSE Group VII - Halogens Group VII elements are all diatomic molecules. Physical Properties of the Halogens Halogen Colour State at Room Formula Temperature pale yellow F2 green gas Cl2 gas Br2 red brown liquid I2 black solid  Don’t conduct electricity as there are no mobile electrons or mobile ions  Density increases down the group  Melting point and boiling point increase down the group  All the members of the group are diatomic Increasing Melting and Boiling Point Down Group VII Going down the group the number of electrons increases and the size of the diatomic molecules increase therefore the intermolecular van der Waals forces of attraction (force between molecules) increase and so more energy has to be supplied to overcome these forces resulting in the melting and boiling points increasing. Displacement Reactions of Group VII A halogen atom higher up Group VII will displace/oxidise a halide lower down in the group. Halogen Displacement Reactions Relative oxidising power F2 displaces chlorine, bromine, oxidises Cl–, Br– and I– to form Cl2, Br2 and I2 Cl2 iodine, astatine Br2 I2 displaces bromine, iodine, oxidises Br– and I– to form Br2 and I2 At2 astatine displaces iodine, astatine oxidises I– to form I2 oxidises At– to form At2 displaces astatine only Cannot displace any halide cannot oxidise any halide 181

NES/Chemistry/IGCSE Results Table for the Displacement Reactions Displacement Reaction Colour change Cl2 + 2KBr  2KCl + Br2 colourless to orange brown colourless to dark brown solution Cl2 + 2KI  2KCl + I2 Br2 + 2KCl  No reaction orange brown solution remains Br2 + 2KI  2KBr + I2 colourless to dark brown solution I2 + 2KCl  No reaction dark brown solution remains I2 + 2KBr  No reaction dark brown solution remains Decreasing Reactivity (oxidising power) down Group VII Non-metal atoms gain electrons in order to have a full valence shell. As we go down Group VII the number of shells of electrons increases. As there are more shells, the electrostatic force of attraction between the nucleus and the electrons decrease therefore larger atoms are less likely to take in electrons than the smaller halogens.  Reactivity decreases down the group  Oxidising power decreases down the group (F2 is the strongest oxidising agent) Uses of the Halogens Uses of chlorine Uses of bromine Uses of iodine  used to sterilise water  to make anti-freeze  to make antiseptics  to make pvc (a polymer)  to make domestic bleach 182

NES/Chemistry/IGCSE 9.4 - The Transition elements These are described as a collection of metals having:  high densities physical properties of transition metals  high melting points  variable valences chemical properties of transition metals  form coloured compounds except Zn  the elements or compounds of the transition metals act as catalysts Variable Valency There are various transition metal compounds. In naming these compounds it is essential to put the oxidation state in the name as a roman numeral in order for us to know the state of oxidation of the transition metal cation. For example we have to say copper(I) oxide, because there is also a compound called copper(II) oxide. The same goes for iron. Iron(II) oxide is a green solid, whereas iron(III) oxide is a red/brown solid. Not all elements in the central block behave like transition metals.  Example Zinc is volatile - It has a low boiling point. Zinc does not have variable valency - It has one valency of 2. Zinc hydroxide is a white precipitate, not coloured. Uses of Copper (a typical transition metal)  In electrical wiring as it is a good conductor of electricity and is ductile.  Cooking utensils as it is a good conductor of heat.  Domestic heating systems 183

NES/Chemistry/IGCSE 9.5 - Noble Gases Group 0 elements are monatomic gases that make up about 1% of air. Properties of Group 0  Don’t conduct electricity  Melting point and boiling point increase down the group  Density increases down the group  Do not take part in chemical reactions Increasing Melting Point and Boiling Point down Group 0 Going down Group 0, the size of the atoms increase and there are more electrons in the atoms, so the van der Waals force of attraction between monatomic atoms increases, therefore more energy is required to overcome the forces. Melting and boiling points increase down Group 0. Unreactive The noble gases have full valence shells, and so do not lose, gain or share electrons. This means that Group 0 elements are chemically unreactive. Uses of the Noble Gases  Helium is used to fill airships or weather balloons because it is less dense than air and is unreactive. Hydrogen is also less dense than air, but is very explosive.  Argon is used to provide an inert atmosphere in lamps, so the filament does not burn out.  Neon is used in lamps as it glows when an electrical current passes through it. 184

NES/Chemistry/IGCSE Topic 10 - Metals 10.1 Properties of metals  List the general physical properties of metals  Describe the general chemical properties of metals e.g. reaction with dilute acids and reaction with oxygen  Explain in terms of their properties why alloys are used instead of pure metals  Identify representations of alloys from diagrams of structure 10.2 Reactivity series  Place in order of reactivity: potassium, sodium, calcium, magnesium, zinc, iron, (hydrogen) and copper, by reference to the reactions, if any, of the metals with: – water or steam – dilute hydrochloric acid and the reduction of their oxides with carbon  Deduce an order of reactivity from a given set of experimental results  Describe the reactivity series as related to the tendency of a metal to form its positive ion, illustrated by its reaction, if any, with: – the aqueous ions – the oxides of the other listed metals  Describe and explain the action of heat on the hydroxides, carbonates and nitrates of the listed metals  Account for the apparent unreactivity of aluminium in terms of the oxide layer which adheres to the metal 185

NES/Chemistry/IGCSE 10.3 Extraction of metals  Describe the ease in obtaining metals from their ores by relating the elements to the reactivity series  Describe and state the essential reactions in the extraction of iron from hematite  Describe the conversion of iron into steel using basic oxides and oxygen  Know that aluminium is extracted from the ore bauxite by electrolysis  Discuss the advantages and disadvantages of recycling metals, limited to iron/steel and aluminium  Describe in outline, the extraction of zinc from zinc blende  Describe in outline, the extraction of aluminium from bauxite including the role of cryolite and the reactions at the electrodes 10.4 Uses of metals  Name the uses of aluminium: – in the manufacture of aircraft because of its strength and low density – in food containers because of its resistance to corrosion  Name the uses of copper related to its properties (electrical wiring and in cooking utensils)  Name the uses of mild steel (car bodies and machinery) and stainless steel (chemical plant and cutlery)  Explain the uses of zinc for galvanising and for making brass  Describe the idea of changing the properties of iron by the controlled use of additives to form steel alloys 186

NES/Chemistry/IGCSE 10.1 - Properties of Metals An alloy is a mixture of a metal and at least one other substance Metals are elements on the left hand side of the Periodic Table. Physical Properties Chemical Properties Good conductor of heat and electricity Usually form basic oxides (bases) Shiny Form positive ions Malleable React with acid to form salt and hydrogen (except copper, silver, gold and platinum) Sonorous React with non-metals forming ionic bonds Ductile Usually have high melting and boiling points (except Group I) Usually high density (except Group I) Malleable means can be hammered and bent into shapes without breaking. Ductile means can be pulled into thin wires. Sonorous means makes a ringing sound when hit. Alloys Alloys are mixtures, not compounds. They are made molten and mixed together, rather than react to form a compound. Alloys have the properties of the metals used to make them. So the properties of a metal can be modified by mixing it with other metals. positive ions delocalised mobile valence 'sea' of electrons 187

NES/Chemistry/IGCSE Testing for Metals (or alloys) 1. Physical Test - conducts electricity cell bulb Test - place in a circuit Result - bulb will light up place sample to be tested across here 2. Chemical Test - reacts with acid to make hydrogen All metals above hydrogen in the reactivity series will react with acid. There will be effervescence and the gas made will give a 'squeaky pop' with a lit splint. 188

NES/Chemistry/IGCSE 10.2 Reactivity Series The reactivity of a metal is a measure of its ability to form positive ions. A reactive metal like potassium will readily lose its valence electron to form a cation. K  K+ + e– A less reactive metal will slowly, if at all, to form a cation. Cu  Cu2+ + 2e – Reactivity Series The reactivity series is when metals are listed in their order of reactivity. You will notice that Group I is the most reactive, followed by Group II, Group III and then the Transition Metals. Metal Position on Periodic Table Reactivity K Group I Most reactive Na Ca Group II Least reactive Mg Group III Al Group IV Transition metal *Carbon No Group/Position Zn Fe Transition metal *Hydrogen Cu Ag Au Pt *Carbon and hydrogen are included in the list because they can also form positive ions. 189

NES/Chemistry/IGCSE Reactions of Metals 1. Reaction with oxygen 2. Reaction with water (or steam) 3. Reaction with acid 4. Displacement reactions with metal compounds 5. Thermal decomposition 1. Reaction with Oxygen All metals will react with oxygen to form a metal oxide. metal + oxygen  metal oxide 2. Reaction with Water The more reactive metals react with cold water (from potassium to magnesium in the reactivity series). The Group I metals are so reactive the reaction is dangerous and can only be done by your teacher. metal + water  metal hydroxide + hydrogen Less reactive metals will react with steam as it is hotter, but the products of the reaction are different. metal + steam  metal oxide + hydrogen Metals below hydrogen will not react with water, or steam. 3. Reaction with Acid Group I metals are too reactive to react with acids - they will explode. Metals below hydrogen in the reactivity series will not react with acid. metal + acid  salt + hydrogen 190

NES/Chemistry/IGCSE Reaction Summary Metal Reaction with Reaction with Water, or Reaction with Acids Potassium Oxygen Steam violent reaction to Sodium burns with a lilac reacts violently with cold form a salt and Calcium flame to form water producing potassium hydrogen potassium oxide hydroxide and hydrogen Magnesium burns with a vigorous reaction to yellow flame to reacts vigorously with cold form a salt and Zinc form a sodium water producing sodium hydrogen Iron oxide hydroxide and hydrogen Copper burns with a red reacts with dilute flame to form a reacts with cold water acids to form a salt calcium oxide producing calcium hydroxide and hydrogen and hydrogen burns with a white reacts very slowly with cold reacts with dilute flame to form water to form magnesium acids to form a salt magnesium oxide hydroxide and hydrogen. and hydrogen Magnesium is heated to react forms zinc oxide it with steam to produce reacts with dilute on heating magnesium oxide and acids to form a salt hydrogen and hydrogen reacts slowly heated to react with steam reacts with dilute producing zinc oxide and acids to form a salt forming iron(III) hydrogen and hydrogen reacts slowly with steam oxide producing iron(III) oxide and no reaction hydrogen reacts very slowly no reaction forming copper(II) oxide 191

NES/Chemistry/IGCSE 4. Displacement Reactions A reactive metal will displace, or reduce the less reactive metal ions from a solution of its salt.  Example 1: zinc metal reacting with copper(II) sulphate Zinc is a more reactive metal that copper, so it will displace the copper(II) ions from the compound. Symbol equation zinc + copper(II) sulfate  zinc sulfate + copper Ionic equation Zn + CuSO4  ZnSO4 + Cu Ionic half-equation Zn  Zn2+ + 2e- Cu2+ + 2e-  Cu Observations The zinc decreases in size A pink brown solid forms The reaction is exothermic The blue colour of the solution fades to colourless The bigger the difference in reactivity between the two metals the faster and more exothermic the reaction is.  Example 2: silver reacting with calcium chloride No reaction happens because silver is less reactive than calcium, so silver cannot displace the calcium ions from the compound. 192

NES/Chemistry/IGCSE 5. Thermal Decomposition Thermal decomposition reactions are when one compound breaks down into two, or more, smaller chemicals when heated. Group Nitrates Hydroxides Carbonates decompose do not decompose do not decompose Group I  a metal nitrite oxygen gas (except lithium)  2NaNO3  2NaNO2 + O2 decompose decompose  metal oxide decompose  CO2 Group II, lithium  a metal oxide  metal oxide MgCO3  MgO + CO2 and the  NO2 gas  water transition  oxygen gas metals 2Mg(NO3)2  2MgO + 4NO2 + O2 2LiOH  Li2O + H2O Thermal stability is the resistance of a chemical to decompose when heated. More thermally stable compounds required a higher temperature to react. 193

NES/Chemistry/IGCSE 10.3 Extraction of Metals Metals are found either as metals, or in compounds in rocks in the ground. Metals found native, or uncombined, are metals found as elements; rather than in compounds. Silver, gold and platinum are usually found native and do not need to be processed. They are just dug out of the ground and used directly. An ore is a rock which contains a compound from which a metal can be extracted. Ore Compound Processed by haematite iron (III) oxide carbon reduction zinc blende carbon reduction zinc sulfide bauxite aluminium oxide electrolysis The ores must be processed in order to get the metals out of them. There are two different methods of extracting metals from ores, based on the metal's reactivity 1. Metals less reactive than carbon can be extracted in a blast furnace using carbon reduction. 2. Metals more reactive than carbon are extracted using electrolysis. A blast furnace is like a very large, hot oven. Hot air is blown in at the sides to increase the temperature. 194

NES/Chemistry/IGCSE 1. Carbon Reduction For metals less reactive than carbon (zinc to platinum), carbon reduction is used to extract the metal. Carbon, or carbon monoxide is a reducing agent, which reduces the metal ions in the compound to metal atoms. a. Iron iron ore (hematite) limestone coke CO2(g) + N2(g) 700 oC 1100 oC hot air in 1900 oC molten calcium hot air in silicate (slag) out molten pig iron out Haematite is a mixture of iron(III) oxide and impurities, mainly sand (SiO2). Coke is a pure form of coal - it contains mainly carbon Limestone is a rock from the ground - it contains mainly calcium carbonate Haematite, coke and limestone enter at the top of the furnace. Carbon dioxide gas and nitrogen gas leave the furnace as waste gases. Carbon monoxide is the reducing agent and reduces iron(III) oxide to iron in the reaction, and therefore will be oxidised to carbon dioxide gas. 195

NES/Chemistry/IGCSE Reactions in the Blast Furnace Coke (combustion to produce heat and carbon monoxide)  C + O2  CO2  CO2 + C  2CO Limestone (thermal decomposition to form calcium oxide)  CaCO3  CaO + CO2 Haematite (reduction of iron(III) oxide)  Fe2O3 + 3CO  2Fe + 3CO2  2Fe2O3 + 3C  4Fe + 3CO2 Silicon(IV) oxide, an acidic impurity is removed by basic calcium oxide forming slag, calcium silicate.  CaO + SiO2  CaSiO3 Slag is used to make roads and cement; as well as electronic chips. The iron that is produced (called cast iron) contains about 4% carbon impurities. Cast iron is quite brittle and is usually processed again to make steel. Steel is more malleable and stronger than cast iron, making it a more useful material. 196

NES/Chemistry/IGCSE b. Steel oxygen gas in CO(g) SO2(g) + CO2(g) C Fe S Si SiO2 In the Basic Oxygen Furnace, oxygen is blown into the furnace to oxidise the impurities (the iron is not oxidised as it is the least reactive element present). Calcium oxide is also added to remove acidic impurities. The carbon content in steel is less than 1% making it more malleable than cast iron with 4% carbon. Reactions in the Basic Oxygen Furnace Carbon  C + O2  CO2 Sulfur  S + O2  SO2 These gases leave the top of the furnace. Silicon  Si + O2  SiO2  CaO + SiO2  CaSiO3 197

NES/Chemistry/IGCSE c. Zinc Zinc blende is roasted in air first as zinc sulfide cannot be reduced by carbon, but zinc oxide can be reduced by carbon.  2ZnS + 3O2  2ZnO + 2SO2 The sulfur dioxide is used to make sulfuric acid in the Contact Process (Topic 12). The zinc oxide is then reduced in the blast furnace to zinc gas (zinc has quite a low boiling point). The zinc gas is then collected in side arms and condensed to liquid zinc. zinc oxide + coke Zn(g) liquified zinc impurities hot air in hot air in Reactions in the Blast Furnace Coke (combustion to produce heat and carbon monoxide)  C + O2  CO2  CO2 + C  2CO Zinc blende  ZnO + C  Zn(g) + CO 198

NES/Chemistry/IGCSE 10.4 Uses of Metals Composition of alloys: Alloy Composition Steel Iron and Carbon Stainless steel Iron, Carbon, Chromium and Nickel Brass Copper and Zinc Bronze Copper and Tin Alloys are always stronger than the metals they are made from. They are also less malleable. This is because of the different sized ions in the lattice that make it harder for the layers to slide over each other. Uses of metals and alloys: Property Uses Metal, or Alloy low density / strong aircraft manufacture Aluminium corrosion resistance electrical conductor food containers Copper thermal conductor / electrical wiring corrosion resistant / Zinc cooking utensils Mild Steel malleable Stainless Steel strong / ductile making brass reactive metal galvanising car bodies strong chemical plants / cutlery resists corrosion 199