Chemical Reactions

1. Chemical Reactions

2. Types of chemical change thermal decomposition oxidation and reduction exothermic and endothermic displacement reactions: metals neutralisation displacement reactions: non-metals precipitation reversible reactions

3. A thermal decomposition is when heat causes a chemical to break down to simpler substances. Compounds – but not elements - undergo thermal decomposition. For compounds that contain metals we usually find: the more reactive the metal, the harder it is to decompose its compounds. For example: Potassium carbonate is not thermally decomposed. Calcium carbonate decomposes on strong heating Silver carbonate decomposes on gentle heating Gets harder Thermal decomposition

4. Generally, the more reactive the metal, the more difficult it is to decompose its compounds. Fill in the last column: easy, medium or hard. Thermal decomposition Potassium sodium calcium magnesium aluminium zinc iron copper mercury silver gold Increasing reactivity easy hard medium easy medium

5. When carbonates are heated they release carbon dioxide. This reaction is performed industrially to make calcium oxide (quicklime) from calcium carbonate (limestone). Quicklime is used to make concrete and to make calcium hydroxide (slaked lime). limestone waste air and carbon dioxide Hot air 1500°C calcium oxide (lime) Thermal decomposition of carbonates Calcium Carbonate Calcium oxide Carbondioxide + 

6. Most metal oxides are thermally stable (i.e. do not decompose when heated). Oxides of the least reactive metals can be thermally decomposed more easily. For example, silver oxide begins to break up at about 160oC and mercury oxide decomposes when heated strongly. mercury metal and oxygen formed O O O O Hg Hg Hg Hg Hg Hg Hg Hg O O O O O O O O Hg Hg Hg Hg Hg Hg O O O O Hg Hg Hg Hg O O O O Hg Hg mercury oxide decomposes Thermal decomposition of metal oxides Heat  + Mercury Oxide Mercury oxygen

7. Exothermic reactions give out heat (gets hot). Endothermic reactions take in heat (gets cold). Many chemical reactions need some energy to get them started (activation energy) but then the majority of chemical reactions are exothermic. Ex = out (as in exit) En = in (as in entrance) Shuttle fuel burning- highly exothermic Exothermic and endothermic reactions

8. It is hard to think of examples of endothermic reactions but there are lots of exothermic ones that occur in the laboratory and in everyday life. List 8 exothermic reactions. Exothermic and endothermic reactions Burning wood on a fire Burning petrol in a car Burning butane in a cigarette lighter Burning gas in a gas hob Reacting an acid and alkali together Burning magnesium Rotting compost etc etc

9. These are reactions where two metals are competing to be combined with a non-metal. The more reactive metal wins the competition and becomes part of a compound. The less reactive metal is displaced and so is present as the metal at the end of the reaction. Potassium sodium calcium magnesium aluminium zinc iron copper silver gold Increasing reactivity Displacement reactions: metals A more reactive metal (higher in the reactivity series) will displace a less reactive metal from its compound.

10. Copper is quite low in the activity series. Several metals will displace it from its compounds. magnesium magnesium sulphate solution copper sulphate solution copper metal more reactive less reactive Displacement reactions: metals K Na Ca Mg Al Zn Fe Cu Ag Au Copper sulphate + Magnesium sulphate  + Copper Magnesium Magnesium wins the competition. Copper is displaced.

11. Here are some actual photos. The colour changes from blue to red/black as copper metal is displaced. Displacement reactions: metals more reactive less reactive photograph at start of reaction photograph at end of reaction K Na Ca Mg Al Zn Fe Cu Ag Au Copper sulphate + Magnesium sulphate  + Copper Magnesium Magnesium wins the competition. Copper is displaced

12. The thermit reaction takes place between aluminium and iron oxide. It is so exothermic that molten iron is produced and the reaction is used to repair broken railway tracks. Displacement reactions: metals more reactive less reactive magnesium fuse iron oxide + aluminium powder K Na Ca Mg Al Zn Fe Cu Ag Au Iron Oxide + Aluminium Oxide  + Iron Aluminium Aluminium wins the competition. Iron is displaced and melts at the high temperatures produced.

13. Here is a photo of the thermit reaction being carried out in a laboratory. Displacement reactions: metals magnesium fuse iron oxide + aluminium powder

14. Predict which mixtures will result in a reaction. Yes Yes No No No No Yes No No Yes Yes Yes

15. These are displacement reactions where two halogens are competing to be combined with a metal. It is the more reactive halogen that will win and become part of a compound. The less reactive halogen remains (or becomes) the element. Fluorine Chlorine Bromine Iodine Displacement reactions: halogens Increasing reactivity • We can often tell which halogen is present from the colour of the solution.

16. For example, if chlorine solution is added to sodium bromide. chlorine solution sodium chloride solution sodium bromide solution bromine more reactive less reactive Displacement reactions: halogens F Cl Br I At Sodium Bromide + Sodium Chloride  + Bromine Chlorine Chlorine wins the competition. Bromine (red) is displaced.

17. The compounds of the halogens with Group 1 metals are all colourless. Predict what colour these will be after mixing. Br2 I2 I2 Br2 I2 I2

18. When writing equations for halogen displacement reactions you must remember that – when in the form of the element – halogens exist in pairs. For chlorine and sodium bromide: Chlorine + Displacement reactions: halogens F Cl Br I At Sodiumbromide sodium chloride + bromine  Cl2(aq) + 2NaBr(aq)  2NaCl(aq) + Br2(aq) Solution goes yellow/brown as bromine is produced. Cl More reactive Br Less reactive

19. If no reaction - not write “no reaction.” Where there is a reaction write the names of the products and then write a chemical equation underneath. Predict whether or not a chemical reaction will occur. 1) iodine + sodium bromide solution  2) bromine + sodium chloride solution  3) chlorine + sodium iodide solution  + iodine sodium chloride F Cl Br I At No reaction No reaction Cl2(g) + 2NaI(aq)  2NaCl(aq) + I2(aq)

20. Most chemical reactions are considered irreversible in that the new products are not readily changed back into reactants. For example, once you have reacted magnesium with hydrochloric acid it is very hard to get the magnesium back. Reversible and irreversible reaction • In the equations for irreversible reactions reactants and products are joined by a “one-way arrow.” magnesium + hydrochloric  magnesium + hydrogen acidchloride

21. Although most chemical reactions are difficult to reverse it is possible to find reactions ranging from irreversible through to the fully reversible. One of the best known reversible processes is heating copper sulphate. Note the double arrow symbol in the chemical equation Reversible reactions Heat anhydrous copper sulphate steam hydrated copper sulphate CuSO4.5H20 CuSO4 + 5H2O these decompose these combine

22. There are some reactions in which both the “forward and backward” reactions occur to a substantial extent under the same conditions. These lead to equilibrium mixtures of reactants and products. One of the most important of these reactions occurs in the Haber Process. N2(g) + 3H2(g) 2 NH3(g) Equilibrium reactions However long you leave the reaction going you still get a mixture of nitrogen, hydrogen and ammonia.

23. There are some simple rules that can be used to move the position of an equilibrium towards reactants or products: What conditions will favour formation of more ammonia? 3H2(g) + N2 (g)  2NH3 (g) (exothermic) Getting more product at equilibrium • Exothermic reactions give more product at lower temperatures. (Endothermic – the opposite) • Increasing the pressure in gas reactions favours whichever side of the chemical equation has least gas molecules. Low temperature High pressure

24. A precipitation reaction is any reaction that produces an insoluble compound when two aqueous solutions are mixed. It is impossible to predict whether or not we will get precipitation reactions unless we know something about the physical states (especially solubility) of the various reactants and products. Precipitation reactions • Here are the symbols that we use in chemical equations to say what the physical state is: • (s) solid • (l) liquid • (g) gas • (aq) aqueous (dissolved in water)

25. A precipitation reaction that is often used to measure reaction rates occurs between sodium thiosulphate and hydrochloric acid. Precipitation reactions – first example Sulphur is insoluble and precipitates. This makes the solution go cloudy. Both reactants are colourless and dissolved (aq) Sodium + hydrochloric  sodium + sulphur + water + sulphur thiosulphate acid chloride dioxide aqueous aqueous aqueous solid liquid gas solid

26. Most metal hydroxides (except sodium, potassium and calcium) are insoluble. Reactions leading to their formation give precipitates. Precipitation reactions – second example Copper hydroxide is insoluble and precipitates. A pale blue solid settles at the bottom of the test tube. Both reactants are dissolved (aq). Copper sulphate is blue. Copper + ammonium  copper + ammonium sulphate hydroxide hydroxide sulphate solid aqueous aqueous solid aqueous

27. Another metal hydroxide that precipitates is iron(III) hydroxide. Like many transition metals its compounds are coloured. Precipitation reactions – third example Iron hydroxide is insoluble and precipitates. A deep brown solid settles at the bottom of the test tube. Both reactants are dissolved (aq) (iron chloride is yellow). Iron + sodium  iron + sodium chloride hydroxide hydroxide chloride aqueous aqueous solid aqueous solid

28. To work out whether a precipitate will be formed we need to know the solubility of the compounds that may be formed. Here are a few general guidelines: Precipitation and solubility

29. To work out whether a precipitate will be formed when many ionic compounds react there are four stages: Precipitation and solubility 1 Write down the names of the reactants. Sodium chloride & lead nitrate Na+ Cl- Pb2+ NO3- 2 Write down the ions in the reactants. (Ignore numbers) 3 Swap over the + and – ions. Pb2+ Cl- Na+ NO3- 4 Are the products going to be soluble or insoluble? Lead chloride is insoluble so there will be a precipitate Sodium + lead  lead + sodium chloride nitrate chloride nitrate aqueous aqueous solid aqueous solid

30. Will there be a precipitate if I mix sodium sulphate and magnesium nitrate? 1 Write down the names of the reactants. 2 Write down the ions in the reactants. 3 Swap over the + and – ions. 4 Are the products going to be soluble or insoluble? Sodium nitrate & Magnesium sulphate Na+ SO42- Mg2+ NO3- Mg2+ SO42- Na+ NO3- Both the products are soluble there will be no precipitate. Sodium + magnesium  magnesium + sodium sulphate nitrate sulphate nitrate aqueous aqueous aqueous aqueous

31. Will there be a precipitate if I mix sodium sulphate and barium nitrate? 1 Write down the names of the reactants. 2 Write down the ions in the reactants. 3 Swap over the + and – ions. 4 Are the products going to be soluble or insoluble? Sodium sulphate & barium nitrate Na+ SO42- Ba2+ NO3- Ba2+ SO42- Na+ NO3- Barium sulphate is insoluble so there will be a precipitate. Sodium + barium  barium + sodium sulphate nitrate sulphate nitrate aqueous aqueous solid aqueous solid

32. Separating Precipitates – reminder!

33. Acids are substances that: Turn litmus red. Turn universal indicator yellow, orange or red. Have a pH below 7. Form solutions containing H+ ions. Bases are substances that: Turn litmus blue. Turn universal indicator dark green, blue or purple. React with the H+ ions in acids. Are called alkalis if they dissolve in water. 1 2 3 4 5 6 7 8 9 10 11 12 13 14 Neutralisation reactions Increasingly acid Increasingly alkali

34. Common Acids are Neutralisation reactions: acids H2SO4 strong HCl strong HNO3 strong CH3COOH weak • Salts Sulphuric acid Nitric acid Hydrochloric acid Sulphates Nitrates Chlorides

35. Common alkalis are Neutralisation reactions: bases NaOH strong strong KOH strong Ca(OH)2 NH4OH weak • Common bases (neutralise acids but don’t dissolve) are OH- water + a salt O2- water + a salt CO32- water + a salt + CO2

36. A neutralisation reaction is where an acid reacts with a base to produce a neutral solution of a saltandwater. 1 2 3 4 5 6 7 8 9 10 11 12 13 14 sodium hydroxide pH 14 hydrochloric acid pH 1 sodium chloride pH 7 Increasingly acid Increasingly alkali Neutralisation reactions: acid + base neutralisation

37. To name the salt formed in a neutralisation: Neutralisation - naming salts 1 The first part of the name of the salt comes from the first name of the base So Ammonium hydroxide gives ammonium ………… Magnesium oxide gives magnesium …………... 2 The acid gives the last part of the name of the salt. So Sulphuric acid make sulphates Nitric acid makes nitrates Hydrochloric acid makes chlorides Eg. Sodium hydroxide + nitric acid forms: Calcium carbonate + sulphuric acid forms: Sodium nitrate calcium sulphate

38. Name the salt formed in these neutralisations: +  Calcium chloride Magnesium nitrate Calcium sulphate Aluminium nitrate Potassium sulphate

39. Each OH- ion reacts with one H+ ion. Neutralisation reactions: hydroxides Reaction with hydroxides: H+ + OH- H2O Eg. Potassium +hydrochloric  water + potassium hydroxide acid chloride KOH+ HCl H2O+ KCl Eg. Calcium + sulphuric  water + calcium hydroxide acid sulphate Ca(OH)2+ H2SO4 2H2O+ CaSO4

40. Neutralisation reactions usually lead to water being formed. Neutralisation Reactions: oxides Reaction with oxides: 2H+ + O2- H2O Eg. Calcium + hydrochloric  water + calcium oxide acid chloride CaO+ 2HCl  H2O+ CaCl2 Eg. Sodium + sulphuric  water + sodium oxide acid sulphate Na2O+H2SO4 H2O+ Na2SO4

41. Each carbonate ion provides one oxygen to join with two H+ ions. At the same time carbon dioxide is released. Neutralisation Reactions: carbonates Carbonates: 2H+ + CO32- H2O + CO2 Eg. Potassium + hydrochloric  water + carbon + potassium carbonate acid dioxide chloride K2CO3+ 2HClH2O + CO2 + 2KCl Eg. calcium + nitric  water + carbon + calcium carbonate acid dioxide nitrate CaCO3+ 2HNO3 H2O+ CO2 +Ca(NO3)2

42. Complete the word equation Neutralisation equations Eg. Potassium + hydrochloric  + hydroxide acid Potassium chloride water Replace the words with the correct formula Eg. KOH + HCl  + H2O KCl Check that it balances (same number of each type of atom each side).  Eg. KOH + HCl  H2O + KCl

43. Complete the word equation Neutralisation equations Eg. Magnesium + nitric  + oxide acid Magnesium nitrate water Replace the words with the correct formula Eg. MgO + HNO3 + H2O Mg(NO3)2 Check that it balances (Same number of each type of atom each side.  Eg. MgO + HNO3 H2O + Mg(NO3)2 2 2

44. Write balanced equations going through the same stages as the previous examples. word equation formulae balance a) sodium hydroxide + hydrochloric acid  b) magnesium oxide + hydrochloric acid  c) sodium hydroxide + sulphuric acid  d) ammonium hydroxide + hydrochloric acid  e) calcium hydroxide + nitric acid 

45. Insoluble salts can be separated by filtering. Soluble salts are obtained by evaporating. vapour evaporating basin gauze tripod bunsen burner heat-proof mat • Put these in the correct order. • Check the pH frequently by testing drops of the solution. • Add the acid slowly to the alkali. • When neutral pour into the evaporating basin. • Put on safety specs. • Allow to cool • Heat. D B A C F E

46. Redox is a short way of saying: Oxidation meant adding oxygen to a substance. Reduction meant taking oxygen away. Extracting iron from iron oxide in the blast furnace is reduction. Rusting (iron becoming iron oxide) is an example of oxidation. Redox Reactions Reduction and oxidation • Early on in chemistry these words had very straightforward meanings.

47. Many redox reactions involve metals and their oxides. Whenever metals react with oxygen they form ionic compounds and the metal loses electrons to form positively charged ions. Eg. When magnesium burns to form magnesium oxide magnesium atoms (no charge) become magnesium ions (2+ charge) by losing 2 electrons to oxygen atoms. Mg Mg2+ O2- Redox reactions: oxidation and ions O 2 e- to give Oxidation involves loss of electrons.

48. Think about what has happened to the magnesium when it reacts with oxygen. It has been oxidised. It has lost electrons by changing from Mg  Mg2+ Magnesium can also lose electrons to things other than oxygen (e.g. to chlorine or sulphur) and since these also involve Mg  Mg2+ these too must be oxidation. O2- O2- Mg2+ Mg2+ O O Mg S S Mg2+ Mg2+ S2- S2- Cl Cl Mg2+ Mg2+ Cl- Cl- Cl- Cl- Redox reactions: electron loss Mg Oxidation is the loss of electrons.

49. Exactly the same reasoning applies to reduction. Reduction can be the removal of oxygen (e.g. from iron oxide to form iron or from aluminium oxide in the electrolysis to extract aluminium.) When this happens the metal gets back its electrons. Aluminium has been reduced. Aluminium has gained electrons O2- Al O Al3+ 1½ O2- O Al Al3+ Oxygen removed O2- Redox reactions: electron gain Reduction is the gain of electrons.

50. An easy way of remembering this is “Oil Rig”! Redox Reactions: oil rig of electrons O oxidation I is L loss R reduction I is G gain