Chemical Reactions Boardworks Ltd 2001 Oxidation and reduction

Chemical Reactions © Boardworks Ltd 2001

Oxidation and reduction Thermal decomposition Types of chemical change neutralisation precipitation Reversible reactions Exothermic and endothermic Displacement Reactions: • metals Displacement Reactions: • non-metals © Boardworks Ltd 2001

Thermal Decomposition Potassium carbonate is not thermally decomposed. Calcium carbonate decomposes on strong heating Silver carbonate decomposes on gentle heating Gets harder • 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. Eg. © Boardworks Ltd 2001

• Generally, the more reactive the metal is, the more difficult it is to decompose it’s compounds. • Fill in the last column: easy, medium or hard Compound Mercury oxide Sodium oxide Iron oxide Silver oxide Zinc oxide How easy to decompose easy hard medium Potassium sodium calcium magnesium aluminium zinc iron copper mercury silver gold Increasing reactivity Thermal Decomposition easy medium © Boardworks Ltd 2001

Thermal Decomposition of carbonates limestone • 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). Calcium Carbonate Calcium oxide + 1500°C Waste air and carbon dioxide Hot air Carbon dioxide Calcium oxide (lime) © Boardworks Ltd 2001

Thermal Decomposition of metal oxides • Most metal oxides are thermally stable. (I. e. Do not decompose when heated. ) • Oxides of the least reactive metals are thermally decomposed. • E. g. silver oxide begins to break up at about 160 o. C and mercury oxide decomposes when heated strongly. Mercury Oxide Mercury + Mercury metal and oxygen formed Heat Hg O Hg O O Hg Hg Hg O O Hg Mercury oxide decomposes oxygen © Boardworks Ltd 2001

Exothermic and Endothermic Reactions • Ex = out (as in exit. ) • En = in (as in entrance. ) • Exothermic reactions give out heat (get hot. ) • Endothermic reactions take in heat (get cold. ) • Many chemical reactions need some energy to get them started (activation energy) but then the majority of chemical reactions are exothermic. Shuttle fuel burning. Highly exothermic © Boardworks Ltd 2001

Exothermic and Endothermic Reactions • 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. Some examples of exothermic 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 © Boardworks Ltd 2001

• 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 © Boardworks Ltd 2001

Displacement Reactions: Metals • Copper is quite low in the activity series. • Several metals will displace it from its compounds. magnesium K Na Ca Mg Al Zn Fe Cu Ag Au Magnesium more reactive Magnesium sulphate solution copper sulphate solution + Copper sulphate less reactive Copper metal Magnesium sulphate + Copper Magnesium wins the competition. Copper is displaced © Boardworks Ltd 2001

Displacement Reactions: Metals • Here are some actual photos. • The colour changes from blue to red/black as copper metal is displaced. K Na Ca Mg Al Zn Fe Cu Ag Au Photograph at start of reaction Magnesium + more reactive Photograph at end of reaction Copper sulphate less reactive Magnesium sulphate + Copper Magnesium wins the competition. Copper is displaced © Boardworks Ltd 2001

Displacement Reactions: Metals • 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 K Na broken railway tracks. Ca Mg Al Zn Fe Cu Ag Au Aluminium more reactive + Iron Oxide less reactive Aluminium Oxide Magnesium fuse iron oxide + aluminium powder + Iron Aluminium wins the competition. Iron is displaced and melts at the high temperatures produced. © Boardworks Ltd 2001

Displacement Reactions: Metals • Here is a photo of thermit reaction being carried out in a laboratory Magnesium fuse iron oxide + aluminium powder © Boardworks Ltd 2001

Activity • Predict which mixtures will result in a reaction Metal Solution Iron chloride Magnesium nitrate Zinc nitrate Copper sulphate Magnesium Yes No No Yes Yes Zinc Copper Yes No No Yes © Boardworks Ltd 2001

Displacement Reactions: Halogens • Fluorine Chlorine Bromine Iodine Increasing reactivity • 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. We can often tell which halogen is present from the colour of the solution. © Boardworks Ltd 2001

Displacement Reactions: Halogens • Eg. If chlorine solution is added to sodium bromide chlorine solution F Cl Br I At Sodium chloride solution Sodium bromide solution Bromine Chlorine more reactive + Sodium Bromide less reactive Sodium Chloride + Bromine Chlorine wins the competition. Bromine (red) is displaced © Boardworks Ltd 2001

Activity • Predict what colour these will be after mixing. • (The compounds of the halogens with Group 1 metals are all colorless. ) Halogen Halide Chlorine solution Potassium chloride Bromine solution Iodine Solution Br 2 Potassium bromide Br 2 Potassium Iodide I 2 I 2 © Boardworks Ltd 2001

Displacement Reactions: Halogens • When writing equations for halogen displacement reactions you must remember that – when in the form of the element – halogens exist in pairs. • Eg. For chlorine and sodium bromide: Chlorine + Sodium sodium bromide chloride F Cl Br I At + bromine Cl 2(aq) + 2 Na. Br(aq) 2 Na. Cl(aq) + Br 2(aq) Cl More reactive Br Less reactive Solution goes yellow/brown as bromine is produced. © Boardworks Ltd 2001

Activity • Predict whether or not a chemical reaction will occur, • 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. . 1) iodine + sodium bromide solution No reaction 2) bromine + sodium chloride solution No reaction 3) chlorine + sodium iodide solution Sodium chloride Cl 2(g) + 2 Na. I(aq) F Cl Br I At + iodine 2 Na. Cl(aq) + I 2(aq) © Boardworks Ltd 2001

Reversible and Irreversible Reaction • Most Chemical reactions are considered irreversible in that the new products are not readily changed back into reactants. Eg. Once you have reacted magnesium with hydrochloric acid it is very hard to get the magnesium back. • In the equations for irreversible reactions reactants and products are joined by a “oneway arrow. ” magnesium + hydrochloric magnesium + hydrogen acid chloride © Boardworks Ltd 2001

Reversible Reactions • 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 Heat Hydrated copper sulphate Cu. SO 4. 5 H 20 These decompose anhydrous copper sulphate Cu. SO 4 + steam 5 H 2 O These combine © Boardworks Ltd 2001

Equilibrium Reactions. • 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. N 2(g) + 3 H 2(g) 2 NH 3(g) However long you leave the reaction going you still get a mixture of nitrogen, hydrogen and ammonia. © Boardworks Ltd 2001

Getting more product at equilibrium. • There are some simple rules that can be used to move the position of an equilibrium towards reactants or products: 1. Exothermic reactions give more product at lower temperatures. (Endothermic – the opposite) 2. Increasing the pressure in gas reactions favours whichever side of the chemical equation has least gas molecules. What conditions will favour formation of more ammonia? 3 H 2(g) + N 2 (g) Low temperature 2 NH 3 (g) (exothermic) High pressure © Boardworks Ltd 2001

Precipitation Reactions • 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. • Here are the symbols that we use in chemical equations to say what the physical state is: –(s) –(l) –(g) –(aq) solid liquid gas aqueous (dissolved in water) © Boardworks Ltd 2001

Precipitation Reactions – 1 st example • A precipitation reaction that is often used to measure reaction rates occurs between sodium thiosulphate and hydrochloric acid. Sodium + thiosulphate hydrochloric acid aqueous Both reactants are colourless and dissolved (aq) sodium + sulphur + water + sulphur chloride dioxide solid aqueous solid liquid gas Sulphur is insoluble and precipitates. This makes the solution go cloudy. © Boardworks Ltd 2001

Precipitation Reactions – 2 nd example • Most metal hydroxides (except sodium, potassium and calcium) are insoluble. Reactions leading to their formation give precipitates. Copper + sulphate aqueous ammonium hydroxide copper hydroxide aqueous solid Both reactants are dissolved (aq). Copper sulphate is blue. + ammonium sulphate aqueous Copper hydroxide is insoluble and precipitates. A pale blue solid settles at the bottom of the test tube. © Boardworks Ltd 2001

Precipitation Reactions – 3 rd example • Another metal hydroxide that precepitates is iron(III) hydroxide. Like many transition metals its compounds are coloured. Iron + chloride sodium hydroxide aqueous Both reactants are dissolved (aq) (Iron chloride is yellow. ) iron hydroxide solid + sodium chloride aqueous Iron hydroxide is insoluble and precipitates. A deep brown solid settles at the bottom of the test tube. © Boardworks Ltd 2001

Precipitation and solubility • To work out whether a precipitate will be formed we need to know the solubility of the compounds that may be formed. Here a few general guidelines: Soluble Insoluble All sodium, potassium and ammonium salts All nitrates Most chlorides, bromides and iodides. (halides) Silver and lead halides Most sulphates Lead, barium and calcium sulphates Sodium, potassium and ammonium carbonates Most carbonates Sodium, potassium, ammonium and calcium hydroxide Most hydroxides © Boardworks Ltd 2001

Precipitation and solubility • To work out whether a precipitate will be formed when many ionic compounds react there are four stages: 1. Write down the names of the reactants. Sodium chloride & lead nitrate 2 Write down the ions in the reactants. (Ignore numbers) Na+ Cl- Pb 2+ NO 3 - 3 Swap over the + and – ions. Pb 2+ Cl- Na+ NO 3 - 4 Are the products going to be soluble or insoluble? Sodium + chloride aqueous lead nitrate aqueous Lead chloride is insoluble so there will be a precipitate lead chloride solid + sodium nitrate aqueous © Boardworks Ltd 2001

Activity • 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. Sodium nitrate & Na+ SO 42 - Magnesium sulphate Mg 2+ NO 3 - 3 Swap over the + and – ions. Mg 2+ SO 42 - 4 Are the products going to be soluble or insoluble? Both the products are soluble there will be no precipitate Sodium + sulphate aqueous magnesium nitrate aqueous magnesium sulphate aqueous Na+ NO 3 - + sodium nitrate aqueous © Boardworks Ltd 2001

Activity • 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. Na+ SO 42 - 3 Swap over the + and – ions. 4 Are the products going to be soluble or insoluble? Sodium + sulphate aqueous barium nitrate aqueous Sodium sulphate & barium nitrate Ba 2+ NO 3 - Ba 2+ SO 42 - Barium sulphate is insoluble so there will be a precipitate barium sulphate solid Na+ NO 3 - + sodium nitrate aqueous © Boardworks Ltd 2001

Activity Separating Precipitates – reminder! © Boardworks Ltd 2001

Neutralisation Reactions • Acids are substances that: • Turn litmus red. • Turn universal indicator yellow, orange or red. • Have a p. H 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 Increasingly acid 6 7 8 9 10 11 12 13 14 Increasingly alkali © Boardworks Ltd 2001

Neutralisation Reactions: Acids • Common Acids are Name of acid Formula Sulphuric H 2 SO 4 strong Hydrochloric HCl strong Nitric HNO 3 strong CH 3 COOH weak Ethanoic (vinegar) Strong or Weak? • Salts Sulphuric acid Nitric acid Hydrochloric acid Sulphates Nitrates Chlorides © Boardworks Ltd 2001

Neutralisation Reactions: Bases • Common Alkalis are Name of alkali Formula Strong or Weak? Sodium Hydroxide Na. OH strong Potassium Hydroxide KOH strong Calcium Hydroxide Ca(OH)2 strong Ammonium Hydroxide NH 4 OH weak • Common Bases (neutralise acids but don’t dissolve) are Type of compound Contain React with acids to give Metal Hydroxides OH- water + a salt Metal Oxides O 2 - water + a salt Metal Carbonates CO 32 - water + a salt + CO 2 © Boardworks Ltd 2001

Neutralisation Reactions: acid + base • A neutralisation reaction is where an acid reacts with a base to produce a neutral solution of a salt and water Sodium hydroxide p. H 14 neutralisation Hydrochloric acid p. H 1 1 2 3 Sodium chloride p. H 7 4 5 Increasingly acid 6 7 8 9 10 11 12 13 14 Increasingly alkali © Boardworks Ltd 2001

Neutralisation Naming Salts • To name the salt formed in a neutralisation: 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 © Boardworks Ltd 2001

Activity • Name the salt formed in these neutralisations: + Base Acid Calcium hydroxide Hydrochloric acid Magnesium oxide Nitric acid Calcium carbonate Sulphuric acid Aluminium hydroxide Nitric acid Potassium hydroxide Sulphuric acid Salt? Calcium chloride Magnesium nitrate Calcium sulphate Aluminium nitrate Potassium sulphate © Boardworks Ltd 2001

Neutralisation Reactions: hydroxides • Each OH- ion reacts with one H+ ion Reaction with hydroxides: H+ + OH- H 2 O Eg. Potassium +hydrochloric water + hydroxide acid KOH + HCl H 2 O Eg. Calcium + sulphuric water + hydroxide acid Ca(OH) OH 2 + H 2 SO 4 + potassium chloride KCl calcium sulphate 2 H 2 O + Ca. SO 4 © Boardworks Ltd 2001

Neutralisation Reactions: oxides • Neutralisation reactions usually lead to water being formed. Reaction with oxides: 2 H+ + O 2 - H 2 O Eg. Calcium + hydrochloric water + oxide acid Ca. O + 2 HCl 2 H H 2 O Eg. Sodium + sulphuric water + oxide acid Na 2 O + H 2 SO 4 H 2 O calcium chloride + Ca. Cl 2 sodium sulphate + Na 2 SO 4 © Boardworks Ltd 2001

Neutralisation Reactions: carbonates • Each carbonate ion provides one oxygen to join with two H+ ions. At the same time carbon dioxide is released Carbonates: 2 H+ + CO 32 - H 2 O + CO 2 Eg. Potassium + hydrochloric water + carbon + potassium carbonate acid dioxide chloride K 2 CO 3 Eg. calcium + carbonate Ca. CO 3 + 2 HCl 2 H nitric acid + 2 HNO 2 H 3 H 2 O + CO 2 + 2 KCl water + carbon + calcium dioxide nitrate H 2 O + CO 2 +Ca(NO 3)2 © Boardworks Ltd 2001

Neutralisation Equations • Complete the word equation Eg. Potassium + hydrochloric acid water + hydroxide Potassium chloride • Replace the words with the correct formula KOH Eg. + HCl H 2 O + KCl • Check that it balances (same number of each type of atom each side). Eg. KOH 1*K + HCl Reactants 1*O 2*H H 2 O + KCl 1*Cl 2*H Products 1*O 1*K 1*Cl © Boardworks Ltd 2001

Neutralisation Equations • Complete the word equation Eg. Magnesium + nitric acid + water oxide Magnesium nitrate • Replace the words with the correct formula Eg. Mg. O + HNO 3 H 2 O + Mg(NO 3)2 • Check that it balances (Same number of each type of atom each side. Eg. Mg. O 1*Mg + 2 HNO 3 Reactants 1*O 1*H 2 H 2 O + 1*NO 3 2*H Mg(NO 3)2 Products 1*O 1*Mg 2*NO 3 © Boardworks Ltd 2001

Activity • Write balanced equations going through the same stages as the previous examples 1. word equation, 2. formulae, 3. 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 © Boardworks Ltd 2001

Activity • Insoluble salts can be separated by filtering. • Soluble salts are obtained by evaporating. Put these in the correct order D A. Check the p. H frequently by testing drops of the solution. B. Add the acid slowly to the alkali. C. When neutral pour into the evaporating basin. D. Put on safety specs. B A C F E vapour gauze Evaporating basin tripod Heat-proof mat Bunsen burner E. Allow to cool F. Heat. © Boardworks Ltd 2001

Redox Reactions: • Redox is a short way of saying: • Early on in chemistry these words had very straightforward meanings Oxidation meant adding oxygen to a substance. Rusting (iron becoming iron oxide) is an example of oxidation Reduction and oxidation. Reduction meant taking oxygen away. Extracting iron from iron oxide in the blast furnace is reduction © Boardworks Ltd 2001

Redox Reactions: Oxidation and ions • 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 2 e- O to give Mg 2+ O 2 - Oxidation involves loss of electrons. © Boardworks Ltd 2001

Redox Reactions: Electron loss. • 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 2+ • Magnesium can also lose electrons to things other than oxygen (e. g. to chlorine or sulphur) and since these also involve Mg 2+ these too must be oxidation. O 2 - O Mg S Mg 2+ S 2 - Cl Oxidation is loss of electrons. Cl- Mg 2+ Cl© Boardworks Ltd 2001

Redox Reactions: Electron gain • 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 O 2 Al Al 3+ O 2 - Al 1½ O O Reduction is gain of electrons. Oxygen removed © Boardworks Ltd 2001

Redox Reactions: Oil Rig • An easy way of remembering this is “Oil Rig” Rig O oxidation I is L loss of electrons R Reduction I is G gain © Boardworks Ltd 2001

Redox Reactions: Two for one! • Whenever something is oxidised, something else is reduced. • This should be obvious if we use the oil rig definition. • If something loses electrons – then something else must have gained them. • E. g. When burning magnesium: – Magnesium loses electrons (Mg Mg 2+ …. . oxidation) – Oxygen gains electrons (O O 2 - ……. reduction) • The overall reaction is both Reduction and Red Oxidation = Redox Ox © Boardworks Ltd 2001

Activity • Say whether the substance in red type is oxidised or reduced. Calcium + oxygen calcium oxide oxidised Zinc oxide + hydrogen zinc + water reduced Copper chloride copper + chlorine reduced Aluminium + iron oxide iron + aluminium oxide oxidised © Boardworks Ltd 2001

Activity • If the first substance is oxidised, what has been reduced or vice versa. (Use whichever definition of oxidation and reduction seems easier to apply. ). Calcium + oxygen calcium oxide oxidised Oxygen is reduced. Each oxygen atom gains 2 e- Zinc oxide + hydrogen zinc + water reduced Hydrogen is oxidised. It gains oxygen Copper chloride reduced copper + chlorine Chlorine is Oxidised. It gains an electron Cl- ½Cl 2 Aluminium + iron oxide iron + aluminium oxide oxidised Iron is reduced. It loses oxygen © Boardworks Ltd 2001

Activity • Across: 5 tells us whether acid or alkali 11 reaction of an acid with a base • 1 2 3 4 6 7 8 9 10 Down a solid forms in a solution loss of electrons competition reaction gives solutions containing H+ ions to break down into smaller particles removal of oxygen state of balance soluble base ionic compound formed in neutralisations © Boardworks Ltd 2001

Activity • Match them up Thermal decomposition Endothermic Metal displacement Reversible reaction Dehydrating copper sulphate A solid forms within a solution A salt and water is formed Alkali Precipitation Neutralisation Oxidation Reduction Soluble base Equilibrium Reaction in a state of balance Thermit reaction Removal of oxygen Breaking up with heat Takes in energy – gets cold Loss of electrons © Boardworks Ltd 2001

When heated the orange powder erupted like a volcano producing a huge pile of green powder that had less mass than the orange material. Is this? 1. 2. 3. 4. Neutralisation Thermal decomposition Displacement Precipitation © Boardworks Ltd 2001

When the two colourless solutions mixed a yellow solid formed which sank to the bottom of the test tube Is this? 1. 2. 3. 4. Neutralisation Thermal decomposition Displacement Precipitation © Boardworks Ltd 2001

When the copper was placed in the silver nitrate solution snow like crystals of silver seemed to grow out from the copper. Is this? 1. 2. 3. 4. Equilibrium Thermal decomposition Displacement Precipitation © Boardworks Ltd 2001

When the washing soda was added to the lemon juice it fizzed and the p. H rose towards 7. Is this? 1. 2. 3. 4. Neutralisation Thermal decomposition Displacement Oxidation © Boardworks Ltd 2001

Which of the oxides shown will thermally decompose most easily? 1. 2. 3. 4. Mercury oxide Potassium oxide Iron oxide Gold oxide © Boardworks Ltd 2001

Which of the salts below might be formed when nitric acid neutralises a metal hydroxide? 1. 2. 3. 4. Potassium hydroxide Potassium nitrate Ammonium nitrate Calcium sulphate © Boardworks Ltd 2001

Which of the mixtures below will result in a metal displacement reaction? 1. Potassium oxide and gold 2. Magnesium and sodium nitrate 3. Copper and gold nitrate 4. Aluminium and calcium sulphate © Boardworks Ltd 2001

Which of the mixtures below will result in a nonmetal displacement reaction? 1. Potassium chloride and iodine 2. Potassium bromide and iodine 3. Potassium fluoride and chlorine 4. Potassium iodide and chlorine © Boardworks Ltd 2001

Which of the elements in red (below) is oxidised in the reaction? (Oil Rig!)? 1. Ca + 2. 2 Li + 3. 2 Al + 4. HNO 3 Cu. O Cu Ca. O + Cu 2 HCl 2 Li. Cl + H 2 Fe 2 O 3 Al 2 O 3 + 2 Fe + Cu. O Cu Cu. NO 3 + H 2 O © Boardworks Ltd 2001

Which compound can you be sure is soluble in water? 1. 2. 3. 4. Manganese nitrate Osmium iodide Thallium chloride Palladium sulphate © Boardworks Ltd 2001