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KS4 Chemistry Chemical Reactions 1 of 69 © Boardworks Ltd 2005 Contents Chemical Reactions Introducing chemical reactions Thermal decomposition Displacement and precipitation Neutralization Redox Summary activities 2 of 69 © Boardworks Ltd 2005 What is a chemical reaction? What is a chemical reaction? A chemical reaction is a change that takes place when one or more substances (called reactants) form one or more new substances (called products). chemical reaction reactants products There are many different types of chemical reactions. 3 of 69 © Boardworks Ltd 2005 Types of chemical reaction How many types of chemical reaction can you name? 4 of 69 © Boardworks Ltd 2005 Types of chemical reaction oxidation & reduction neutralization thermal decomposition chemical reaction exothermic & endothermic displacement: metals displacement: non-metals precipitation reversible 5 of 69 © Boardworks Ltd 2005 Exothermic and endothermic reactions What are exothermic and endothermic reactions? exothermic reactions give out energy – they get hot ex = out (as in ‘exit’) thermic = relating to heat endothermic reactions take in energy – they get cold en = in (as in ‘entrance’) Most chemical reactions are exothermic. 6 of 69 © Boardworks Ltd 2005 Examples of exothermic reactions Many exothermic reactions occur in the lab and in everyday life. Can you think of six exothermic reactions? Exothermic reactions Burning wood on a fire Burning petrol in a car Burning gas on a gas hob Reacting an acid and alkali together Burning magnesium Rotting compost 7 of 69 © Boardworks Ltd 2005 Irreversible reactions Most chemical reactions are considered to be irreversible because the products cannot easily be changed back into reactants. For example, once magnesium has reacted with hydrochloric acid, it is difficult to get the magnesium back. magnesium + hydrochloric magnesium acid chloride + hydrogen In equations for irreversible reactions, reactants and products are joined by a ‘one-way’ arrow. 8 of 69 © Boardworks Ltd 2005 Reversible reactions Although most chemical reactions are difficult to reverse, there are some reactions that are fully reversible. One of the best known reversible reactions occurs when copper sulfate crystals are heated. hydrated copper sulfate anhydrous copper sulfate + water CuSO4.5H2O CuSO4 + 5H2O In equations for reversible reactions, reactants and products are joined by a ‘two-way’ arrow. 9 of 69 © Boardworks Ltd 2005 Equilibrium reactions In some reversible reactions, the forward and backward reactions largely occur in the same conditions and at the same rate. These reaction are said to be in equilibrium – there is no overall change in the amount of products and reactants. One of the most important equilibrium reactions occurs in the production of ammonia in the Haber process: nitrogen + hydrogen ammonia N2 (g) + 3H2 (g) 2NH3 (g) No matter how long the reaction is left, there will always be a mixture of nitrogen, hydrogen and ammonia. 10 of 69 © Boardworks Ltd 2005 Contents Chemical Reactions Introducing chemical reactions Thermal decomposition Displacement and precipitation Neutralization Redox Summary activities 11 of 69 © Boardworks Ltd 2005 Thermal decomposition Thermal decomposition is a reaction in which a compound is broken down by heat into two or more simpler substances. Generally, the more reactive a metal, the harder it is to decompose its compounds by heating. For example: silver carbonate: decomposes on gentle heating calcium carbonate: decomposes on strong heating increase in reactivity of metal potassium carbonate: is not thermally decomposed 12 of 69 © Boardworks Ltd 2005 Thermal decomposition – easy or hard? Compound Decomposition mercury oxide easy sodium oxide hard iron oxide medium silver oxide easy zinc oxide medium 13 of 69 potassium sodium calcium magnesium aluminium zinc iron copper mercury silver gold increase in reactivity How easy will these metal compounds be to decompose: easy, medium or difficult? © Boardworks Ltd 2005 Thermal decomposition of carbonates When metal carbonates are heated, they decompose to produce metal oxides and carbon dioxide. This reaction is performed industrially to make calcium oxide (quicklime) from calcium carbonate (limestone): calcium carbonate heat calcium oxide + carbon dioxide CaCO3 CaO + CO2 Quicklime is used to make concrete and calcium hydroxide (slaked lime). 14 of 69 © Boardworks Ltd 2005 Thermal decomposition of metal oxides Most metal oxides are thermally stable – they do not decompose when heated. Oxides of the least reactive metals can be thermally decomposed more easily. For example, mercury oxide decomposes when heated strongly: mercury oxide 2HgO heat mercury condenses at the top of the test tube, where it is cooler 15 of 69 mercury + oxygen 2Hg + O2 oxygen gas escapes © Boardworks Ltd 2005 Contents Chemical Reactions Introducing chemical reactions Thermal decomposition Displacement and precipitation Neutralization Redox Summary activities 16 of 69 © Boardworks Ltd 2005 The reactivity series potassium sodium calcium magnesium aluminium zinc iron copper mercury silver gold 17 of 69 increase in reactivity The reactivity series is a list of metals in order of their reactivity. The reactivity series can be used to make predictions about the reactivity of metals – for example, how a metal will react with oxygen, water and acids. © Boardworks Ltd 2005 Displacement reactions: metals A metal displacement reaction occurs when a metal is added to a compound of a less reactive metal. more reactive metal + compound less of less reactive reactive metal metal + compound of more reactive metal The less reactive metal is displaced from the compound and becomes elemental metal. The more reactive metal forms a new compound. A metal will always displace another metal that is lower in the reactivity series. 18 of 69 © Boardworks Ltd 2005 Displacement reactions – examples Will magnesium react with copper chloride? magnesium + copper magnesium chloride chloride + copper Magnesium is more reactive than copper, so it displaces copper from its compound. Will silver react with magnesium chloride? silver + magnesium chloride no reaction Silver is less reactive than magnesium, so it does not displace magnesium from its compound. 19 of 69 © Boardworks Ltd 2005 The Thermit process The Thermit process is a displacement reaction between aluminium and iron (III) oxide. magnesium fuse iron oxide aluminium powder (thermite) aluminium Al + iron oxide aluminium oxide + iron + + Fe2O3 Al2O3 Fe Aluminium is more reactive than iron and displaces it from the oxide. 20 of 69 © Boardworks Ltd 2005 The Thermit process The reaction between aluminium and iron oxide is so exothermic that the displaced iron melts. The reaction is used to weld iron and steel together; for example, railway tracks. 21 of 69 © Boardworks Ltd 2005 Is there a displacement reaction? 22 of 69 © Boardworks Ltd 2005 Displacement reactions: halogens A halogen displacement reaction occurs when a halogen is added to a metal halide containing a less reactive halogen. The less reactive halogen is displaced from the compound and the more reactive halogen bonds with the metal to form a new metal halide. F Cl decrease in reactivity Br I For example: 23 of 69 fluorine + F2 (aq) + sodium chloride sodium fluoride 2NaCl (aq) 2NaF (aq) + chlorine + Cl2 (aq) © Boardworks Ltd 2005 Displacement reactions of halogens 24 of 69 © Boardworks Ltd 2005 Is there a displacement reaction? 25 of 69 © Boardworks Ltd 2005 Precipitation reactions When two aqueous solutions are mixed, they may react to form a product that is insoluble in water. The solid is called a precipitate and the reaction is called a precipitation reaction. To predict whether a precipitation reaction will occur, information on the solubility of the products is needed. What are the symbols for these physical states? solid (g) liquid (l) gas (g) aqueous (aq) (dissolved in water) 26 of 69 © Boardworks Ltd 2005 Precipitation reactions: sulfur The precipitation reaction between solutions of sodium thiosulfate and hydrochloric acid is often used to measure rates of reaction. sodium sodium hydrochloric chloride thiosulfate + acid Na2S2O3 (aq) + 2HCl (aq) Both reactants are colourless. 27 of 69 2NaCl (aq) + sulfur dioxide + SO2 (g) + sulfur + S (s) + water + H2O (l) Sulfur is insoluble and precipitates, turning the solution cloudy. © Boardworks Ltd 2005 Precipitation reactions: copper hydroxide Many metal hydroxides are insoluble and can be formed by precipitation reactions. For example: copper (II) sulfate CuSO4 (aq) + + copper (II) ammonium hydroxide hydroxide 2NH4OH (aq) Copper (II) sulfate solution is blue. 28 of 69 Cu(OH)2 (s) + ammonium sulfate + (NH4)2SO4 (aq) Copper (II) hydroxide is insoluble and forms a blue solid at the bottom. © Boardworks Ltd 2005 Precipitation reactions: iron hydroxide Iron (III) hydroxide is another insoluble metal hydroxide that can be formed by a precipitation reaction. iron (III) chloride + FeCl3 (aq) + iron (III) hydroxide + sodium chloride 3NaOH (aq) Fe(OH)3 (s) + 3NaCl (aq) sodium hydroxide Iron (III) chloride solution is yellow. 29 of 69 Iron (III) hydroxide is insoluble and forms a deep brown solid at the bottom. © Boardworks Ltd 2005 Precipitation and solubility To help work out whether a precipitate will form in a reaction, there are some general rules about solubility. Soluble All compounds of sodium, potassium and ammonium. All nitrates. Most chlorides, except silver and lead chlorides. Most sulfates, except lead, barium and calcium sulfates. Insoluble Most carbonates and hydroxides, except those of sodium, potassium and ammonium. 30 of 69 © Boardworks Ltd 2005 Predicting precipitation reactions There are three steps to working out whether a precipitate will be formed in a reaction: Example 1. Write down the names sodium chloride & lead nitrate of the reactants. 2. Swap over the non-metal. sodium nitrate & lead chloride 3. Are the products soluble or insoluble? Lead chloride is insoluble and will form a precipitate. sodium chloride + 2NaCl (aq) + 31 of 69 lead nitrate sodium nitrate Pb(NO3)2 (aq) 2NaNO3 (aq) + lead chloride + PbCl2 (s) © Boardworks Ltd 2005 Will there be a precipitation reaction? (1) Will a precipitate be formed when sodium nitrate and magnesium sulfate react? 1. Write down the names of the reactants. sodium nitrate & magnesium sulfate 2. Swap over the non-metal. sodium sulfate & magnesium nitrate 3. Are the products soluble or insoluble? Both products are soluble so no precipitate will form. sodium sulfate sodium nitrate + magnesium sulfate 2NaNO3 (aq) + MgSO4 (aq) Na2SO3 (aq) 32 of 69 + magnesium nitrate + Mg(NO3)2 (aq) © Boardworks Ltd 2005 Will there be a precipitation reaction? (2) Will a precipitate be formed when sodium sulfate and barium nitrate react? 1. Write down the names of the reactants. sodium sulfate & barium nitrate 2. Swap over the non-metal. sodium nitrate & barium sulfate 3. Are the products soluble or insoluble? Barium sulfate is insoluble and will form a precipitate. sodium sulfate + Na2SO4 (aq) + 33 of 69 barium nitrate sodium nitrate Ba(NO3)2 (aq) 2NaNO3 (aq) + barium sulfate + BaSO4 (s) © Boardworks Ltd 2005 Precipitation: true or false? 34 of 69 © Boardworks Ltd 2005 Contents Chemical Reactions Introducing chemical reactions Thermal decomposition Displacement and precipitation Neutralization Redox Summary activities 35 of 69 © Boardworks Ltd 2005 Acids What are acids? They are substances that: Have a pH below 7. Turn litmus red. Turn universal indicator yellow, orange or red. Form solutions containing H+ ions. The more H+ ions in the solution, the stronger the acid. Universal indicator pH scale 1 2 3 4 5 6 strong acid 7 8 9 10 11 12 13 14 weak acid neutral 36 of 69 © Boardworks Ltd 2005 Common acids What are some common acids? Acid Formula Strength sulfuric acid H2SO4 strong hydrochloric acid HCl strong nitric acid HNO3 strong ethanoic acid (vinegar) CH3COOH weak 1 2 3 4 5 6 37 of 69 7 8 9 10 11 12 13 14 © Boardworks Ltd 2005 Bases What are bases? They are substances that: Have a pH above 7. Turn litmus blue. Turn universal indicator dark green, blue or purple. Are capable of neutralizing acids. Universal indicator pH scale 1 2 3 4 5 6 7 8 9 10 11 12 13 14 weak base strong base neutral 38 of 69 © Boardworks Ltd 2005 More about bases Bases are usually oxides, hydroxides or carbonates of metals. Ammonia is a base that doesn’t contain a metal. Some bases are soluble in water – they are called alkalis. bases pH > 7 alkalis soluble bases 39 of 69 Most carbonates and hydroxides are insoluble in water, except those of sodium, potassium and ammonium. All alkaline solutions contain OH– ions. The more OH– ions in the solution, the stronger the alkali. © Boardworks Ltd 2005 Common bases What are some common bases? Base Formula Strength sodium hydroxide * NaOH strong potassium hydroxide * KOH strong calcium hydroxide * Ca(OH)2 strong ammonia * NH3 weak calcium carbonate CaCO3 weak * = the base is also an alkali 1 2 3 4 5 6 40 of 69 7 8 9 10 11 12 13 14 © Boardworks Ltd 2005 Neutralization reactions A neutralization reaction occurs when an acid reacts with a base or alkali to produce a neutral solution of salt and water. acid 41 of 69 + alkali salt + water © Boardworks Ltd 2005 Naming salts The salt formed in a neutralization reaction takes its name from both the base and the acid. The first part of the salt comes from the first part of the base, for example: magnesium oxide magnesium salt ammonium hydroxide ammonium salt The second part of the salt comes from the acid: 42 of 69 sulfuric acid sulfate hydrochloric acid chloride nitric acid nitrate © Boardworks Ltd 2005 What is the name of the salt? What are the names of the salts formed from these bases and acids? Base Acid Salt calcium hydroxide hydrochloric acid calcium chloride magnesium oxide nitric acid magnesium nitrate copper oxide sulfuric acid copper sulfate aluminium hydroxide nitric acid aluminium nitrate potassium hydroxide sulfuric acid potassium sulfate 43 of 69 © Boardworks Ltd 2005 Neutralization reactions: hydroxides When a hydroxide is mixed with an acid, OH– ions react with H+ ions from the acid to form water: H+ + OH– H2O For example: potassium hydroxide + KOH (aq) + calcium hydroxide Ca(OH)2 (aq) 44 of 69 potassium hydrochloric chloride acid HCl (aq) + sulfuric acid + water H2O (aq) KCl (aq) + calcium sulfate + + H2SO4 (aq) CaSO4(aq) water + 2H2O (aq) © Boardworks Ltd 2005 Neutralization reactions: oxides When an oxide is mixed with an acid, O2– ions react with H+ ions from the acid to form water: 2H+ + O2– H2O For example: 45 of 69 calcium oxide + CaO (aq) + copper oxide + CuO (aq) + hydrochloric acid 2HCl (aq) sulfuric acid calcium chloride CaCl2 (aq) copper sulfate H2SO4 (aq) CuSO4 (aq) + water + H2O (aq) + water + H2O (aq) © Boardworks Ltd 2005 Neutralization reactions: carbonates When a carbonate is mixed with an acid, CO32– ions react with H+ ions from the acid to form water and carbon dioxide: 2H+ + CO32– H2O + CO2 For example: calcium carbonate CaCO3 (aq) 46 of 69 + calcium nitric nitrate acid + water + + 2HNO3 Ca(NO3)2 (aq) (aq) + H2O (aq) + carbon dioxide CO2 (g) © Boardworks Ltd 2005 Complete the neutralization reaction 47 of 69 © Boardworks Ltd 2005 Obtaining insoluble salts How can insoluble salts be obtained? 1. The acid and alkali are mixed, and the salt forms by precipitation. 2. The mixture of products is filtered, trapping the salt. 3. The salt is rinsed in distilled water, then left to dry. For example, barium sulfate can be obtained by mixing barium chloride with sulfuric acid. 48 of 69 barium hydroxide + BaOH2 (aq) + sulfuric acid barium sulfate H2SO4 (aq) BaSO4 (s) + water + 2H2O (aq) © Boardworks Ltd 2005 Obtaining soluble salts from bases How can soluble salts be obtained following a reaction between an acid and a base (insoluble)? For example, obtaining copper sulfate from copper oxide and sulfuric acid: 1. Copper oxide is added to sulfuric acid. When it is heated, the copper oxide dissolves, forming a blue solution. 2. More copper oxide is added until no more will dissolve. This means that all the acid has been used up. 3. The mixture is filtered to remove the excess copper oxide. 4. The filtrate is heated to evaporate some of the water. When it cools, copper sulfate crystals will form. 49 of 69 © Boardworks Ltd 2005 Obtaining soluble salts from alkalis (1) How can soluble salts be obtained following a reaction between an acid and an alkali (soluble)? For example, the reaction between sodium hydroxide and hydrochloric acid produces sodium chloride, which is soluble. sodium hydroxide + NaOH (aq) + hydrochloric acid sodium chloride + water NaCl (aq) + H2O (aq) HCl (aq) There is no obvious sign when the reaction is complete, so an indicator must be used to show when the solution is neutral. This process is called titration. 50 of 69 © Boardworks Ltd 2005 Obtaining soluble salts from alkalis (2) To run the titration: 1. 25 cm3 of sodium hydroxide is added to a flask. Two drops of the indicator phenolphthalein are added. This turns pink. 2. Hydrochloric acid is added to the flask, a little at a time, from a burette. 3. When all the alkali has reacted with the acid, the indicator turns colourless. The amount of acid used is noted. 4. The experiment is repeated, but without adding the indicator, as this makes the salt impure. 5. The salt solution is heated to evaporate the water. Crystals of sodium chloride will remain. 51 of 69 © Boardworks Ltd 2005 Matching reactants and salts 52 of 69 © Boardworks Ltd 2005 Neutralization: true or false? 53 of 69 © Boardworks Ltd 2005 Contents Chemical Reactions Introducing chemical reactions Thermal decomposition Displacement and precipitation Neutralization Redox Summary activities 54 of 69 © Boardworks Ltd 2005 What is redox? What does redox mean? reduction and oxidation So far, these terms have been used to describe the gain and loss of oxygen. 55 of 69 Oxidation = adding oxygen to a substance. Reduction = removing oxygen from a substance. E.g. production of iron oxide during rusting E.g. extracting iron from iron oxide in a blast furnace. © Boardworks Ltd 2005 Oxidation and ions When a metal burns, it is oxidized to form a metal oxide. What happens to the metal atoms during the reaction? Metal atoms lose electrons to form positive ions. The oxygen atoms accept these electrons and form negative ions. For example, when magnesium burns to form magnesium oxide, each magnesium atom loses 2 electrons and becomes a magnesium ion with a +2 charge. 2 electrons Mg 56 of 69 O Mg2+ O2- © Boardworks Ltd 2005 Oxidation and electron loss What has happened to magnesium when it reacts with oxygen? 2 electrons O Mg Mg2+ O2- It has been oxidized. It has lost electrons by changing from Mg to Mg2+. Oxidation is the loss of electrons. Magnesium can lose electrons to substances other than oxygen, e.g. when it reacts with chlorine or sulfur. These reactions both involve Mg Mg2+, so they are also oxidation. 57 of 69 © Boardworks Ltd 2005 Reduction and ions When iron is extracted from iron oxide (iron ore), the oxygen is removed and the iron is said to be reduced. When the oxygen is removed, 3 electrons are transferred back to each iron ion, which become atoms. O2O2O2- 58 of 69 Fe3+ 2 electrons from each ion Fe3+ O Fe O O Fe © Boardworks Ltd 2005 Reduction and electron gain What has happened to iron when oxygen is removed? O2O2O2- Fe3+ 2 electrons from each ion Fe3+ O Fe O O Fe It has been reduced. It has gained electrons by changing from Fe3+ to Fe. Reduction is the gain of electrons. 59 of 69 © Boardworks Ltd 2005 Electrons and OILRIG An easy way to remember what happens to electrons during oxidation and reduction is to think… Oxidation …OILRIG Is Loss …of electrons Reduction Is Gain …of electrons 60 of 69 © Boardworks Ltd 2005 Redox reactions When a substance is oxidized, it loses electrons. Another substance must gain these electrons and become reduced. For example, when magnesium burns: magnesium loses electrons: Mg Mg2+ = oxidation oxygen gains electrons: O O2– = reduction The overall reaction is reduction and oxidation = redox. Oxidation and reduction always take place together. 61 of 69 © Boardworks Ltd 2005 Oxidized or reduced? For each reaction, decide which product is oxidized and which product is reduced. calcium + oxidized zinc oxide reduced aluminium oxidized 62 of 69 oxygen calcium oxide reduced + hydrogen zinc + water oxidized + iron oxide iron + aluminium oxide reduced © Boardworks Ltd 2005 Oxidized or reduced? 63 of 69 © Boardworks Ltd 2005 Contents Chemical Reactions Introducing chemical reactions Thermal decomposition Displacement and precipitation Neutralization Redox Summary activities 64 of 69 © Boardworks Ltd 2005 Glossary (part 1) displacement – A type of reaction in which a metal or halogen replaces a less reactive metal or halogen in a compound. equilibrium – A type of reaction in which products are broken down at the same rate as reactants are combining. endothermic – A type of reaction that takes in energy. exothermic – A type of reaction that gives out energy. neutralization – A type of reaction in which an acid and base react to form a salt and water. oxidation – A type of reaction involving the loss of electrons. 65 of 69 © Boardworks Ltd 2005 Glossary (part 2) precipitation – A type of reaction in which two aqueous solutions react to form an insoluble product. reaction – A change that takes place when one or more substances form one or more new substances. redox – A type of reaction in which oxidation and reduction take place together. reduction – A type of reaction involving the gain of electrons. decomposition – A type of reaction in which a substance is broken down into two or more simpler substances. 66 of 69 © Boardworks Ltd 2005 Anagrams 67 of 69 © Boardworks Ltd 2005 Identify the reactions 68 of 69 © Boardworks Ltd 2005 Multiple-choice quiz 69 of 69 © Boardworks Ltd 2005