Chemistry 30
Unit 7:
REDOX Reactions
A Simple Redox Experiment
When a copper wire is placed in an aqueous solution of silver nitrate, a spectacular reaction results.
Within minutes a large number of shiny metallic crystals form on the wire, and the solution turns pale
blue. When we shake off the crystals, we observe that much of the copper wire has disappeared. The
crystals are silver crystals and the blue colour of the solution indicates the presence of hydrated
copper ions. Copper metal has displaced silver from its salt.
 Cu(NO3)2(aq) + 2 Ag(s)
Cu(s) + 2 AgNO3(aq) 
 Cu2+(aq) + 2 Ag(s)
Cu(s) + 2 Ag+(aq) 
Copper atoms in the wire each lose two electrons and go into solution as hydrated copper ions. The
electrons are transferred to hydrated silver ions, which crystallize as metallic silver.
Any reaction in which electrons are lost is called an oxidation reaction. A reaction that involves the
gain of electrons is called a reduction reaction. In the copper-silver nitrate reaction, the copper metal
is oxidized to copper ions and the silver ions are reduced to silver metal. Because oxidation and
reduction reactions always take place concurrently, we refer to these reactions as redox reactions.
Electronegativity and Oxidation Numbers





The oxidation number tells us how many electrons an atom in a compound would ______________
or _________________________.
In assigning oxidation numbers we assume that the electrons of each bond are transferred to
the _________________________ electronegative atom in the bond.
A positive oxidation number indicates a _________________________ of electrons and a negative
number indicates a _________________________ of electrons.
The oxidation number of an atom in its pure form, as an element, is _________________________.
The oxidation number of a monatomic ion is equal to the charge on the _________________________,
since the charge indicates the number of electrons lost or gained.
Example
What are the oxidation numbers of each element in calcium chloride?
Example
What are the oxidation numbers of iron in the two iron oxides, FeO and Fe2O3? Assume that oxygen
has an oxidation number of −2 in each case.
The sum of the oxidation numbers in each compound will be _________________________.
-2-
Oxidation Number Rules
1. The oxidation number of an atom in its pure form, as an element, is _________________________.
2. The oxidation number of a monatomic ion is equal to the _________________________ on the ion.
3. The algebraic sum of the oxidation numbers in a neutral polyatomic compound is _______________.
4. The algebraic sum of the oxidation numbers in a polyatomic ion is equal to the
_________________________ on the ion.
5. The most common oxidation number for hydrogen is ___________; in metal hydrides it is __________.

In most of its compound, hydrogen has the lowest electronegativity value and an oxidation
number of +1. But in _________________________ hydrides, such as sodium hydride, NaH, the
hydrogen is bonded to a less electronegative metal and has an oxidation number of −1.
6. The most common oxidation number for oxygen is ________; in peroxides it is _________, and in
combination with fluorine it is +2.
7. In combinations of nonmetals, the oxidation number of the more electronegative atom is
_________________________. The oxidation number of the less electronegative atom is
_________________________.
Example
H
He
2.01
-
What are the oxidation numbers of oxygen and
fluorine in oxygen difluoride, OF2?
Li
Be
0.98 1.57
2.04 2.55 3.04 3.44 3.98 -
The electronegativity values of oxygen and
fluorine are __________ and __________respectively.
The fluorine atoms attract electrons more
strongly than oxygen; therefore it is assigned
the _________________________ oxidation number,
and the oxygen, the positive oxidation number.
Na
Al
Mg
B
C
Si
P
O
S
F
Cl
Ne
Ar
0.93 1.31
1.61 1.90 2.19 2.58 3.16 -
K
Ga
Ca
Ge
As
Se
Br
Kr
0.82 1.00
1.81 2.01 2.18 2.55 2.96 3.0
Rb
In
Sr
Sn
Sb
Te
I
Xe
0.82 0.95
1.78 1.96 2.05 2.1
2.66 2.6
Cs
Tl
Po
At
Rn
2.04 2.33 2.02 2.0
2.2
-
Ba
0.79 0.89
Pb
Example
What is the oxidation number of carbon in each of the following compounds:
a) ethane, C2H6
N
b) ethyne, C2H2
-3-
Bi
Oxidation Numbers of Atoms in Polyatomic Ions or Molecules
Example
What are the oxidation numbers of the elements in the fluorosulfate ion, FSO3−?
Example
What are the oxidation numbers of each of the elements in sodium perxenate, Na4XeO6?
Example
What are the oxidation numbers of each of the elements in ammonium carbonate, (NH4)2CO3?
Determine Oxidation Numbers Assign
Redox Reactions





An increase in the oxidation number of an atom indicates that it has _________________________
electrons and has undergone _________________________.
A decrease in the oxidation number of an atom means an _________________________ in the
number of electrons and indicates _________________________.
Since electrons are neither created nor destroyed in chemical reactions, the total number of
electrons will always remain unchanged. This means that the loss of electrons must always
be accompanied by the _________________________ process – the gain of electrons. In other
words, oxidation and reduction are processes that always take place together.
The element that undergoes oxidation is the _________________________ agent.
The element that undergoes reduction is the _________________________ agent.
-4-
Example
Which of the following reactions is a redox reaction?
 2 NaCl(aq) + Br2(aq)
a) 2 NaBr(aq) + Cl2(aq) 
 2 NaCl(aq) + H2S(g)
b) 2 HCl(aq) + Na2S(aq) 
Activity Series of Metals
Example
Use the activity series to predict whether a
reaction takes place when the following
substances are mixed. If there is a reaction,
complete and balance the equation, and
then write the net ionic equation.

a) Fe(s) + ZnSO4(aq) 

b) Al(s) + H2SO4(aq) 

c) Sn(s) + CuSO4(aq) 
The Activity Series
Lithium
Potassium
These metals displace hydrogen from water
Barium
Calcium
Sodium
Magnesium
Aluminium
Zinc
Chromium
These metals displace hydrogen from acids.
Iron
Cadmium
Nickel
Tin
Lead
Hydrogen
Copper
These metals do not displace hydrogen from
Mercury
Acids or water.
Silver
Gold
-5-
Oxidizing and Reducing Agents





A good oxidizing agent must be a good electron _________________________ and a good reducing
agent must be a good electron _________________________.
The cheapest and most widely-used oxidizing agent is atmospheric _________________________.
Elemental chlorine is also a good _________________________ agent.
The most commonly-used reducing agent is _________________________ gas.
Reactive metals such as magnesium and aluminium are also good _________________________ agents.
Example
Consider the addition or removal of oxygen atoms, as well as the change in oxidation numbers, to
determine which substance is oxidized and which is reduced in the following reaction:

CH3OH(aq) + 2 HClO(aq) 
HCO2H(aq) + 2 HCl(aq) + H2O(l)
methanol
formic acid
hypochlorous acid
Balancing Redox Equations Using Oxidation Numbers
In a balanced redox equation, the total _________________________ in oxidation number of the element that
is oxidized must equal the total _________________________ in oxidation number of the element that is
reduced.
Procedure:
1. Assign _________________________ numbers to all the atoms in the equation.
2. Identify which atoms undergo a _________________________ in oxidation number.
3. Determine the _________________________ in which these atoms must react so that the total increase
in oxidation numbers equals the decrease.
4. Balance the _________________________ participants in the equation.
5. Balance the other atoms by the _________________________ method.
Example: Balance the following redox reaction
HNO3 + H2S

 NO + S + H2O
Step 1:
Step 2:
Step 3:
Step 4:
Step 5:
-2-
Example: Balance the following redox reaction
 Mn2+ + HSO4− + H2O
MnO4− + H2SO3 + H+ 
Step 1:
Step 2:
Step 3:
Step 4:
Step 5:
Examples
a)
 HIO3 + NO2 + H2O
I2 + HNO3 
b)
 Br2 + H2O
HBr + HBrO3 
Balancing Using Oxidation #s Assign
-3-
Half-Reactions
By separating a redox reaction into two half-reactions, we can get a better understanding of how
redox reactions take place.
 MgCl2(s)
Mg(s) + Cl2(g) 
 Cu(NO3)2(aq) + 2 Ag(s)
Cu(s) + 2 AgNO3(aq) 
Example
Write the half-reactions for the reaction:
 2 AlCl3(aq) + 3 H2(g)
2 Al(s) + 6 HCl(aq) 
Ion-Electron Method
We can use half-reactions to _________________________ redox equations. When we do so, we are using the
ion-electron method.
1. Write balanced oxidation and reduction __________________________________________________.
2. Change the coefficients in the balanced half-reactions so that the number of
_________________________ produced in the oxidation half-reaction equals the number of
electrons consumed in the reduction half-reaction.
3. Add the two half-reactions to obtain a _________________________ net ionic equation for the
redox reaction.
-4-
In acidic solution: To write a balanced half-reaction for the oxidation of ethanol, C2H5OH to acetic
acid, CH3CO2H, in acidic solution, we would follow the steps outlined below.
1. Write the _________________________ equation:
 CH3CO2H
C2H5OH 
2. Balance for species other than oxygen and hydrogen:
 CH3CO2H (no change)
C2H5OH 
3. Balance for oxygen using one _________________________ molecule for each oxygen you
require:
4. Balance for hydrogen using a hydrogen _______________ for each hydrogen you require:
5. Balance for charge by adding _________________________ to either the product or reactant
side:
In basic solution: We will use the reduction of the permanganate ion to manganese(IV) oxide in basic
solution as an example.
1. Write the skeletal equation:
 MnO2
MnO4− 
2. Balance for species other than oxygen and hydrogen:
 MnO2 (no change)
MnO4− 
3. Balance for oxygen atoms by adding one water molecule for each oxygen
that you require:
4. Balance for hydrogen atoms by adding one hydrogen ion for every
hydrogen atom:
5. For every hydrogen ion add one hydroxide ion to both sides of the equation:
6. Combine hydrogen and hydroxide ions to form water:
7. Balance for charge by adding electrons:
8. Finally, cancel water molecules:
-5-
Example
Write a balanced half-reaction for the oxidation of aluminum metal to the aluminate ion (AlO2−) in
basic solution.
 AlO2−(aq)
Al(s) 
Balancing Half Reactions Assign
Combining Half-Reactions
We can combine the half-reactions for the reduction of iron(III) to iron(II) with the oxidation of
hydrogen gas to hydrogen ion:
The Full Ion-Electron Method
A reaction between the purple permanganate ion and iron(II) ion in acid solution to produce the
colourless manganese(II) ion and the iron (III) ion.
 Mn2+ + Fe3+
MnO4− + Fe2+ 
-6-
Example
Use the ion-electron method to write a balanced equation for the reaction between copper metal and
concentrated nitric acid to produce copper(II) ions and nitrogen dioxide gas.
See Ion Electron Assign
Electrochemical Cells



An electrochemical cell uses energy released from a spontaneous redox reaction to generate an
_________________________ current.
The current is derived from the flow of electrons through _________________________, called metallic
conduction, and the movements of ions in solution, called electrolytic _________________________.
A battery consists of a single electrochemical cell or a number of cells connected in
_________________________.
How Electrochemical Cells Work
When a piece of zinc metal is placed in an aqueous solution of copper(II) sulfate, a rapid redox
reaction takes place. Electrons are _________________________ from the zinc metal to the copper ions. The
reaction is spontaneous since zinc oxidizes more readily than copper. It appears above copper in the
activity series:
 Zn2+(aq) + Cu(s)
Zn(s) + Cu2+(aq) 
Oxidation:
Reduction:

The electrical circuit of an electrochemical cell consists of two parts, an _________________________
circuit and an _________________________ circuit.
-7-
_________________________: an electrically-conducting medium, usually a solid, that is used to make
electrical contact
_________________________: the site of oxidation and electron release
_________________________: the site of reduction and electron consumption
__________________________________________________: the force that pushes the electrons from the anode to the
cathode (a.k.a. emf)
_________________________: a unit used to measure potential difference
___________________________________: a device filled with a salt solution that connects the half-cells to allow
the passage of ions and to maintain the electrical neutrality of the two solutions
As the reaction continues, the voltage decreases and finally becomes zero when the reaction reaches
equilibrium. The cell will then be “dead.”
Notation for an Electrochemical Cell
By convention the anode appears on the left, the cathode on the right. The single vertical
lines indicate the contact boundaries between phases in each half-cell. The double
vertical lines represent the _________________________or separator. An electrochemical cell is
referred to as a standard cell when the reactant and product ions or molecules are
present in concentrations of 1.0 M.
-8-
Example
When the metal electrodes of an electrochemical cell are connected by a voltmeter, the meter indicates
a potential difference of 0.46 V.
a)
using the activity series, identify the half-reaction taking place at each electrode.
b)
Write the shorthand notation for the cell, identifying the anode and cathode, and
indicate the direction of movement of electrons and ions.
See Cell Assign
-9-
Standard Electrode Potentials


An almost unlimited number of half-cells can be constructed and combined into
electrochemical cells. However, the reduction potential of each half-cell, the tendency of its ions
and molecules to gain electrons, will be _________________________. As a result, the electrodes will
have different electrical potentials.
Unfortunately we cannot measure the electrical potential of one half-cell. We can only measure
the voltage difference, or emf, between _________________________ half-cells by combining them into
an electrochemical cell. For example, a cell voltage of 1.10 V in the standard zinc-copper ion cell
tells us that the reduction potential below zinc in the activity series. It also means that the
oxidation potential (the tendency to lose electrons) is 1.10 V greater for _________________________
than copper.
Standard Half-cell Potentials

The standard hydrogen electrode has been assigned a standard potential (E °) of 0.00 V.
 H2(g)
2 H+(aq) + 2 e− 
1M



E ° = 0.00 V
100 kPa
T = 25°C
This electrode is then used as a reference to assign potentials to other half-cells. The measured
electromotive force of a standard electrochemical cell containing a _________________________
half-cell is the standard potential (E °) of the second half-cell.
The standard reduction potential of a half-cell is a measure of the tendency to
_________________________ electrons from the hydrogen half-cell.
Ions or molecules with negative reduction potentials gain electrons less easily than hydrogen
ions. Those with positive values gain electrons more easily than hydrogen ions.
o Lithium ions in aqueous solution have the weakest tendency to gain electrons (E ° =
−3.05 V); molecular fluorine has the greatest potential to gain electrons (E ° = +2.87 V).
Calculating Standard Cell Potentials

In an electrochemical cell, one half-cell undergoes _________________________ while the other
undergoes _________________________. Thus when we are combining half-reactions and calculating
the cell potential, one of the tabulated reduction equations will have to be reversed and the sign
of the potential changed. Let us find the potential for the following standard cell:
Fe( s ) Fe2 ( aq ) 1 M Cu 2 ( aq ) (1 M ) Cu( s )
From table:
 Fe(s)
Fe2+(aq) + 2 e− 
E°=
 Cu(s)
Cu2+(aq) + 2 e− 
E°=
The Fe 2 Fe half-cell is the _________________________. Thus we _________________________ the Fe 2 Fe
standard half-cell reduction reaction and the sign of E ° to give:
Anode (oxidation)
Cathode (reduction)
The sum of these half-cell potentials provides the potential of the cell.
- 10 -
Example
Calculate the potential of the standard cell:
Zn( s ) Zn 2 ( aq ) (1 M ) Cl  ( aq ) (1 M ) Cl2( g ) (100 kPa), Pt( s )
First we determine the half-reactions from the shorthand notation of the cell. We reverse the Zn Zn2+
standard half-cell reduction and the sign of E ° given in table thus,
Predicting Spontaneity of Redox Reactions
Reduction potentials can also be used to predict whether a redox reaction is spontaneous or
nonspontaneous. A redox reaction is spontaneous when the sum of the potentials of the half-reactions
is _________________________. The reaction is nonspontaneous when the sum is negative.
Example
Predict whether sulfur dioxide and bromine will react in aqueous solution given the equation:
 SO42−(aq) + 2 Br−(aq) + 4 H+(aq)
SO2(g) + Br2(aq) + 2 H2O(l) 
Half Cell Potentials and Spontaneity Assign
- 11 -
Standard Reduction Potentials at 25oC*
 4 OH-(aq)
O2(g) + 2 H2 + 4 e- 
+0.40
Half-Reaction
E0 (V)
 2 I-(aq)
I2(s) + 2 e- 
+0.53
 Li(s)
Li+(aq) + e- 
-3.05
+0.59
 K(s)
K+(aq) + e- 
-2.93
 MnO2(s) +
MnO4-(aq) + 2 H2O + 3 e- 
4 OH (aq)
 Ba(s)
Ba2+(aq) + 2 e- 
-2.90
 H2O2(aq)
O2(g) + 2 H+(aq) + 2 e- 
+0.68
 Sr(s)
Sr2+(aq) + 2 e- 
-2.89
 Fe2+(aq)
Fe3+(aq) + e- 
+0.77
 Ca(s)
Ca2+(aq) + 2 e- 
-2.87
 Ag(s)
Ag+(aq) + e- 
+0.80
 Na(s)
Na+(aq) + e- 
-2.71
 2 Hg(l)
Hg22+(aq) + 2 e- 
+0.85
 Mg(s)
Mg2+(aq) + 2 e- 
-2.37
 Hg22+(aq)
2 Hg2+(aq) + 2 e- 
+0.92
 Be(s)
Be2+(aq) + 2 e- 
-1.85
+0.96
 Al(s)
Al3+(aq) + 3 e- 
-1.66
 NO(g) + 2
NO3-(aq) + 4 H+(aq) + 3 e- 
H2O
 2 Br-(aq)
Br2(l) + 2 e- 
+1.07
 2 H2O
O2(g) + 4 H+(aq) + 4 e- 
+1.23
 Mn2+(aq) +
MnO2(s) + 4 H+(aq) + 2 e- 
2 H2O
+1.2`3
+1.33
Mn2+(aq)
+2
2 H2O + 2
e-
e-
-1.18

 Mn(s)

 H2(g) + 2
OH-(aq)
-0.83

 Zn(s)
-0.76
 Cr(s)
Cr3+(aq) + 3 e- 
-0.74
 Fe(s)
Fe2+(aq) + 2 e- 
-0.44
2
Cr2O72-(aq) + 14 H+(aq) + 6 e- 
Cr3+(aq) + 7 H2O
 Cd(s)
Cd2+(aq) + 2 e- 
-0.40
 2 Cl-(aq)
Cl2(g) + 2 e- 
+1.36
 Pb(s) + SO42-(aq)
PbSO4(s) + 2 e- 
-0.31
 Au(s)
Au3+(aq) + 3 e- 
+1.50
 Co(s)
Co2+(aq) + 2 e- 
-0.28
+1.51
 Ni(s)
Ni2+(aq) + 2 e- 
-0.25
 Mn2+(aq)
MnO4-(aq) + 8 H+(aq) + 5 e- 
+ 4 H2O
+1.61

 Sn(s)
-0.14
 Ce3+(aq)
Ce4+(aq) + e- 
+1.70
 Pb(s)
Pb2+(aq) + 2 e- 
-0.13

PbO2(s) + 4 H+ + SO42-(aq) + 2 e- 
PbSO4(s) + 2 H2O
 H2(g)
2 H+(aq) + 2 e- 
0.00
 2 H2O
H2O2(aq) + 2 H+(aq) + 2 e- 
+1.77
 Sn2+(aq)
Sn4+(aq) + 2 e- 
+0.13
 Co2+(aq)
Co3+(aq) + e- 
 Cu+(aq)
Cu2+(aq) + e- 
+0.13
+0.14
 2 SO42−
S2O82− + 2 e− 
+1.82
+2.01
 O2(g) + H2O
O3(g) + 2 H+(aq) + 2 e- 
+2.07
 SO2(g) + 2
SO42-(aq) + 4 H+(aq) + 2 e- 
H 2O
+0.20
 F-(aq)
F2(g) + 2 e- 
+2.87
 Ag(s) + Cl-(aq)
AgCl(s) + e- 
+0.22
*For
 Cu(s)
Cu2+(aq) + 2 e- 
+0.34
Zn2+(aq)
Sn2+(aq)
+2
+2
e-
e-
 H2S
S + 2 H+ 2 e− 
- 12 -
all half-reactions the concentration is 1M
for dissolved species and the pressure is 1 atm
for gases. These are standard state values.
Name: ___________________________
Determining Oxidation Numbers
1. What are the oxidation numbers for each element in the following ionic compounds:
a) SnCl4
b) Ca3P2
c) SnO
d) Ag2S
2. What are the formulas of the oxides of nitrogen in which nitrogen has oxidation numbers
of +1, +2, +3, +4, and +5?
3. Determine the oxidation numbers for each element in the following substances or ions:
a)
Cu2 S
d)
Mn O42−
b)
S8
e)
Cr O42−
c)
Li Al H4
f)
N H4 Cl
4. Assign oxidation numbers to all elements in the following compounds:
a) nitrous acid
b) potassium permanganate
c) ammonium sulfate
d) sodium aluminate
e) hydronium ion
5. What is the oxidation number of carbon in each of the following compounds:
a) methane, CH4
b) formaldehyde, CH2O
c) carbon monoxide
6. What are the oxidation numbers of the following:
a)
copper in Cu2 SO4 and in Cu SO4
b)
lead in Pb Br2 and in Pb Br4
- 13 -
Name:_______________________________
Oxidation and Reduction
1. Which of the following are redox reactions?
 2 NaCl(aq) + FeS(s)
a)
Na2S(aq) + FeCl2(aq) 
b)
 2 NaOH(aq) + H2(g)
2 Na(s) + 2 H2O(l) 
c)
 2 KCl(s) + 3 O2(g)
2 KClO3(s) 
d)
 Na2 S + H2 O
H2 S + Na2 O 
e)
 Ca O + H2
Ca + H2 O 
f)
 C H3 O H
C O + 2 H2 
g)
 H2 S O4
S O3 + H2 O 
h)
 Ca S O3 + C O2
Ca C O3 + S O2 
2. Which reactant in each of the following is oxidized? Which reactant is the oxidizing
agent?
 2 HI(g)
a)
H2(g) + I2(g) 
b)
 2 F2O(g)
O2(g) + 2 F2(g) 
c)
 Mg3N2(s)
3 Mg(s) + N2(g) 
3. Identify the reducing agents in the following reactions. Which of the reactants are
reduced?
4.
a)
 CO2(g) + Pb(s)
CO(g) + PbO(s) 
b)
 2 HBr(aq) + H2SO4(aq)
Br2(aq) + SO2(g) + 2 H2O(l) 
c)
 2 NaCl(aq) + Br2(aq)
2 NaBr(aq) + Cl2(aq) 
Show that in the following balanced redox reactions the increase in oxidation numbers
equals the decrease.
a)
 2 NH3
N2 + 3 H2 
b)
 2 Fe2+ + I2
2 Fe3+ + 2 I− 
c)
 2 KCl + 3 O2
2 KClO3 
- 14 -
Name:____________________________________
Balancing Using Oxidation Numbers
1. Using the activity series, predict whether a reaction takes place, and if so, give a
balanced reaction equation.

a)
Ag(s) + HCl(aq) 
b)

Mg(s) + FeSO4(aq) 
c)

Cu(s) + AuCl3(aq) 
d)

Sn(s) + Al2(SO4)3(aq) 
2. Use the oxidation number method to balance the following equations:
a)
K2Cr2O7(aq) +
b)
HNO3(aq) +

HI(aq) 
K2CrO4(aq) +
KI(aq) +
CrI3(aq) +

Fe(NO3)2(aq) 
I2(s) +
KNO3(aq) +
c)
 KNO3(aq) +
HNO3(aq) + C2H6O(l) + K2Cr2O7(aq) 
d)
HCl(aq) +

NH4Cl(aq) + K2Cr2O7(aq) 
- 15 -
H2O(l)
KCl(aq) +
Fe(NO3)3(aq) +
Cr(NO3)3(aq) +
C2H4O(l) +
CrCl3(aq) +
H2O(l)
H2O(l) +
Cr(NO3)3(aq)
N2(g) +
H2O(l)

HBr(aq) 
e)
KClO3(aq) +
f)
Se(s) +
g)
Ce4+(aq) +
h)

I2 + HNO3 
i)
MnO4− + H+ +
Br2(l) +
H2O(l) +

H N O3(aq) 
Se O2(s) +
N O(g) +

Sn2+(aq) 
Ce3+(aq) +
Sn4+(aq)
KCl(aq)
HIO3 + NO2 + H2O

Cl− 
Mn2+ + Cl2 + H2O
j) CrO2- + ClO- + OH-  CrO42- + Cl- + H2O
- 16 -
H2 O(l)
Name:_______________________________
Balancing Half Reactions
1. Write the oxidation and reduction half-reactions for the preparation of the following ionic
compounds:
 2 NaBr(s)
2 Na(s) + Br2(l) 
a)
b)
 ZnS(s)
Zn(s) + S(s) 
c)
 2 Al2O3(s)
4 Al(s) + 3 O2(g) 
2. Write net ionic equations and half-reactions for each of the following reactions:
 Cr2(SO4)3(aq) + 2 Ag(s)
a) 2 CrSO4(aq) + Ag2SO4(aq) 
 Ca(OH)2(aq) + H2(g)
b) Ca(s) + 2 H2O(l) 
- 17 -
 2 NaCl(aq) + I2(s)
c) Cl2(g) + 2 NaI(aq) 
 MgCl2(aq) + H2(g)
d) Mg(s) + 2 HCl(aq) 
e)
 Zn(NO3)2(aq) + 2 Ag(s)
Zn(s) + 2 AgNO3(aq) 
3. Balance the following half-reactions:
 Au3+(aq)
Au+(aq) 
a)
b)
 S(s)
S2−(aq) 
c)
 Ag(s) + OH−(aq)
Ag2O(s) + H2O(l) 
d)
 BrO3−(aq) + H+(aq)
Br2(l) + H2O(l) 
- 18 -
4.
Write balanced equations for the following half-reactions in acidic aqueous solution:
a)
the reduction of sulfurous acid, H2SO3, to elemental sulfur
b)
the oxidation of formic acid, HCO2H, to carbon dioxide
c)
the conversion of lead(IV) oxide into lead(II) ions
5. The following half-reactions proceed in basic aqueous solution. Write a balanced
reaction equation for each.
a)
the reduction of sulfur dioxide to elemental sulfur
b)
the conversion of methane into carbon dioxide
c)
the oxidation of nickel into nickel(II) oxide
- 19 -
Name:_______________________________________
Ion Electron Method
1. Complete and balance the following redox reactions by the ion-electron method.
a)
 Ag+(aq) + Cr3+(aq); acidic solution
Ag(s) + Cr2O72−(aq) 
b)
 MnO2(s) + NO3−(aq): basic solution
MnO4−(aq) + NO2−(aq) 
c)
 CO2(g) + Mn2+(aq); acidic solution
CH3OH(aq) + MnO4−(aq) 
d)
 Ce3+(aq) + IO3−(aq); basic solution
Ce4+(aq) + I−(aq) 
e)
 O2(g) + Cl2(g); acidic solution
H2O2(aq) + ClO3−(aq) 
f)
 Ni(OH)2(s) + SO32−(aq); basic solution
NiO2(s) + S2O32−(aq) 
- 20 -
2. Balance the reactions by the ion-electron method. Reactions take place in acidic aqueous
solutions.
 CO2 + I−
a)
H2C2O4 + IO3− 
b)
 NO3− + Cr3+
NO + Cr2O72− 
3. Write a balanced equation for the reaction in which zinc metal is oxidized to zinc(II) ion
by a dilute acidic solution containing nitrate ion. Dinitrogen monoxide gas is also formed.
4. Balance the reactions by the ion-electron method. Reactions take place in basic aqueous
solutions.
 MnO2 + N2
a)
MnO4− + N2H4 
b)
 CrO42− + I−
Cr(OH)3 + IO3− 
- 21 -
Cells
Name:____________________________
1. An electrochemical cell consists of two half-cells connected by a salt bridge. One half
consists of an aluminum electrode placed in a 1.0 M aluminum ion solution. The other
half-cell contains a nickel electrode in a 1.0 M nickel ion solution. The salt bridge contains
sodium sulfate.
a)
from the activity series determine the reaction taking place in each half-cell
b)
write a shorthand notation for the cell
c)
make a sketch of the cell marking in the anode and cathode, and indicate the
direction of electron and ion movement
Name:__________________________________
- 22 -
Half Cell Potentials and Spontaneity
1. Write the anode and cathode half reactions for each of the following electrochemical cells and
calculate the cell potential (E °) for each cell. In b), c) and d) the electrodes are platinum.
a)
Cr( s ) Cr 3 ( aq ) (1 M ) Ag  ( aq ) (1 M ) Ag ( s )
b)
Pb( s ) Pb 2  ( aq ) (1 M ) Cl  ( aq ) (1 M ) Cl2( g ) , Pt( s )
c)
Pt( s ) , H 2( g ) OH  ( aq ) (1 M ) OH  ( aq ) (1 M ) O2( g ) , Pt( s )
d)
Fe( s ) Fe 2 ( aq ) (1 M ) Cl  ( aq ) (1 M ) Cl2( g ) (100 kPa), Pt( s )
2. Use half-reaction potentials to predict whether the following reactions are spontaneous or
nonspontaneous in aqueous solutions.
a)
 Sn(s) + O2(g) + 2 H+(aq)
H2O2(aq) + Sn2+(aq) 
b)
 2 I−(aq) + Cl2(g)
I2(aq) + 2 Cl−(aq) 
c)
 Mn2+(aq) +
MnO4−(aq) + 5 Br−(aq) + 8 H+(aq) 
d)
 2 H+(aq) + 2 SO42−(aq)
H2(g) + S2O82−(aq) 
- 23 -
5
Br2(g) + 4 H2O(l)
2