Chapter 11
Reactivity of metals
11.1 Comparing reactivities of common metals
11.2 The metal reactivity series
11.3 Chemical equations
11.4 Metal reactivity series and the tendency of
metals to form positive ions
11.5 Displacement reactions of metals
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11.6 Ionic equations
11.7 Relation between the extraction method and
the metal reactivity series
Key terms
Progress check
Summary
Concept map
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11.1 Comparing reactivities of common
metals
Different reactivities of metals
Some metals are more reactive than others.
Figure 11.1 Potassium reacts
vigorously with water.
Figure 11.2 Gold is a very
unreactive metal. It does not
react with air and water.
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Learning tip
Potassium metal catches fire easily when placed
in oxygen or when added to water. Hence, it is
flammable.
Reactivity of a metal refers to the readiness of it
to react with other substances.
Three factors are considered when comparing the
reactivities of metals:
1. The lowest temperature at which the reaction
starts
2. The rate of reaction
3. The amount of heat given out during reaction
11.1 Comparing reactivities of common metals
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Metal
Conditions
for reaction
Observation
Word equation
gentle
heating
It burns vigorously
with a lilac (pale
purple) flame to
produce a white
powder.
potassium + oxygen
→ potassium oxide
Sodium
gentle
heating
It burns vigorously
with a golden yellow
flame to produce a
white powder.
sodium + oxygen
→ sodium oxide
strong
heating
It burns quite
vigorously with a
brick-red flame to
produce a white
powder.
calcium + oxygen
→ calcium oxide
BURN
Potassium
Calcium
Table 11.1 Reactions of some common metals with air.
11.1 Comparing reactivities of common metals
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Metal
Conditions
for reaction
Observation
Word equation
strong
heating
It burns with a very magnesium + oxygen
bright white flame → magnesium oxide
to produce a white
powder.
Aluminium
strong
heating
Aluminium powder aluminium + oxygen
burns to give out
→ aluminium oxide
much heat; a white
powder forms.
strong
heating
Zinc powder burns
to give out some
heat; a powder
(yellow when hot,
white when cold)
forms.
BURN
Magnesium
Zinc
zinc + oxygen
→ zinc oxide
Table 11.1 Reactions of some common metals with air.
11.1 Comparing reactivities of common metals
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Metal
BURN
Iron
Conditions
for reaction
strong
heating
Lead
strong
heating
Mercury
DO NOT BURN
Copper
Observation
Word equation
Iron powder burns iron + oxygen
with yellow
→ iron(II, III) oxide
sparks to produce
a black solid.
It melts; a powder
(orange when hot,
yellow when cold)
is seen on the
surface.
very strong Its surface turns
heating
black.
lead + oxygen
→ lead(II) oxide
copper + oxygen
→ copper(II) oxide
very strong A red powder
mercury + oxygen
heating
forms on the
→ mercury(II) oxide
surface.
Table 11.1 Reactions of some common metals with air.
11.1 Comparing reactivities of common metals
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Metal
Conditions
for reaction
Silver
Gold
NO REACTION
Platinum
Observation
Word equation
__
No observable
change even on
very strong
heating.
__
__
No observable
change even on
very strong
heating.
__
__
No observable
change even on
very strong
heating.
__
Table 11.1 Reactions of some common metals with air.
11.1 Comparing reactivities of common metals
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(a) Potassium burns
with a lilac flame.
(b) Sodium burns with
a golden yellow flame.
(d) Magnesium burns with
a very bright white flame.
(c) Calcium burns with
a brick-red flame.
(e) Iron burns with
yellow sparks
Figure 11.3 Different
metals burn in air to give
different flame colours.
11.1 Comparing reactivities of common metals
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Metal
Reaction with air Word equation
most reactive Potassium, K
Sodium, Na
Calcium, Ca
react and
burn in air
(oxygen)
Magnesium, Mg
Reactivity
of metals
towards air
decreases
Aluminium, Al
Zinc, Zn
metal + oxygen
→ metal oxide
Iron, Fe
Lead, Pb
Copper, Cu
Mercury, Hg
react but do not
burn in air
(oxygen)
Silver, Ag
Platinum, Pt
least reactive Gold, Au
no reaction
with air (oxygen)
__
Table 11.2 The reactivity of common metals towards air (oxygen).
11.1 Comparing reactivities of common metals
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Appearance of metals and storage methods
All metals look shiny when they are freshly cut.
But the shiny surface of very reactive metals soon
becomes dull when exposed to air.
Metals react with oxygen in air, forming an oxide
layer on the metal surface.
Very reactive metals (such as potassium and
sodium) are stored under paraffin oil.
11.1 Comparing reactivities of common metals
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Calcium, which is a quite reactive metal, is kept in
an airtight container.
Figure 11.4 Sodium is stored
under paraffin oil while calcium
is stored in an airtight container.
Gold is the least reactive of all metals. It does not
react with oxygen and is always shiny and
attractive in appearance.
Class practice 11.1
11.1 Comparing reactivities of common metals
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Action of water on potassium
Melts to form a silvery ball
Moves about very quickly on the water surface
with a hissing sound
Burns with a lilac flame
potassium + water → potassium hydroxide + hydrogen
The resultant solution is alkaline.
∵ potassium hydroxide is produced which turns
red litmus paper blue.
11.1 Comparing reactivities of common metals
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Learning tip
A large amount of heat is produced when potassium
reacts with water. The heat causes potassium to melt.
Action of water on sodium
Less vigorously than potassium
Melts to form a silvery ball.
Moves about quickly on the water surface
Burns with a golden yellow flame
11.1 Comparing reactivities of common metals
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sodium + water → sodium hydroxide + hydrogen
The resultant solution is alkaline.
∵ sodium hydroxide is produced.
Figure 11.5 Potassium reacts
vigorously with water.
Figure 11.6 Sodium reacts with water
less vigorously than potassium.
11.1 Comparing reactivities of common metals
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Action of water on calcium
When we add a calcium granule to cold water,
it sinks to the bottom.
Colourless gas bubbles form at a moderate rate.
The gas can be collected using an inverted funnel.
When the gas is tested with a burning splint, it
burns with a ‘pop’ sound.
calcium + water → calcium hydroxide + hydrogen
11.1 Comparing reactivities of common metals
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hydrogen
water
inverted
funnel
Figure 11.7 Calcium reacts with
calcium
water at a moderate rate.
granule
Calcium granule gradually decreases in size and
eventually disappears.
A milky suspension (calcium hydroxide solid) is
produced, which is only slightly soluble in water.
Skill corner 11.1
11.1 Comparing reactivities of common metals
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Action of steam on magnesium
Magnesium has almost no reaction with cold
water.
It reacts slowly with hot water to give magnesium
hydroxide (only slightly soluble in water) and
hydrogen.
magnesium + water → magnesium hydroxide + hydrogen
Magnesium reacts more vigorously with steam.
11.1 Comparing reactivities of common metals
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wet sand magnesium ribbon
delivery tube
hydrogen
water
heat
trough
Figure 11.8 The reaction of magnesium with steam.
11.1 Comparing reactivities of common metals
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With strong heating, the water in the wet sand
changes into steam.
Steam then reacts with magnesium to give an
intense white light.
A white solid, magnesium oxide, forms.
magnesium + steam → magnesium oxide + hydrogen
11.1 Comparing reactivities of common metals
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Action of steam on aluminium, zinc and iron
Zinc and iron do not react with cold or hot water.
They react with steam.
The reaction is less vigorous for zinc, and even
less for iron.
zinc + steam → zinc oxide + hydrogen
iron + steam → iron(II, III) oxide + hydrogen
11.1 Comparing reactivities of common metals
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Learning tip
Iron(II, III) oxide refers to a mixture of iron(II) oxide
and iron(III) oxide.
Aluminium metal is usually covered with a very
thin layer of aluminium oxide which protects the
metal from reaction.
If the oxide layer is removed, the aluminium
obtained would be more reactive than zinc.
aluminium + steam → aluminium oxide + hydrogen
(after removing
the oxide layer)
11.1 Comparing reactivities of common metals
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Lead, copper, mercury, silver, platinum and gold,
even if heated strongly, do not react with steam.
Learning tip
Remember that magnesium can react with hot water
and steam.
11.1 Comparing reactivities of common metals
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Metal
most reactive
Reaction with
water or steam
Equation
metals react with
cold water
metal + water →
metal hydroxide +
hydrogen
heated metals
react with steam
metal + steam →
metal oxide +
hydrogen
heated metals do
not react with
water or steam
___
Potassium, K
Sodium, Na
Calcium, Ca
Magnesium, Mg
Reactivity
of metals
towards
water or
steam
decreases
Aluminium, Al
Zinc, Zn
Iron, Fe
Lead, Pb
Copper, Cu
Mercury, Hg
Silver, Ag
Platinum, Pt
least reactive
Gold, Au
Table 11.3 The reactivity of common
metals towards water or steam.
Class practice 11.2
11.1 Comparing reactivities of common metals
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Reactions of metals with dilute hydrochloric acid
and dilute sulphuric acid
When we add a magnesium ribbon to dilute
hydrochloric acid, magnesium dissolves and
many colourless gas bubbles are given out.
The test tube quickly becomes warm as heat is
given out.
magnesium + dilute hydrochloric acid
→ magnesium chloride + hydrogen
11.1 Comparing reactivities of common metals
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Similar observations can be made when we add
a magnesium ribbon to dilute sulphuric acid.
magnesium + dilute sulphuric acid
→ magnesium sulphate + hydrogen
iron
magnesium
dilute
dilute
hydrochloric
hydrochloric
acid
acid
(a) Magnesium reacts vigorously
(b) Iron reacts slowly with
with dilute hydrochloric acid.
dilute hydrochloric acid.
Figure 11.9
11.1 Comparing reactivities of common metals
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copper
dilute
hydrochloric
acid
(c) There is no
observable change
when copper is added to
dilute hydrochloric acid.
Figure 11.9 (a) Magnesium, (b) iron and (c) copper react
differently with dilute hydrochloric acid.
11.1 Comparing reactivities of common metals
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Metal
Reaction with dilute acid
most reactive Potassium, K
Sodium, Na
very slow reaction
metal + dilute
hydrochloric
acid → metal
chloride +
hydrogen
or
metal + dilute
sulphuric acid →
metal sulphate +
hydrogen
no reaction
___
explosive reaction
Calcium, Ca
Reactivity of
metals
towards
dilute
hydrochloric/
sulphuric
acid
decreases
Magnesium, Mg
Aluminium, Al
Zinc, Zn
reacts with dilute acid,
more slowly down the
series
Iron, Fe
Lead, Pb
Equation
Copper, Cu
Mercury, Hg
Silver, Ag
Platinum, Pt
least reactive Gold, Au
Table 11.4 The reactivity of common
metals towards dilute hydrochloric
acid or dilute sulphuric acid.
Experiment 11.1
Class practice 11.3
11.1 Comparing reactivities of common metals
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Experiment 11.1
11.2 The metal reactivity series
Comparing their reactions with air, water and
dilute acids
arrange common metals in order
of reactivity
Metal reactivity series
Metals at the top of the series are the most
reactive.
Metals at the bottom are the least reactive.
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most reactive
Potassium, K
Sodium, Na
Calcium, Ca
Magnesium, Mg
Aluminium, Al
Zinc, Zn
Iron, Fe
decreasing reactivity
Lead, Pb
Copper, Cu
Mercury, Hg
Silver, Ag
Platinum, Pt
Figure 11.10 Metal
reactivity series for
common metals.
Gold, Au
least reactive
Example 11.1
Class practice 11.4
11.2 The metal reactivity series
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11.3 Chemical equations
What is a chemical equation?
When magnesium burns in air (or oxygen),
magnesium oxide is produced.
magnesium + oxygen → magnesium oxide
magnesium
oxide
magnesium
Figure 11.11 Magnesium burns
in air to form magnesium oxide.
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In the reaction between magnesium and oxygen,
two magnesium atoms react with one oxygen
molecule to form two formula units of magnesium
oxide.
OO
Mg
1 oxygen
molecule
Mg
2 magnesium
atoms
Mg2+ O2– Mg2+ O2–
2 formula units of
magnesium oxide
Figure 11.12 A simple diagram showing what happens to the
particles during the reaction between magnesium and oxygen.
11.3 Chemical equations
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Learning tip
One formula unit of magnesium oxide consists of
one magnesium ion (Mg2+) and one oxide ion (O2–).
Chemists usually use chemical equations to
represent reactions instead.
Chemical equation for this reaction is written as:
2Mg(s) + O2(g) → 2MgO(s)
Think about
11.3 Chemical equations
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What does a chemical equation tell us?
Example
2Mg(s) + O2(g) → 2MgO(s)
1. The reactants
involved
These are magnesium (Mg) and oxygen (O2),
written on the left-hand side of the arrow.
2. The products
formed
This is magnesium oxide (MgO), written on the
right-hand side of the arrow.
3. Physical states of
the substances
involved
Mg and MgO are solids, represented by a state
symbol (s); O2 is a gas (g). Other state symbols
are: liquid (l) and aqueous solution (aq).
4. The relative number 2 atoms of Mg would react with 1 molecule of
of particles (i.e. atoms, O2 to produce 2 formula units of MgO.
molecules, ions or
formula units)
Table 11.5 The information that a chemical equation contains.
11.3 Chemical equations
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Single arrow ‘→’ between the two sides of an
equation indicates that the reaction goes one
way only.
Double arrow ‘
’ is also used, e.g.
N2(g) + 3H2(g)
2NH3(g)
The ‘
’ means that the reaction is reversible
i.e. both forward and backward reactions occur at
the same time.
Class practice 11.5
11.3 Chemical equations
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Balancing a chemical equation
Example
2Mg(s) + O2(g) → 2MgO(s)
2Mg(s) + O2(g)
→
2MgO(s)
left-hand side
right-hand side
2 Mg atoms
2 Mg atoms
2 O atoms
2 O atoms
The numbers of magnesium atoms and oxygen
atoms are the same on both sides of the equation.
The equation is balanced.
11.3 Chemical equations
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Learning tip
Atoms cannot be created or destroyed in a reaction.
They just rearrange to give new substances.
The numbers in front of the formulae of reactants
and products in a balanced chemical equation
are called stoichiometric coefficients.
11.3 Chemical equations
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Steps in writing a chemical equation
1. Stoichiometric coefficients must be placed in
front of the formulae where necessary. The
formulae themselves must not be changed.
2. The stoichiometric coefficient in front of a
chemical formula is different from the subscript in
a chemical formula.
Problem-solving strategy 11.1
11.3 Chemical equations
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stoichiometric coefficient (affects both H and O)
4H2O
subscript (affects only H)
The coefficient ‘4’ means that there are 4 water
molecules.
There are totally 8 hydrogen atoms and 4
oxygen atoms.
The subscript ‘2’ means that there are 2 hydrogen
atoms in a water molecule.
11.3 Chemical equations
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3. Some equations involve polyatomic ions (e.g.
SO42–, NO3–, OH–). When balancing such
equations, we should consider the polyatomic
ions as a single unit. For example, if there are
two SO42– ions on the reactant side of the
equation, there should be two SO42– ions on the
product side.
Example 11.2
Class practice 11.6
11.3 Chemical equations
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11.4 Metal reactivity series and the tendency
of metals to form positive ions
Metals react by losing electrons to form positive ions
Each magnesium atom loses two outermost
shell electrons
magnesium ion (Mg2+)
Electrons lost from magnesium atoms are gained
by oxygen atoms
oxide ions (O2–)
Figure 11.13 Electron diagrams showing the formation of magnesium
oxide from the reaction between magnesium and oxygen.
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Learning tip
A magnesium atom loses two outermost shell
electrons in order to get the electronic arrangement
of a noble gas atom.
Key point
Metals react by losing electrons to form
positive ions
____________.
11.4 Metal reactivity series and the tendency of metals
to form positive ions
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Reactivity and readiness to lose electrons
The more readily the metal loses electrons, the
more reactive is the metal.
The readiness of elements to lose electrons
decreases as we go across a period.
E.g. Across Period 3, reactivity of metals:
Na > Mg > Al
The readiness of elements to lose electrons
increases down a group.
E.g. Reactivity of metals:
Li < Na < K (Group I)
Be < Mg < Ca (Group II)
11.4 Metal reactivity series and the tendency of metals
to form positive ions
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increasing readiness to lose electrons
increasing reactivity of metals
increasing
readiness to
lose electrons
increasing
reactivity
of metals
Figure 11.14 Readiness of metals to lose electrons (and hence reactivity
of metals) decreases across a period and increases down a group.
11.4 Metal reactivity series and the tendency of metals
to form positive ions
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From the above reasoning, the order of reactivity
of some metals can be explained:
K > Na, Ca > Mg > Al...
Key point
higher in the reactivity series has a
A metal _______
lose
higher reactivity, and its atoms would _____
_________
electrons to form positive ions more easily.
11.4 Metal reactivity series and the tendency of metals
to form positive ions
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11.5 Displacement reactions of metals
Copper in silver nitrate solution
When we place copper in silver nitrate solution,
the copper slowly dissolves.
Some shiny silver deposits form on the copper
surface.
The solution gradually turns pale blue.
Cu(s) + 2AgNO3(aq) → Cu(NO3)2(aq) + 2Ag(s)
colourless
pale blue
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copper
wire
silver
deposits
Figure 11.15 Copper displaces silver
from silver nitrate solution. Note the
silver deposits formed on the copper
wire and the pale blue colour of the
resultant solution.
Key point
Displacement reaction is a reaction in which
another element from
one element displaces _______________
its compound.
11.5 Displacement reactions of metals
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Zinc in copper(II) sulphate solution
A similar displacement reaction occurs when we
place zinc into copper(II) sulphate solution.
Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)
silvery
blue
colourless
reddish brown
When putting copper into zinc sulphate solution
no reaction occurs.
11.5 Displacement reactions of metals
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reddish brown
copper coated on
zinc
zinc
copper(II)
sulphate
solution
Figure 11.16 The reaction between zinc and copper(II) sulphate solution.
The part of the zinc metal dipped into the solution is coated with reddish
brown copper. The blue colour of copper(II) sulphate solution becomes paler.
Experiment 11.2
Experiment 11.2
11.5 Displacement reactions of metals
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Key point
higher in the reactivity series
A metal (M1) _______
lower in the
will displace any metal (M2) _______
series from the solution of a compound of M2.
A metal higher in the reactivity series is more
reactive.
its atoms lose electrons more readily to form
positive ions.
The positive ions of the less reactive metal would
accept these electrons.
forming the atoms of the less reactive metal.
Example 11.3
11.5 Displacement reactions of metals
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11.6 Ionic equations
Example
Displacement reaction between copper and silver
nitrate solution
copper
Ag Ag
Cu Cu
Ag Ag Ag
Cu Cu Cu
Ag Ag
copper
Cu Cu
NO3–
–
Ag+ NO3
Cu2+
Ag+
silver
nitrate
solution
Ag+
Cu2+
NO3–
NO3–
Ag+
NO3–
NO3–
NO3–
NO3–
silver
deposits
coated on
copper
P. 51 / 72
Chemical equation:
Cu(s) + 2AgNO3(aq) → Cu(NO3)2(aq) + 2Ag(s)
Silver nitrate and copper(II) nitrate are ionic
compounds which are soluble in water.
In aqueous solutions, they appear as mobile ions
i.e. Ag+, Cu2+ and NO3–.
Rewrite the chemical equation:
Cu(s) + 2Ag+(aq) + 2NO3–(aq) → Cu2+(aq) + 2NO3–(aq) + 2Ag(s)
silver nitrate solution
copper(II) nitrate solution
11.6 Ionic equations
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Nitrate ions (NO3–) remain unchanged in the
displacement reaction.
Ions which do not actually take part in the reaction
spectator ions
By cancelling out the spectator ions from the
equation, we get:
Cu(s) + 2Ag+(aq) → Cu2+(aq) + 2Ag(s)
This is called an ionic equation.
11.6 Ionic equations
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Learning tip
An ionic equation must be balanced with respect to
the net ionic charges.
Key point
An ionic equation is an equation which includes
formed or _________
changed
only those ions that are _______
during the reaction.
Example 11.4
Problem-solving strategy 11.2
Example 11.5
Class practice 11.7
11.6 Ionic equations
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11.7 Relation between the extraction method
and the metal reactivity series
most difficult
Ease of
extraction
easiest
most reactive
Potassium, K
Sodium, Na
Calcium, Ca
Magnesium, Mg
Aluminium, Al
Zinc, Zn
Reactivity
Iron, Fe
Lead, Pb
Copper, Cu
Mercury, Hg
Silver, Ag
Platinum, Pt
Gold, Au
least reactive
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Metals at the top of the reactivity series (e.g.
potassium, sodium) give stable metal ores and
are more difficult to be extracted.
Electrolysis is used for extraction of these
metals from their molten ores.
Metals in the middle of the series (e.g. zinc, iron)
are easier to be extracted.
These metals are often extracted by heating
their metal ores with carbon (i.e. carbon
reduction)
11.7 Relation between the extraction method and the
metal reactivity series
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Metals near the bottom of the series (e.g. copper,
mercury) are the easiest to be extracted.
These metals are often extracted by heating
the ore alone (or in air) or displacement
from solutions.
11.7 Relation between the extraction method and the
metal reactivity series
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Metals at the bottom of the series (e.g. gold) are
extracted by mechanical separation.
Key point
The lower the position of the metal in the
more easily it can
reactivity series, the ____________
be extracted from its ore.
Example 11.6
Class practice 11.8
11.7 Relation between the extraction method and the
metal reactivity series
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Key terms
1.
2.
3.
4.
5.
6.
7.
8.
9.
balanced chemical equation 平衡化學方程式
chemical equation 化學方程式
displace 置換
displacement reaction 置換反應
ionic equation 離子方程式
metal reactivity series 金屬活性序
reactivity 活性
spectator ion 旁觀離子
stoichiometric coefficient 計量系數
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Progress check
1.
2.
3.
4.
How can we compare the reactivity of metals?
How do some common metals react with oxygen?
How do some common metals react with water?
How do some common metals react with dilute
acids?
5. How can we construct a metal reactivity series?
6. How can we write the chemical equations for the
reactions of common metals with oxygen, water
and dilute acids?
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7. How can we transcribe word equations into
chemical equations?
8. How can we write balanced chemical equations to
describe various reactions?
9. What is the relation between the reactivity of a
metal and its tendency to form a positive ion?
10. What is a displacement reaction?
11. How can we write balanced ionic equations?
12. What is the relation between the extraction
methods and the reactivity of metals?
Progress check
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Summary
11.1 Comparing reactivities of common metals
1.
2.
Reactivity of a metal is the readiness of it to
react with other substances.
The reactivity of metals can be found by
comparing their reactions with air, water and
dilute acids. Refer to p.4–11 for the results of
the reactions.
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11.2 The metal reactivity series
3.
4.
The metal reactivity series is a series of
common metals arranged in a decreasing order
of reactivity.
The following table summarizes the
appearances and reactions of metals in the
reactivity series.
Summary
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Metal
Appearance
of metal
K
Na
dull (stored
under paraffin
oil)
Ca
Mg
Al
Zn
generally dull
Fe
Pb
Cu
Hg
Ag
Reaction of metal with
air (oxygen)
water/steam
burns
vigorously,
forming metal
oxide
(Example 1)
metal +
water →
metal
hydroxide +
hydrogen
(Example 4)
reacts with
decreasing
vigour,
forming metal
oxide
(Example 2)
a layer of
metal oxide
formed on the
generally shiny surface
(Example 3)
Au
metal +
steam →
metal oxide
+ hydrogen
(Example 5)
no reaction
dilute
hydrochloric acid
reacts explosively,
forming metal
chloride and
hydrogen
(Example 6)
reacts with
decreasing vigour:
metal +
hydrochloric acid
→ metal chloride +
hydrogen
(Example 7)
no reaction
no reaction
Summary
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Displacement
reaction
not applicable
∵ these three
metals react
with water to
give hydrogen
a metal
displaces any
other metal
lower in the
series from a
solution of its
compound
(Example 8)
Example 1: 4Na(s) + O2(g) → 2Na2O(s)
Example 2: 2Ca(s) + O2(g) → 2CaO(s)
Example 3: 2Cu(s) + O2(g) → 2CuO(s)
Example 4: 2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
Example 5: Zn(s) + H2O(g) → ZnO(s) + H2(g)
Example 6: 2K(s) + 2HCl(aq) → 2KCl(aq) + H2(g)
(NEVER attempt this experiment!)
Example 7: Fe(s) + 2HCl(aq) → FeCl2(aq) + H2(g)
Example 8: Mg(s) + 2AgNO3(aq)
→ Mg(NO3)2(aq) + 2Ag(s) or
Ionic equation: Mg(s) + 2Ag+(aq) → Mg2+(aq) + 2Ag(s)
Summary
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11.3 Chemical equations
5.
6.
7.
A chemical equation shows the physical states
and relative numbers of particles of the
reactants and products in a chemical reaction.
A reversible reaction is represented by a double
arrow ‘
’.
The steps in writing a chemical equation are
shown in ‘Problem-solving strategy 11.1’ on
p.16.
Summary
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11.4 Metal reactivity series and the tendency of
metals to form positive ions
8.
Metals react by losing electrons to form positive
ions. Different metals have different reactivities
because they have different tendencies to lose
electrons. Atoms of a reactive metal lose
electrons readily.
Summary
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11.5 Displacement reactions of metals
9.
A metal (M1) higher in the reactivity series will
displace any metal (M2) lower in the series from
the solution of a compound of M2. This is
because a more reactive metal loses electrons
more readily.
11.6 Ionic equations
10.
An ionic equation is an equation which includes
only those ions that are formed or changed
during the reaction.
Summary
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11.
An ionic equation must be balanced with
respect to the ionic charges as well as the
number of atoms. (Refer to ‘Problem-solving
strategy 11.2’ on p.23.)
11.7 Relation between the extraction method and
the metal reactivity series
12.
13.
The method used to extract a metal from its
ores depends on the position of the metal in the
reactivity series.
The lower the position of the metal in the
reactivity series, the more easily it can be
extracted from its ore.
Summary
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Concept map
Extraction
method
of metal
Readiness of
metals to
lose
________
electrons and
form
metal ions
___________
depends on
the position
of metal in
Metal
reactivity
series
arranged in
decreasing
order
REACTIVITY
OF
METALS
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Metal
reactivity
series
can be used
to predict
arranged in
decreasing
order
REACTIVITY
OF
METALS
Displacement
_____________
reactions
can be
compared by
The lowest
temperature at
____________
which the reaction
starts
The rate of reaction
The amount of
heat given out
_____
Concept map
P. 71 / 72
REACTIVITY
OF METALS
reaction of
metals with
Oxygen
gives
Metal
oxide
Water
if no reaction,
allow metal to
react with
gives
Metal hydroxide
_______________
and hydrogen
Steam
Dilute acid
gives
gives
Metal oxide
and
hydrogen
Salt
_______
and
hydrogen
Concept map
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