Download Unit 10

Chapter 12
Reactivity of metals
Criteria for comparing the reactivity
of metals

The temperature at which the reaction starts.
–

The rate/speed of the reaction
–

The more reactive the metal is, the lower the
temperature required.
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.
Chemical reactions for determining
the reactivity series



Reaction with air / oxygen
Reaction with water
Reaction with dilute acid
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.
Reaction with air (Heating in air)

Some metals / their metal compounds burn
with a characteristic coloured flame.
Metal / metal
compound
Sodium
Potassium
Calcium
Magnesium (only metal)
Copper
Barium
Strontium
Colour of the flame
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).
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) →
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.
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.
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)
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
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)
Test for hydrogen gas


Put a burning splint near the mouth of the test
tube.
A pop sound is heard.
Never add sodium / potassium into
dilute acids


Why?
Sodium / potassium (Group I metal) reacts
explosively with dilute acids.
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.
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
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.
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.
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+ + eNon-metal reacts by gaining electrons.
Chlorine reacts by gaining electrons.
Cl2 + 2e-  2Cl-
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.)
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.
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.
Application of reactivity series




Extraction of metals from ores
Thermit reaction (reduction with metals)
Metal displacement reaction
Predicting the stability of metal compounds
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.
Different methods of extracting
metals



Heating metal oxides alone.
Heating metal oxides with carbon (coke)
Electrolysis of hot molten ores
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?
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)
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)
Heating mercury(II) sulphide in air


Reacts with air to form mercury and sulphur
dioxide.
HgS(s) + O2(g)  Hg(l) + SO2(g)
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.
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
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)
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.
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.
Heating copper(II) oxide with
carbon


The black copper(II) oxide turns reddish brown.
2CuO(s) + C(s)  2Cu(s0 + CO2(G)
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.
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)
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.
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).
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)
Can copper reduce magnesium
oxide?


No. Why?
Copper is less reactive than magnesium.
Thermit / Thermite Reaction



For welding railway lines.
Heating aluminium powder with iron(III) oxide
2Al(s) + Fe2O3(s)  Al2O3(s) + 2Fe(s)
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)
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)
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)
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.