Reactivity of metals

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.

Metal / metal compound Colour of the flame
Sodium
Potassium
Calcium
Magnesium (only metal)
Copper
Barium
Strontium
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 gt Mg gt 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 gt Na gtLi)
• 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.