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Chapter 12

Chapter 12. Reactivity of metals. Criteria for comparing the reactivity of metals. The temperature at which the reaction starts. The more reactive the metal is, the lower the temperature required. The rate/speed of the reaction The more reactive the metal, the faster is the reaction rate.

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Chapter 12

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  1. Chapter 12 Reactivity of metals

  2. Criteria for comparing the reactivity of metals • The temperature at which the reaction starts. • The more reactive the metal is, the lower the temperature required. • The rate/speed of the reaction • The more reactive the metal, the faster is the reaction rate. • The amount of heat given out during reaction. • The more reactive the metal, the more heat will be given out during reaction.

  3. Chemical reactions for determining the reactivity series • Reaction with air / oxygen • Reaction with water • Reaction with dilute acid

  4. Reaction with air (exposing to air) • Metals are usually dull in colour after exposing to air for a long time. • Ready to react with oxygen in air to form an oxide layer. (i.e., tarnish in air) • Shown shiny surface only when freshly cut or polished / scratched. • Reactive metals (such as sodium and potassium) are stored under paraffin oil.

  5. Reaction with air (Heating in air) • Some metals / their metal compounds burn with a characteristic coloured flame.

  6. Products formed in the reaction • Metal oxides are formed. • Metal + oxygen → Metal oxide • Mg + O2 → ________________ • Oxides of transition metals are usually coloured. • No apparent reaction for silver, gold and platinum (unreactive metals).

  7. Reaction with water (at room temperature) • Reactive metals such as potassium, sodium and calcium react with cold water to form metal hydroxide and hydrogen. • Metal + water → metal hydroxide + hydrogen • Na(s) + H2O(l) → • Ca(s) + H2O(l) →

  8. Reaction of sodium with cold water • Briefly describe the reaction of sodium with cold water? • The small piece of sodium melts into a silvery ball. It moves across the surface of water with a hissing sound. If its movement is stopped, it burns with a golden yellow flame.

  9. Reaction of calcium with cold water • Briefly describe the reaction of calcium with cold water? • Calcium metal sinks to the bottom of the beaker. Why ? • It reacts moderately with cold water giving out colourless babbles of hydrogen. • A white suspension of calcium hydroxide is formed as calcium hydroxide is slightly soluble in water.

  10. Reaction with hot steam • Less reactive metals (such as magnesium, zinc and iron) have little or no reaction with cold water. • Readily react with hot steam to form metal oxide and hydrogen. • Metal + steam  metal oxide + hydrogen • Mg(s) + H2O(g) MgO(s) + H2(g)

  11. Reaction with dilute acids • Dilute acid: hydrochloric acid and sulphuric acid • Metals that are more reactive than copper, react with dilute acids to give hydrogen. • Metal + hydrochloric acid  metal chloride + hydrogen • Metal + sulphuric acid  metal sulphate + hydrogen

  12. What do you observe when magnesium ribbon is added into dilute hydrochloric acid? • Magnesium ribbon dissolves rapidly in dilute acid. Colourless gas bubbles are given out. The tube becomes warm. • It is an exothermic reaction. • Mg(s) + 2HCl(aq) MgCl2(aq) + H2(g) • Zn(s) + H2SO4(aq)  ZnSO4(aq) + H2(g)

  13. Test for hydrogen gas • Put a burning splint near the mouth of the test tube. • A pop sound is heard.

  14. Never add sodium / potassium into dilute acids • Why? • Sodium / potassium (Group I metal) reacts explosively with dilute acids.

  15. Reaction of dilute sulphuric acid with calcium / lead • Colourless gas bubbles are given out at a moderate rate. • But, the reaction stops after a short while. Why? • A layer of insoluble calcium sulphate is formed on the surface of calcium. This insoluble layer prevents the further attack of the acid. • All metal sulphates are soluble in water, except calcium sulphate, barium sulphate and lead(II) sulphate.

  16. Chemical Equations • Formulae of reactants – on the left hand side of the arrow • Formulae of products – on the right hand side of the arrow • ‘+’on the LHS – react with • ‘+’on the RHS – and • ‘’– change to & equal to

  17. Useful information from a balanced equation • The reactants involved. • The products formed. • The physical states of substances involved. • The relative number of particles (atoms, ions, & molecules) of each substance iinvolved.

  18. Rules for writing an equation • Determine the types of reactants involved and the products formed in the reaction. • Write down the correct formulae of reactants on the left hand side of the arrow. • Write down the correct formulae of products on the right hand side of the arrow. • Balance the equation with simple whole numbers such that the total number of each type of atoms are equal on both sides of the arrow. • Put in the physical states for each substance.

  19. Why metals have different reactivity? • Atoms tend to attain stable octet (an inert gas structure) either by gaining or losing electrons; or by sharing electron pairs. • Metal reacts by losing electrons. • Sodium reacts by losing one electron. • Na  Na+ + e- • Non-metal reacts by gaining electrons. • Chlorine reacts by gaining electrons. • Cl2 + 2e- 2Cl-

  20. Why metals have different reactivity? • The relative reactivity of metals depends on the readiness (ease / tendency) of losing electrons. • The relative ease of losing electrons is related to the number of outermost shell electrons and the number of electron shells (i.e., the size of the atoms.)

  21. Relative reactivity of metals across the Periodic Table from left to right • Third Period (from Na to Al) ??? • The reactivity of metals decreases from right to left. (i.e., Na > Mg > Al) • The relative reactivity of metals decreases with increasing group number (increasing number of outermost shell electrons.) • More difficult to remove all the outermost shell electrons.

  22. Relative reactivity of metals down a group in the Periodic Table • Group I metals ??? • The relative reactivity of metals increases down the group as the number of inner shells increases. (K > Na >Li) • The attractive force between the nucleus and the outermost shell electron decreases with increasing atomic size. Thus, the reactivity of metals increases down the group.

  23. Application of reactivity series • Extraction of metals from ores • Thermit reaction (reduction with metals) • Metal displacement reaction • Predicting the stability of metal compounds

  24. Extraction of metals • Extracting metal – getting metal from ores. • What are the metal compounds from mineral ores? • Are they soluble in water? • Insoluble metal oxides, carbonates and sulphides – found in ores • Which metals are found free in nature? • Unreactive metals such as gold and platinum found free (as elements) in nature.

  25. Different methods of extracting metals • Heating metal oxides alone. • Heating metal oxides with carbon (coke) • Electrolysis of hot molten ores

  26. Heating metal oxides alone • Which oxides, magnesium oxide or silver(I) oxide, is more stable to heat? • Why? • The more reactive the metal, the more stable is its compounds. • The less reactive the metal, the less stable is its compounds. • Only fit for metals that are at the bottom of the reactivity series. Why?

  27. Heating silver(I) oxide alone • A colourless gas which relights a glowing splint is given out. • The brown solid turns silvery grey. • 2Ag2O(s) 4Ag(s) + O2(g)

  28. Heating mercury(II) oxide alone • A colourless gas which relights a glowing splint is given out. • The red powder turns silvery. • 2HgO(s) 2Hg(l) + O2(g)

  29. Heating mercury(II) sulphide in air • Reacts with air to form mercury and sulphur dioxide. • HgS(s) + O2(g) Hg(l) + SO2(g)

  30. Redox reaction • Oxidation-reduction reaction • Oxidation and reduction take place at the same time (simultaneously). • Reduction is the removal of oxygen from a substance. • Oxidation is the addition of oxygen to a substance.

  31. Heating metal oxides with coke / carbon • What is reduction? • What is oxidation? • Give examples of oxidation-reduction reaction. • Burning of fuels / candles • Respiration • Rusting • Burning of hydrogen / carbon

  32. Heating lead(II) oxide with carbon • The yellow lead(II) oxide changes to silvery beads of hot molten lead. • 2PbO(s) + C(s)  2Pb(s0 + CO2(G)

  33. Role of carbon • What is the role of carbon? • Carbon is the reducing agent. • What is a reducing agent? • A reducing agent helps to remove oxygen from other substances.

  34. Oxidizing / oxidising agent • What is an oxidizing agent? • An oxidizing agent helps to add oxygen to othe substances. • Name the oxidizing agent in the reaction of lead(II) oxide with carbon. • Lead(II) oxide is the oxidizing agent.

  35. Heating copper(II) oxide with carbon • The black copper(II) oxide turns reddish brown. • 2CuO(s) + C(s)  2Cu(s0 + CO2(G)

  36. Reduction with carbon • With Bunsen flame, carbon can reduce up to lead(II) oxide. (~approx. 1200oC) • In furnace (in factory), (up to 1500oC), carbon can reduce up to zinc oxide.

  37. How to extract lead from lead(II) sulphide (galena)? • Lead(II) sulphide is first heated (roasted) in air. Lead(II) oxide is formed. • 2PbS(s) + 3O2(g) 2PbO(s) + 2SO2(g) • Lead(II) oxide is then heated with carbon. Lead is formed. • 2PbO(s) + C(s)  2Pb(s)+ CO2(g)

  38. Electrolysis of hot molten ores • Reactive metals, such as potassium , sodium, calcium, magnesium and aluminium are extracted from their hot molten ores by electrolysis. • An expensive method. • e.g., aluminium – from the electrolysis of hot molten aluminium oxide.

  39. Year of discovery • The less reactive the metal, the less stable is its compounds and the easier is it to be extracted by Man (the earlier is it to be discovered by Man).

  40. Reaction with more reactive metal • Metal reacts by losing electrons. • Suggest a metal that can be used to extract copper from copper(II) oxide. • Magnesium (a more reactive metal than copper). • Magnesium, a more reactive metal than copper, takes oxygen away from copper(II) oxide. • Mg(s) + CuO(s) MgO(s) + Cu(s)

  41. Can copper reduce magnesium oxide? • No. Why? • Copper is less reactive than magnesium.

  42. Thermit / Thermite Reaction • For welding railway lines. • Heating aluminium powder with iron(III) oxide • 2Al(s) + Fe2O3(s) Al2O3(s) + 2Fe(s)

  43. Metal displacement reaction • What do you observe when copper is added into silver nitrate solution? • Brown copper dissolves slowly. Silvery grey silver crystals form on the surface of copper. The colourless solution turns pale blue. • Copper is more reactive than silver / is higher than silver in the reactivity series. • Cu(s) + 2AgNO3(aq) Cu(NO3)2(aq) + 2Ag(s)

  44. Ionic equation • Which chemical species (ions) do not take part in the above chemical reaction? • Nitrate ion, NO3-, is the spectators ion. • Can be deleted from the balanced equation. • Cu(s) + 2Ag=(aq)  Cu2+(aq) + 2Ag(s)

  45. Adding zinc into copper(II) sulphate solution • What do you see? • Zinc slowly dissolves. Brown solids form on the surface of zinc. The blue solution turns pale blue. • Zinc is more reactive than copper / is higher than copper in the reactivity series. • Zn(s) + CuSO4(aq)  ZnSO4(aq) + Cu(s) • Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s)

  46. Adding copper to magnesium sulphate solution • What do you see? • No observable change • Why? • Copper is less reactive than magnesium / is lower than magnesium in the reactivity series.

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