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Chapter 30
Redox reactions
30.1 Oxidation and reduction
30.2 Oxidation and reduction in terms of changes
in oxidation numbers
30.3 Common oxidizing agents and reducing agents
30.4 Balancing redox equations
30.5 The Electrochemical Series and the relative
strength of oxidizing agents/reducing agents
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30.6 Uses of the Electrochemical Series
30.7 Chlorine as an oxidizing agent
30.8 Nitric acid of different concentrations as
oxidizing agents
30.9 Concentrated sulphuric acid as an oxidizing
agent
30.10 Sulphite ion as a reducing agent
Key terms
Progress check
Summary
Concept map
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30.1 Oxidation and reduction
Examples of oxidation-reduction
reaction (or redox reaction)
Example 1:
Metal objects corrode in the presence of oxygen and
water.
Example 2:
The chemical reactions in chemical cells.
Learning tip
The word ‘redox’ comes from two words —
reduction and oxidation.
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Oxidation and reduction in terms of gain or loss of
oxygen
The reaction between magnesium and copper(II)
oxide involves a gain and loss of oxygen.
(gains oxygen)
oxidation
Mg(s) + CuO(s) → MgO(s) + Cu(s)
reduction
(loses oxygen)
30.1 Oxidation and reduction
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Key point
Oxidation is the process in which a substance
gains oxygen.
Reduction is the process in which a substance
loses oxygen.
Oxidation and reduction must occur together. One
cannot take place without the other.
30.1 Oxidation and reduction
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(gains oxygen)
oxidation
Mg(s) + CuO(s) → MgO(s) + Cu(s)
reducing agent
oxidizing agent
reduction
(loses oxygen)
Copper(II) oxide is the oxidizing agent (or oxidant).
It oxidizes the other reactant and is reduced at
the same time.
Magnesium is the reducing agent (or reductant).
It reduces the other reactant and is oxidized at
the same time.
30.1 Oxidation and reduction
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Key point
An oxidizing agent is a substance which oxidizes
others by losing oxygen.
A reducing agent is a substance which reduces
others by gaining oxygen.
Example 30.1
30.1 Oxidation and reduction
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Oxidation and reduction in terms of gain or loss of
hydrogen
The reaction between hydrogen sulphide and
oxygen involves a gain and loss of hydrogen.
(loses hydrogen)
oxidation
2H2S(g) + O2(g) → 2S(s) + 2H2O(l)
reduction
(gains hydrogen)
30.1 Oxidation and reduction
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Key point
Oxidation is the process in which a substance loses
hydrogen.
Reduction is the process in which a substance gains
hydrogen.
30.1 Oxidation and reduction
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(loses hydrogen)
oxidation
2H2S(g) + O2(g) → 2S(s) + 2H2O(l)
reducing agent oxidizing agent
reduction
(gains hydrogen)
Hydrogen sulphide is a reducing agent.
It reduces oxygen and is oxidized at the same time.
Oxygen is an oxidizing agent.
It oxidizes hydrogen sulphide and is reduced at
the same time.
30.1 Oxidation and reduction
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Key point
An oxidizing agent is a substance which oxidizes
others by gaining hydrogen.
A reducing agent is a substance which reduces
others by losing hydrogen.
30.1 Oxidation and reduction
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Oxidation and reduction in terms of gain or loss of
electrons
The reaction between Mg and CuO:
transfer of 2e−
2+
2−
Mg + Cu O
2+
2−
→ Mg O
+ Cu
2+
Each Mg atom loses 2e− to form a Mg ion.
2+
Each Cu ion gains 2e− to form a Cu atom.
Electrons are transferred from Mg atoms to
Cu2+ ions.
30.1 Oxidation and reduction
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Key point
A redox reaction is a reaction involving gain or loss
of electrons.
Oxidation is the process in which a substance loses
electrons.
Reduction is the process in which a substance gains
electrons.
Learning tip
We may memorize the concept of oxidation and
reduction using the word ‘OIL-RIG’. Oxidation Is
Loss of electron(s). Reduction Is Gain of electron(s).
30.1 Oxidation and reduction
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The redox reaction between Mg and CuO:
(loses electrons)
oxidation
Mg(s) + CuO(s) → MgO(s) + Cu(s)
reducing agent oxidizing agent
reduction
(gains electrons)
The Mg atoms lose electrons and are oxidized.
Mg is the reducing agent.
2+
The Cu ions gain electrons and are reduced.
CuO is the oxidizing agent.
30.1 Oxidation and reduction
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Key point
An oxidizing agent is a substance which oxidizes
others by gaining electrons.
A reducing agent is a substance which reduces
others by losing electrons.
Learning tip
Examples of some common oxidizing agents and
reducing agents are shown in Tables 30.5 and 30.6
on p.13–14.
Example 30.2
Class practice 30.1
30.1 Oxidation and reduction
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30.2 Oxidation and reduction in terms of
changes in oxidation numbers
Oxidation number
Are there any redox reactions that
cannot be defined in terms of gain
or loss of oxygen, hydrogen and
electrons?
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Reaction
Any
Any gain Any gain Redox
gain or
or loss of of loss of
or
loss of
hydrogen? electrons? not?
oxygen?
(1)
2MgO(s) + C(s) → 2Mg(s) + CO2(g)
redox
(2)
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
redox
(3)
Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)
redox
(4)
P4(s) + 10Cl2(g) → 4PCl5(s)
?
?
Table 30.1 Which of the four reactions shown in the table is/are redox?
30.2 Oxidation and reduction in terms of changes in oxidation numbers
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Reaction (4) involves the reaction of only covalent
compounds formed by sharing electrons.
Is reaction (4) a redox reaction?
How can we decide whether this
reaction is redox or not?
30.2 Oxidation and reduction in terms of changes in oxidation numbers
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The concept of oxidation number (O.N.) is used
to decide whether a reaction is redox.
In this concept, each atom is given a charge —
imaginery in many cases.
Key point
The oxidation number of an element in a compound
is the charge an atom of the element would have if
the atom existed as an ion.
30.2 Oxidation and reduction in terms of changes in oxidation numbers
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Rules for assigning oxidation numbers
Rule
Example
(1) The oxidation number of an The oxidation number of
element is zero.
nitrogen in a nitrogen
molecule N2 is zero; that of Cu,
O2, Cl2 or S in the free element
is also zero.
(2) The oxidation number of an
element in a simple ion is
equal to the charge on the
ion.
In Na2O,
oxidation number of sodium in
Na+ is +1
oxidation number of oxygen in
O2– is –2.
Table 30.2 Rules of assigning oxidation numbers.
30.2 Oxidation and reduction in terms of changes in oxidation numbers
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Rule
(3) The oxidation numbers
of some elements in
their compounds are
fixed.
Example
(a) All Group I metals in their
compounds
+1
(b) All Group II metals in their
compounds
+2
(c) Hydrogen in most of its
compounds
+1
(d) Fluorine in all its compounds
−1
−1
(e) Chlorine, bromine and iodine
in most of their compounds
(f) Oxygen in most of its
compounds
Table 30.2 Rules of assigning oxidation numbers.
30.2 Oxidation and reduction in terms of changes in oxidation numbers
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−2
Rule
Example
(4) The sum of oxidation
For,
numbers of all elements in
a compound is zero.
H2O,
(+1) × 2 + (–2) = 0
+1 –2
MgCl2,
(+2) + (–1) × 2 = 0
+2 –1
For,
(5) The sum of oxidation
numbers of all elements in
a polyatomic ion is equal
to the charge on the ion.
–
OH ,
–2+1
2–
CO3 ,
(+4) + (–2) × 3 = –2
+4 –2
(6) The oxidation number of
HNO3, HNO2
+5
+3
an element may vary from
compound to compound.
+
–
NH4 , NO3
–3
(–2) + (+1) = –1
CO, CO2
+2
+4
+5
Table 30.2 Rules of assigning oxidation numbers.
Problem-solving strategy 30.1
Class practice 30.2
30.2 Oxidation and reduction in terms of changes in oxidation numbers
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Elements with different oxidation numbers
An element may have different oxidation numbers
in different compounds.
For example, the oxidation numbers of sulphur in
H2SO4, SO2, SCl2 and H2S are +6, +4, +2 and –2
respectively.
Learning tip
The oxidation number of sulphur in SO2 is +4.
However, it does not mean that there is a
charge of +4 on the sulphur atom.
30.2 Oxidation and reduction in terms of changes in oxidation numbers
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Oxidation number
Sulphur
Nitrogen
Carbon
Iron
+7
+6
H2SO4
+5
+4
HNO3
SO2
+3
+2
SCl2
NO
−3
−4
FeCl3
CO
FeSO4
C
Fe
N2O
S
N2
C2H2
−1
−2
CaCO3
HNO2
+1
0
NO2
H2S
C2H4
NH3
C2H6
CH4
Table 30.3 The different oxidation numbers of some elements in their
compounds.
30.2 Oxidation and reduction in terms of changes in oxidation numbers
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Oxidation number
Copper
Manganese
+7
KMnO4
+6
K2MnO4
Chromium
K2Cr2O7
+5
+4
MnO2
+3
Mn2O3
CrCl3
MnSO4
CrCl2
Mn
Cr
+2
CuSO4
+1
CuCl
0
Cu
Table 30.3 The different oxidation numbers of some elements in their
compounds.
30.2 Oxidation and reduction in terms of changes in oxidation numbers
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Using oxidation number to identify redox reactions
The reaction between Mg and O2 :
(O.N. of Mg increases from 0 to +2)
oxidation
2Mg(s) + O2(g) → 2MgO(s)
reduction
(O.N. of O decreases from 0 to −2)
30.2 Oxidation and reduction in terms of changes in oxidation numbers
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(O.N. of Mg increases from 0 to +2)
oxidation
2Mg(s) + O2(g) → 2MgO(s)
reducing agent oxidizing agent
reduction
(O.N. of O decreases from 0 to −2)
The O.N. of Mg increases from 0 to +2.
Mg undergoes oxidation.
It is oxidized and is a reducing agent.
The O.N. of O2 decreases from 0 to −2.
O2 undergoes reduction.
It is reduced and is a oxidizing agent.
30.2 Oxidation and reduction in terms of changes in oxidation numbers
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Key point
A redox reaction is a reaction in which the reacting
substances undergo changes in oxidation numbers.
Oxidation is the process in which the oxidation
number of an element in a substance increases.
Reduction is the process in which the oxidation
number of an element in a substance decreases.
30.2 Oxidation and reduction in terms of changes in oxidation numbers
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Key point
An oxidizing agent is a substance which oxidizes
others. Its oxidation number decreases.
A reducing agent is a substance which reduces
others. Its oxidation number increases.
Writing practice 30.1
Class practice 30.3
Writing practice 30.2
30.2 Oxidation and reduction in terms of changes in oxidation numbers
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Term
Defined in
terms of gain
or loss of
oxygen
Defined in
terms of gain
or loss of
hydrogen
Defined in
terms of gain
or loss of
electrons
Defined in terms
of changes in
oxidation
numbers
Redox
a reaction
involving gain
and loss of
oxygen
a reaction
involving gain
and loss of
hydrogen
a reaction
involving gain
and loss of
electrons
a reaction in
which the
reacting
substances
undergo changes
in oxidation
numbers
Oxidation
a process in
which a
substance gains
oxygen
a process in
which a
substance
loses hydrogen
a process in
which a
substance
loses electron(s)
a process in
which the
oxidation number
of an element in a
substance
increases
Table 30.4 Different definitions of redox reaction, oxidation, reduction,
oxidizing agent and reducing agent.
30.2 Oxidation and reduction in terms of changes in oxidation numbers
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Term
Reduction
Oxidizing
agent
Defined in
terms of gain
or loss of
oxygen
Defined in
terms of gain
or loss of
hydrogen
Defined in
terms of gain
or loss of
electrons
Defined in terms
of changes in
oxidation
numbers
a process in
which a
substance
loses oxygen
a process in
which a
substance
gains hydrogen
a process in
which a
substance
gains electron(s)
a process in
which the
oxidation number
of an element in a
substance
decreases
a substance
which oxidizes
others by losing
oxygen
a substance
which oxidizes
others by
gaining
hydrogen
a substance
which oxidizes
others by
gaining
electron(s)
a substance
which oxidizes
others by a
decrease in
oxidation number
Table 30.4 Different definitions of redox reaction, oxidation, reduction,
oxidizing agent and reducing agent.
30.2 Oxidation and reduction in terms of changes in oxidation numbers
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Term
Reducing
agent
Defined in
terms of gain
or loss of
oxygen
Defined in
terms of gain
or loss of
hydrogen
Defined in
terms of gain
or loss of
electrons
Defined in terms
of changes in
oxidation
numbers
a substance
which reduces
others by
gaining oxygen
a substance
which reduces
others by losing
hydrogen
a substance
which reduces
others by losing
electron(s)
a substance
which reduces
others by an
increase in
oxidation number
Table 30.4 Different definitions of redox reaction, oxidation, reduction,
oxidizing agent and reducing agent.
30.2 Oxidation and reduction in terms of changes in oxidation numbers
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30.3 Common oxidizing agents and reducing
agents
Oxidizing
agent
Main change &
colour change
Acidified
potassium
permanganate
solution*
MnO4 (aq) → Mn (aq)
purple
colourless (or
very pale pink)
Acidified
potassium
dichromate
solution*
Cr2O7 (aq) → Cr (aq)
orange
green
Dilute nitric
acid
NO3 (aq) → NO(g)
colourless
colourless
Half equation & change in O.N.
−
2+
MnO4 (aq) + 8H (aq) + 5e– → Mn (aq) + 4H2O(l)
+7
+2
2–
3+
Cr2O7 (aq) + 14H (aq) + 6e– → 2Cr (aq) +7H2O(l)
+6
+3
−
−
+
2–
−
2+
+
3+
+
NO3 (aq) + 4H (aq) + 3e– → NO(g) + 2H2O(l)
+5
+2
Table 30.5 Some common strong oxidizing agents.
* Potassium permanganate solution and potassium dichromate solution are
usually acidified with dilute sulphuric acid.
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Oxidizing
agent
Main change &
colour change
Concentrated
nitric acid
NO3 (aq) → NO2(g)
colourless reddish brown
NO3 (aq) + 2H (aq) + e– → NO2(g) + H2O(l)
+5
+4
Concentrated
sulphuric acid
H2SO4(l) → SO2(g)
colourless
colourless
H2SO4(l) + 2H (aq) + 2e– → SO2(g) + 2H2O(l)
+6
+4
Oxygen
−
−
O2(g) → OH (aq)
colourless colourless
Half equation & change in O.N.
+
−
O2(g) + 2H2O(l) + 4e– → 4OH (aq)
0
−2
2−
or O2(g) → O (s)
Chorine
+
−
−
Cl2(g) → Cl (aq)
pale green colourless
2−
O2(g) + 4e– → 2O (s)
0
−2
−
Cl2(g) + 2e– → 2Cl (aq)
0
−1
Table 30.5 Some common strong oxidizing agents.
30.3 Common oxidizing agents and reducing agents
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Oxidizing agent
Bromine (in
aqueous
solution)
Ions of metals
low in the
reactivity series
(e.g. Ag+)
Iron(III) salts
Main change &
colour change
−
Half equation & change in O.N.
−
Br2(aq) → Br (aq)
yellow/brown colourless
Br2(aq) + 2e− → 2Br (aq)
0
−1
positive ion → metal atom
+
e.g. Ag (aq) → Ag(s)
colourless
silvery
Ag+(aq) + e− → Ag(s)
+1
0
3+
2+
Fe (aq) → Fe (aq)
yellow
pale green
3+
2+
Fe (aq) + e− → Fe (aq)
+3
+2
Table 30.5 Some common strong oxidizing agents.
30.3 Common oxidizing agents and reducing agents
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Reducing
agent
Main change &
colour change
Half equation & change in O.N.
+
Metals high in metal atom → positive ion
+
the reactivity
e.g. Na(s) → Na (aq)
series (e.g. Na)
grey
colourless
2−
Sulphur
dioxide
SO2(g) → SO4 (aq)
colourless colourless
Sulphites
SO3 (aq) → SO4 (aq)
colourless
colourless
Iron(II) salts
2–
2+
2–
Na(s) → Na (aq) + e−
0
+1
2−
+
SO2(g) +2H2O(l) → SO4 (aq) + 4H (aq) + 2e−
+4
+6
2–
2–
SO3 (aq) + H2O(l) → SO4 (aq) + 2H+(aq) + 2e−
+4
+6
3+
Fe (aq) → Fe (aq)
pale green
yellow
2+
3+
Fe (aq) → Fe (aq) + e−
+2
+3
Table 30.6 Some common strong reducing agents.
30.3 Common oxidizing agents and reducing agents
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Reducing
agent
Hydrogen
Carbon
Iodides
Main change &
colour change
Half equation & change in O.N.
+
+
H2(g) → H (aq)
colourless colourless
H2(g) → 2H (aq) + 2e−
0
+1
C(s) → CO(g)
black
colourless
(O.N. of carbon increases from 0 to +2)
or C(s) → CO2(g)
black
colourless
(O.N. of carbon increases from 0 to +4)
I−(aq) → I2(aq)*
colourless
brown
−
2I (aq) → I2(aq) + 2e−
−1
0
Table 30.6 Some common strong reducing agents.
* Iodine (dissolved in an organic or a non-aqueous solvent) is purple in colour.
A large number of redox reactions can result from
combinations of oxidizing agents and reducing
agents.
30.3 Common oxidizing agents and reducing agents
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Example
Sodium sulphite solution is mixed with acidified
potassium dichromate solution.
The orange dichromate ions are reduced to green
chromium(III) ions.
The sulphite ions are oxidized to sulphate ions.
sodium sulphite
solution
acidified potassium
dichromate solution
Experiment 30.1
Figure 30.1 The reaction between acidified potassium
dichromate solution and sodium sulphite solution.
30.3 Common oxidizing agents and reducing agents
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Experiment 30.1
30.4 Balancing redox equations
Balancing redox equations by using half equation
method
Redox equations can be balanced by using half
equation method.
Problem-solving strategy 30.2
Class practice 30.4
Balancing redox equations by using oxidation
number method
Redox equations can be balanced by using
oxidation number method.
Problem-solving strategy 30.3
Class practice 30.5
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30.5 The Electrochemical Series and the
relative strength of oxidizing agents/
reducing agents
Arranging oxidizing/reducing agents to construct
the Electrochemical Series
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Half equation
Reduction
Oxidation
very strong
oxidizing agents
K+(aq) + e−
Ca2+(aq) + 2e−
Na+(aq) + e−
Mg2+(aq) + 2e−
Al3+(aq) + 3e−
Zn2+(aq) + 2e−
Fe2+(aq) + 2e−
Pb2+(aq) + 2e−
2H+(aq) + 2e−
SO42−(aq) + 4H+(aq) + 2e−
Cu2+(aq) + 2e−
O2(g) + 2H2O(l) + 4e–
I2(aq) + 2e–
Fe3+(aq) + e–
NO3–(aq) + 2H+(aq) + e–
Ag+(aq) + e−
Br2(aq) + 2e–
Cr2O72–(aq) + 14H+(aq) + 6e–
Cl2(g) + 2e–
MnO4–(aq) + 8H+(aq) + 5e–
S2O82–(aq) + 2e–
F2(g) + 2e–
2H2SO4(l) + 2e–
K(s)
Ca(s)
Na(s)
Mg(s)
Al(s)
Zn(s)
Fe(s)
Pb(s)
H2(g)
SO2(g) + 2H2O(l)
Cu(s)
4OH–(aq)
2I–(aq)
Fe2+(aq)
NO2(g) + H2O(l)
Ag(s)
2Br–(aq)
2Cr3+(aq) + 7H2O(l)
2Cl–(aq)
Mn2+(aq) + 4H2O(l)
2SO42–(aq)
2F–(aq)
SO42–(aq) + SO2(g) + 2H2O(l)
Table 30.7 The Electrochemical Series.
30.5 The Electrochemical Series and the relative strength of oxidizing agents/reducing agents
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very strong
reducing agents
increasing reducing power
increasing ease of gaining electrons
increasing oxidizing power
very weak
oxidizing agents
Reducing agent
increasing ease of losing electrons
Oxidizing agent
very weak
reducing agents
In the Electrochemical Series, each half equation
is written in the form
oxidizing agent + ne–
reducing agent
Electron acceptors or oxidizing agents are arranged
on the left.
Electron donors or reducing agents are arranged
on the right.
30.5 The Electrochemical Series and the relative strength of oxidizing agents/reducing agents
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Oxidizing power of oxidizing agents increases down
the series.
Reducing power of reducing agents increases up
the series.
In Table 30.7,
the strongest oxidizing agent is H2SO4(l)
the strongest reducing agent is K(s)
30.5 The Electrochemical Series and the relative strength of oxidizing agents/reducing agents
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Trend of reducing power of metals and oxidizing
power of metal ions
The metals at the top of the series lose electrons
more readily to form metal ions in a redox reaction.
They are stronger reducing agents than the
metals near the bottom of the series.
The reducing power of metals decreases down
the series.
30.5 The Electrochemical Series and the relative strength of oxidizing agents/reducing agents
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Half equation
Reduction
Oxidation
very strong
oxidizing agents
K+(aq) + e−
Ca2+(aq) + 2e−
Na+(aq) + e−
Mg2+(aq) + 2e−
Al3+(aq) + 3e−
Zn2+(aq) + 2e−
Fe2+(aq) + 2e−
Pb2+(aq) + 2e−
2H+(aq) + 2e−
Cu2+(aq) + 2e−
O2(g) + 2H2O(l) + 4e–
Fe3+(aq) + e–
Ag+(aq) + e−
K(s)
Ca(s)
Na(s)
Mg(s)
Al(s)
Zn(s)
Fe(s)
Pb(s)
H2(g)*
Cu(s)
4OH–(aq)*
Fe2+(aq)
Ag(s)
very strong
reducing agents
very weak
reducing agents
Table 30.8 Trend of reducing power of metals and oxidizing power of metal ions.
* H2(g) and H+(aq); O2(g) and OH–(aq) are not metals nor metal ions. Their half
equations are included here for comparisons.
30.5 The Electrochemical Series and the relative strength of oxidizing agents/reducing agents
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increasing reducing power
increasing ease of
gaining electrons
increasing oxidizing power
very weak
oxidizing agents
Reducing agent
increasing ease of
losing electrons
Oxidizing agent
The metals at the top of the series lose electrons
more readily to form ions.
These ions are less likely to gain electrons to form
metals.
The metal ions at the top of the series are weaker
oxidizing agents than those near the bottom of the
series.
Key point
A metal higher in the Electrochemical Series is a
stronger reducing agent and its ion is a weaker
oxidizing agent.
Class practice 30.6
30.5 The Electrochemical Series and the relative strength of oxidizing agents/reducing agents
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30.6 Uses of the Electrochemical Series
Predicting the feasibility of a redox reaction
Given two half equations in the E.C.S., we can
predict a possible reaction using the following rule:
‘The half reaction lower in the series will go
as written, while the one higher will go in
reverse.’
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Suppose the following two half equations are given:
goes in reverse
2+
Cu (aq) + 2e−
+
Ag (aq) + e−
Cu(s)
(higher in the E.C.S.)
Ag(s)
(lower in the E.C.S.)
goes as written
The lower half equation goes as written:
+
Ag (aq) + e− → Ag(s)
The higher one goes in reverse:
2+
Cu(s) → Cu (aq) + 2e−
30.6 Uses of the Electrochemical Series
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Combining the two half equations to eliminate
electrons, the ionic equation for a possible reaction:
+
2+
2Ag (aq) + Cu(s) → 2Ag(s) + Cu (aq)
copper wire
silver crystals deposit
on the copper wire
silver nitrate solution
turns pale blue
Figure 30.2 The reaction between copper and silver nitrate solution.
30.6 Uses of the Electrochemical Series
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The action of dilute hydrochloric acid on different
metals can be predicted according to the rule.
Metals above hydrogen in the series (e.g. zinc)
react with hydrochloric acid to form hydrogen.
+
2+
Zn(s) + 2H (aq) → Zn (aq) + H2(g)
Metals below hydrogen (e.g. copper, silver) have
no reaction.
30.6 Uses of the Electrochemical Series
P. 50 / 101
zinc
copper
Figure 30.3 The action of dilute hydrochloric acid on two different metals.
(a) Zinc reacts with the acid to give hydrogen gas. (b) Copper does not
react with the acid.
30.6 Uses of the Electrochemical Series
P. 51 / 101
Predicting cell reaction and direction of electron
flow in a chemical cell
light bulb
electron flow
magnesium (negative
electrode or negative pole)
copper (positive electrode
or positive pole)
copper(II) sulphate
solution
Figure 30.4 A simple chemical cell with magnesium and copper electrodes,
and copper(II) sulphate solution as electrolyte.
30.6 Uses of the Electrochemical Series
P. 52 / 101
The two related ionic equations are:
goes in reverse
2+
Mg(s)
(higher in the E.C.S.)
2+
Cu(s)
(lower in the E.C.S.)
Mg (aq) + 2e−
Cu (aq) + 2e−
goes as written
The ionic equation for the overall cell reaction is:
2+
2+
Mg(s) + Cu (aq) → Mg (aq) + Cu(s)
Electrons will flow from magnesium to copper in
the external circuit.
Class practice 30.7
30.6 Uses of the Electrochemical Series
P. 53 / 101
30.7 Chlorine as an oxidizing agent
Chlorine is a strong oxidizing agent.
In most of its reactions,
chlorine acts as an oxidizing agent
its oxidation number decreases from 0 to –1.
–
Cl2 + 2e → 2Cl
0
–
−1
P. 54 / 101
Action of aqueous chlorine on potassium bromide
solution
When aqueous chlorine is added to potassium
bromide solution, bromine is produced.
The colourless solution changes to yellowish brown.
0
−1
0
−1
–
–
Cl2(aq) + 2Br (aq) → 2Cl (aq) + Br2(aq)
very pale green, colourless
almost colourless
colourless
yellowish
brown
–
Chlorine (Cl2) is reduced to chloride ion (Cl ).
The O.N. of chlorine decreases from 0 to –1.
30.7 Chlorine as an oxidizing agent
P. 55 / 101
Since chlorine is a stronger oxidizing agent than
bromine, it displaces bromine out from the
potassium bromide solution.
add Cl2(aq)
Br2 (aq)
Br−(aq)
(a)
(b)
Figure 30.5 (a) and (b) Reaction between aqueous chlorine and potassium
bromide solution.
30.7 Chlorine as an oxidizing agent
P. 56 / 101
A non-aqueous solvent (e.g. heptane) is added to
the solution.
After shaking and allowing it to stand for some time,
the mixture separates into two layers.
The heptane layer on the top extracts most of the
bromine from the bottom aqueous layer.
The orange-red colour in the heptane layer
indicates the presence of bromine.
30.7 Chlorine as an oxidizing agent
P. 57 / 101
heptane layer
(orange-red)
add heptane
(a colourless
non-aqueous
solvent)
add Cl2(aq)
shake and
allow the test
tube to settle
Br−(aq)
(a)
aqueous layer
Br2 (aq)
(b)
(c)
Figure 30.5 Reaction between aqueous chlorine and potassium bromide
solution, followed by the addition of a non-aqueous solvent.
Think about
30.7 Chlorine as an oxidizing agent
P. 58 / 101
Action of aqueous chlorine on potassium iodide
solution
When aqueous chlorine is added to potassium
iodide solution, iodine is produced.
The colourless solution changes to brown.
0
−1
−1
–
–
0
Cl2(aq) + 2I (aq) → 2Cl (aq) + I2(aq)
very pale green,
almost colourless
colourless
add Cl2(aq)
I−(aq)
(a)
(b)
colourless
brown
I2 (aq)
Figure 30.6 (a) and (b) Reaction
between aqueous chlorine and
potassium iodide solution.
30.7 Chlorine as an oxidizing agent
P. 59 / 101
Since chlorine is a stronger oxidizing agent than
iodine, it displaces iodine out from the potassium
iodide solution.
If a non-aqueous solvent (e.g. heptane) is added
and the solution is shaken and allowed to stand for
some time, the heptane layer on the top extracts
most of the iodine from the bottom aqueous layer.
The purple colour in the heptane layer indicates the
presence of iodine.
30.7 Chlorine as an oxidizing agent
P. 60 / 101
heptane layer
(purple)
add heptane
(a colourless
non-aqueous
solvent)
add Cl2(aq)
shake and
allow the test
tube to settle
I−(aq)
(a)
aqueous layer
I2 (aq)
(b)
(c)
Figure 30.6 Reaction between aqueous chlorine and potassium iodide
solution, followed by the addition of a non-aqueous solvent.
30.7 Chlorine as an oxidizing agent
P. 61 / 101
Action of aqueous chlorine on sodium hydroxide
solution
For some reactions, chlorine is reduced and
oxidized at the same time.
The reaction in which a species is simultaneously
reduced and oxidized is called disproportionation.
30.7 Chlorine as an oxidizing agent
P. 62 / 101
When chlorine gas is passed into a cold dilute
sodium hydroxide solution, sodium chloride (NaCl)
and sodium hypochlorite (NaOCl) are produced.
Cl2(g) + 2NaOH(aq) → NaCl(aq) + NaOCl(aq) + H2O(l)
or
–
–
–
Cl2(g) + 2OH (aq) → Cl (aq) + OCl (aq) + H2O(l)
0
−2 +1
−1
−2 +1
+1 −2
The O.N. of chlorine in Cl2 is 0 but it becomes
–1 and +1 in NaCl and NaOCl respectively.
Chlorine is simultaneously reduced and oxidized.
30.7 Chlorine as an oxidizing agent
P. 63 / 101
When Cl2(g) is passed into a hot concentrated
sodium hydroxide solution, sodium chloride (NaCl)
and sodium chlorate (NaClO3) are produced.
3Cl2(g) + 6NaOH(aq) → 5NaCl(aq) + NaClO3(aq) + 3H2O(l)
or
3Cl2(g) + 6OH–(aq) → 5Cl–(aq) + ClO3–(aq) + 3H2O(l)
0
−2 +1
−1
+5 −2
+1 −2
The O.N. of chlorine in Cl2 in NaCl and NaClO3
are –1 and +5.
Chlorine is simultaneously reduced and oxidized.
Example 30.3
Class practice 30.8
30.7 Chlorine as an oxidizing agent
P. 64 / 101
30.8 Nitric acid of different concentrations
as oxidizing agents
Both dilute nitric acid and concentrated nitric acid
are strong oxidizing agents.
They behave quite differently from very dilute nitric
acid.
P. 65 / 101
Dilute nitric acid
Action on metals
Dilute nitric acid is added to copper turnings in a
test tube and the mixture is heated gently.
Copper is oxidized to copper(II) ions and the
solution turns blue.
Nitric acid is reduced to colourless nitrogen
monoxide gas.
30.8 Nitric acid of different concentrations as oxidizing agents
P. 66 / 101
The half equation for the oxidation is:
2+
Cu(s) → Cu (aq) + 2e–
0
+2
The half equation for the reduction is:
+
–
NO3 (aq) + 4H (aq) + 3e– → NO(g) + 2H2O(l)
+5
+2
The equation for the overall reaction is:
3Cu(s) + 2NO3–(aq) + 8H+(aq) → 3Cu2+(aq) + 2NO(g) + 4H2O(l)
0
+5
+2
+2
30.8 Nitric acid of different concentrations as oxidizing agents
P. 67 / 101
The nitrogen monoxide gas evolved reacts readily
with the oxygen in air to form brown nitrogen
dioxide gas.
Brown fumes can be seen at the mouth of the
test tube.
2NO(g) + O2(g) → 2NO2(g)
colourless
(from air)
brown
SBA note
Since both NO(g) and NO2(g) are toxic, the experiment
must be carried out inside the fume cupboard.
30.8 Nitric acid of different concentrations as oxidizing agents
P. 68 / 101
air
nitrogen dioxide gas
(brown)
air
nitrogen monoxide
gas (colourless)
copper
dilute nitric acid
heat
Figure 30.7 The reaction between dilute nitric acid and
copper in a test tube.
30.8 Nitric acid of different concentrations as oxidizing agents
P. 69 / 101
Concentrated nitric acid
Nitric acid is a stronger oxidizing agent when
concentrated.
For concentrated nitric acid, the half equation for
reduction is:
–
NO3 (aq) + 2H+(aq) + e– → NO2(g) + H2O(l)
+5
+4
The oxidation number of nitrogen decreases from
+5 to +4.
30.8 Nitric acid of different concentrations as oxidizing agents
P. 70 / 101
Action on metals
Concentrated nitric acid oxidizes most metals,
even those lower than hydrogen in the E.C.S.
(except gold and platinum).
Brown fumes of nitrogen dioxide are given off at
the same time.
30.8 Nitric acid of different concentrations as oxidizing agents
P. 71 / 101
Examples
Mg(s) + 2NO3–(aq) + 4H+(aq) → Mg2+(aq) + 2NO2(g) + 2H2O(l)
0
+5
+2
+4
Cu(s) + 2NO3–(aq) + 4H+(aq) → Cu2+(aq) + 2NO2(g) + 2H2O(l)
0
+5
+2
+4
brown NO2
fumes
Figure 30.8 The reaction between
concentrated nitric acid and copper.
30.8 Nitric acid of different concentrations as oxidizing agents
P. 72 / 101
Action on non-metals
Hot concentrated nitric acid oxidizes carbon to
carbon dioxide.
Brown fumes of nitrogen dioxide are produced:
C(s) + 4HNO3(aq) → CO2(g) + 4NO2(g) + 2H2O(l)
0
+5
+4
+4
It also oxidizes sulphur to sulphur dioxide:
S(s) + 4HNO3(aq) → SO2(g) + 4NO2(g) + 2H2O(l)
0
+5
+4
+4
Learning tip
A further reaction may occur:
SO2(g) + 2HNO3(aq) → H2SO4(aq) + 2NO2(g)
30.8 Nitric acid of different concentrations as oxidizing agents
P. 73 / 101
Action on iron(II) salts
Concentrated nitric acid oxidizes iron(II) salts to
iron(III) salts:
2+
+
−
3+
3Fe (aq) + NO3 (aq) + 4H (aq) → 3Fe (aq) + NO(g) + 2H2O(l)
+2
(green)
+5
+3
+2
(yellow)
2NO(g) + O2(g) → 2NO2(g)
(from air)
Action on sulphites
Concentrated nitric acid oxidizes sulphites to
sulphates.
2−
+
2−
−
SO3 (aq) + 2H (aq) + 2NO3 (aq) → SO4 (aq) + 2NO2(g) + H2O(l)
+4
+5
Class practice 30.9
+6
+4
Experiment 30.2
Experiment 30.2
30.8 Nitric acid of different concentrations as oxidizing agents
P. 74 / 101
30.9 Concentrated sulphuric acid as an
oxidizing agent
Concentrated sulphuric acid shows the typical
properties of an acid, except towards metals.
Besides, it can react with some non-metals.
This is due to the fact that concentrated sulphuric
acid is a strong oxidizing agent, especially when
hot.
P. 75 / 101
When concentrated sulphuric acid acts as an
oxidizing agent, it is usually reduced to sulphur
dioxide.
−
2−
2H2SO4(l) + 2e → SO4 (aq) + SO2(g) + 2H2O(l)
+6
+4
The O.N. of sulphur decreases from +6 to +4.
Learning tip
Concentrated sulphuric acid neutralizes an alkali to
form salt and water. It reacts with a carbonate to form
salt, carbon dioxide and water.
30.9 Concentrated sulphuric acid as an oxidizing agent
P. 76 / 101
Action on metals
Hot concentrated sulphuric acid oxidizes all metals
(except gold and platinum) to the corresponding
sulphates and sulphur dioxide (not hydrogen).
Examples
Cu(s) + 2H2SO4(l) → CuSO4(s) + SO2(g) + 2H2O(l)
0
+6
+2
+4
Zn(s) + 2H2SO4(l) → ZnSO4(s) + SO2(g) + 2H2O(l)
0
+6
+2
+4
30.9 Concentrated sulphuric acid as an oxidizing agent
P. 77 / 101
Dilute sulphuric acid oxidizes only the metals
above copper in the reactivity series.
+
The oxidizing agent is H (aq), not H2SO4(l).
Example
+
2+
Zn(s) + 2H (aq) → Zn (aq) + H2(g)
0
+1
+2
0
30.9 Concentrated sulphuric acid as an oxidizing agent
P. 78 / 101
Action on non-metals
Hot concentrated sulphuric acid oxidizes non-metals
(such as carbon and sulphur) to their oxides:
C(s) + 2H2SO4(l) → CO2(g) + 2SO2(g) + 2H2O(l)
0
+6
+4
+4
S(s) + 2H2SO4(l) → 3SO2(g) + 2H2O(l)
0
+6
+4
Experiment 30.3
Class practice 30.10
Experiment 30.3
30.9 Concentrated sulphuric acid as an oxidizing agent
P. 79 / 101
30.10 Sulphite ion as a reducing agent
Sulphur dioxide dissolves in water to give a
solution of sulphurous acid, H2SO3(aq).
SO2(g) + H2O(l)
H2SO3(aq)
Sulphurous acid is a weak acid which ionizes to
2–
give sulphite ions (SO3 (aq)).
H2SO3(aq)
+
2–
2H (aq) + SO3 (aq)
P. 80 / 101
Sulphite ion is a strong reducing agent.
It can be oxidized to sulphate ion.
The O.N. of sulphur increases from +4 to +6.
SO3
2–(aq)
2–
+ H2O(l) → SO4 (aq) +
+4
2H+(aq)
+6
30.10 Sulphite ion as a reducing agent
P. 81 / 101
+ 2e
–
Sulphite ions can react with:
(a) bromine, which is reduced to bromide ions:
–
–
Br2(aq) + 2e → 2Br (aq)
yellowish brown
colourless
aqueous sulphur
dioxide
bromine water
Figure 30.9 The reaction between aqueous sulphur dioxide with bromine water.
30.10 Sulphite ion as a reducing agent
P. 82 / 101
(b) acidified potassium dichromate solution, in which
dichromate ions are reduced to chromium(III) ions:
2–
+
Cr2O7 (aq) + 14H (aq) +
orange
6e–
3+
→ 2Cr (aq) + 7H2O(l)
green
aqueous sulphur
dioxide
acidified potassium
dichromate solution
Figure 30.10 The reaction between aqueous sulphur dioxide with acidified
potassium dichromate solution.
30.10 Sulphite ion as a reducing agent
P. 83 / 101
(c) acidified potassium permanganate solution, in
which permanganate ions are reduced to
manganese(II) ions:
+
MnO4–(aq) + 8H (aq) + 5e– → Mn2+(aq) + 4H2O(l)
purple
colourless
aqueous sulphur
dioxide
acidified potassium
permanganate solution
Figure 30.11 The reaction between aqueous sulphur dioxide with acidified
potassium permanganate solution.
Class practice 30.11
Example 30.4
Class practice 30.12
30.10 Sulphite ion as a reducing agent
P. 84 / 101
Key terms
1.
2.
3.
4.
5.
6.
7.
disproportionation 歧化作用
oxidation 氧化作用
oxidation number 氧化數
oxidizing agent 氧化劑
redox reaction 氧化還原反應
reducing agent 還原劑
reduction 還原作用
P. 85 / 101
Progress check
1. What is a redox reaction?
2. How is redox defined in terms of gain or loss of
oxygen?
3. How is redox defined in terms of gain or loss of
hydrogen?
4. How is redox defined in terms of gain or loss of
electrons?
5. What is the concept of oxidation number?
6. How is redox defined in terms of oxidation
number?
P. 86 / 101
7. What are the examples of common oxidizing
agents? How do they change when undergoing
reduction?
8. What are the examples of common reducing
agents? How do they change when undergoing
oxidation?
9. How can we balance redox equations by using
half equation method?
10. How can we balance redox equations by using
oxidation number method?
11. What is the relationship between position of a
species in the Electrochemical Series and its
oxidizing/reducing power?
Progress check
P. 87 / 101
12.How is the reducing power of metals related to
their positions in the Electrochemical Series?
13. How is the oxidizing power of metal ions related
to their positions in the Electrochemical Series?
14. What are the uses of the Electrochemical Series?
15. What are the reactions of chlorine with (i)
potassium bromide solution, (ii) potassium iodide
solution and (iii) sodium hydroxide solution?
16. What is a disproportionation reaction?
17. What are the reactions of nitric acid of different
concentrations with (i) metals, (ii) non-metals and
(iii) some common reducing agents?
Progress check
P. 88 / 101
18. What are the reactions of concentrated sulphuric
acid with (i) metals and (ii) non-metals?
19. What are the reactions of sulphite ions with some
common oxidizing agents?
Progress check
P. 89 / 101
Summary
30.1 Oxidation and reduction
1.
The oxidation-reduction reaction or redox
reaction, is an important type of chemical
reaction. Oxidation and reduction must occur
together. One cannot take place without the
other.
P. 90 / 101
30.2 Oxidation and reduction in terms of changes
in oxidation numbers
2.
The oxidation number of an element in a
compound is the charge an atom of the element
would have if the atom existed as an ion.
3.
There are different definitions of redox reaction,
oxidation, reduction, oxidizing agent and
reducing agent. Refer to Table 30.4 on p.12 for
the definitions.
Summary
P. 91 / 101
30.3 Common oxidizing agents and reducing
agents
4.
A strong oxidizing agent can oxidize most
reducing agents. See Table 30.5 on p.13 for
examples of common oxidizing agents.
5.
A strong reducing agent can reduce most
oxidizing agents. See Table 30.6 on p.14 for
examples of common reducing agents.
Summary
P. 92 / 101
30.4 Balancing redox equations
6.
Redox equations can be balanced by using half
equation method or oxidation number method.
See ‘Problem-solving strategy 30.2’ on p.15
and ‘Problem-solving strategy 30.3’ on p.16.
Summary
P. 93 / 101
30.5 The Electrochemical Series and the relative
strength of oxidizing agents/reducing agents
7.
In the Electrochemical Series (see Table 30.7
on p.18), oxidizing agents are arranged on the
left, and reducing agents on the right. Oxidizing
power of oxidizing agents increases down the
series. Reducing power of reducing agents
increases up the series.
8.
A metal higher in the Electrochemical Series is
a stronger reducing agent and its ion is a
weaker oxidizing agent.
Summary
P. 94 / 101
30.6 Uses of the Electrochemical Series
9.
The Electrochemical Series is very useful in
chemistry. It can be used to predict the
feasibility of a reaction. It can also be used to
predict cell reaction and direction of electron
flow in a chemical cell.
30.7 Chlorine as an oxidizing agent
10.
Chlorine is a strong oxidizing agent. It is usually
reduced to chloride ion.
Summary
P. 95 / 101
11.
The reaction in which a species is
simultaneously reduced and oxidized is called
disproportionation.
30.8 Nitric acid of different concentrations as
oxidizing agents
12.
Dilute nitric acid is an oxidizing agent. It is
reduced to NO(g) which further reacts with O2(g)
in air to form NO2(g).
–
+
–
NO3 (aq) + 4H (aq) + 3e → NO(g) + 2H2O(l)
2NO(g) + O2(g) → 2NO2(g)
colourless
(from air)
brown
Summary
P. 96 / 101
13.
Concentrated nitric acid is a strong oxidizing
agent. It is reduced to NO2(g).
–
+
–
NO3 (aq) + 2H (aq) + e → NO2(g) + H2O(l)
brown
30.9 Concentrated sulphuric acid as an oxidizing
agent
14.
Concentrated sulphuric acid is a strong
oxidizing agent. It is usually reduced to sulphur
dioxide.
–
2–
2H2SO4(l) + 2e → SO4 (aq) + SO2(g) + 2H2O(l)
Summary
P. 97 / 101
30.10 Sulphite ion as a reducing agent
15.
Sulphite ion is a strong reducing agent. It can
be oxidized to sulphate ion.
2–
2–
+
SO3 (aq) + H2O(l) → SO4 (aq) + 2H (aq) + 2e
Summary
P. 98 / 101
–
Concept map
REDOX REACTIONS
Reduction
Oxidation
definitions
definitions
•
•
•
•
Loss of oxygen
Gain of hydrogen
Gain of electrons
____________
in
Decrease
oxidation number
•
•
•
•
Gain of oxygen
Loss of hydrogen
Loss of electrons
Increase
____________
in
oxidation number
P. 99 / 101
Reduction
undergoes
Oxidizing agent
examples
• _____________
Chlorine
• Dilute nitric acid
• Concentrated
nitric acid
• _____________
Concentrated
_____________
nitric acid
Oxidation
undergoes
Reducing agent
example
Sulphite ion
Concept map
P. 100 / 101
Oxidizing agent
Reducing agent
form
Electrochemical Series
uses
Concept map
• predict the feasibility
of a redox reaction
• predict cell reaction
and direction of
_____________
electron flow in
a chemical cell
P. 101 / 101