Catalyst
 Use the following standard free energy of formation for formic
acid (HCO2). Calculate the Ka for the reaction:
ΔGfo (kJ/mol)
HCO2
-372.3
H+
0
CO2-
-351.0
End
Lecture 8.4 – Balancing Redox Reaction
and Galvonic Cells
Today’s Learning Targets
 LT 8.9 – I can identify the compounds that are being oxidized and reduced in a




chemical reaction.
LT 8.10 – I can balance a redox reaction in acidic and basic solutions by utilizing the
half reactions for a given oxidation-reduction reaction.
LT 8.11 – I can draw a voltaic (galvanic) electrochemical cell for a given redox
reaction. In this drawing, I can identify, the anode, cathode, salt bridge, direction of
election movement, and placement of oxidation and reduction half reactions.
LT 8.12 – I can calculate the standard reduction potential for an electrochemical cell
through utilization of a reduction potential table.
LT 8.13 – I can identify an oxidizing agents and reducing agents in a chemical
reaction and utilizing a reduction potential table.
Oxidation – Reduction Reactions
 Oxidation-Reduction reactions (or redox reactions) are
reactions that involve the transfer of electrons from one
element to another element
 A substance either gives or receives electrons
LEO the lion says GER
Lose
Electrons
Oxidized
Gain
Electrons
Reduced
Assigning Oxidation Numbers
 In order to identify substances that are gaining or losing electrons,
we assign oxidation numbers.
1. An atoms in element form have a state of 0
2. A monoatomic ions oxidation number equals its charge (e.g. K+ =
+1 oxidation number). This also applies to monoatomic ionic
compounds.
3. Non-metals have the charge of their group usually:



Oxygen is always -2 except in peroxides (O22-) where it is -1
Hydrogen is usually +1 when bonded to nonmetals and -1 when
bonded to metals
Halogens almost always are -1. Fluorine is always -1, all other
halogens may have a positive charge
4. The sum of the oxidation numbers in a compound equal the
charge of the compound.
Class Example
 Assign the oxidation states to the following elements/compounds
and identify the substances being oxidized and reduced:
2 Al (s) + 6 HBr (aq)  2 AlBr3 (aq) + 3 H2 (g)
Table Talk
 Assign the oxidation states to the following elements/compounds
and identify the substances being oxidized and reduced:
Mg (s) + CoSO4 (aq)  MgSO4 (aq) + Co (s)
Steps to Balancing Redox Reactions
Divide each reaction into two ½ reactions
2. Balance each ½ reaction
1. First balance elements other than H and O
2. Balance O by adding H2O as needed
3. Balance H by adding H+ as needed
4. Finally, balance chare by adding e- as needed
3. Multiply each half reaction so that there are equal number of
electrons in each half reaction
4. Sum the reactions together.
1.
Class Example
 Balance the following reaction for both mass and charge in acidic
solution:
MnO4- (aq) + C2O42- (aq)  Mn2+(aq) + CO2 (aq)
Table Talk
 Balance the following reaction for both mass and charge in acidic
solutions:
Cr2O72- (aq) + Cl- (aq)  Cr3+ (aq) + Cl2 (g)
Table Talk
 Balance the following reaction for both mass and charge in acidic
solutions:
Cu (s) + NO3- (aq)  Cu2+ (aq) + NO2 (g)
Balancing in Basic Solution



By adding hydrogen to each side, we create an acidic solution
If the redox reaction occurs in basic solution, add equal
amounts of OH- on each side.
The OH- reacts with the H+ to make water
Class Example
 Balance the following reaction for both mass and charge in basic
solution:
CN- (aq) + MnO4- (aq)  CNO- (aq) + MnO2 (s)
Table Talk
 Balance the following reaction for both mass and charge in basic
solutions:
NO2- (aq) + Al (s)  NH3 (aq) + Al(OH)4- (aq)
Table Talk
 Balance the following reaction for both mass and charge in basic
solutions:
Cr(OH)3 (s) + ClO- (aq)  CrO42- (aq) + Cl2 (g)
Catalyst
 Balance the following reaction for both mass and charge in acidic
solutions:
Cr2O72- (aq) + Cl- (aq)  Cr3+ (aq) + Cl2 (g)
End
Fill the Box!
 There are nine questions around the room
 Balance the reactions and fill in the handout as you complete the
activity.
Voltaic Cells
 The energy released in a spontaneous redox reaction can be used
to perform electrical work.
 This is done through the construction of a voltaic cell
 In a voltaic cell, the two half reactions are split and they are not in
direct contact with one another
 The movement of electrons between the two cells can be utilized
to do work
Important Parts of Voltaic Cells
 Anode – The site where oxidation takes place. Negatively
charged.
 Cathode – The site where reduction takes place. Positively
charged.
 Salt Bridge – Connects the two ½ cells and contains an
unreactive salt. This ensures that charge does not build up in the
½ cells
 Common salts solutions are NaNO3 or KNO3
Class Example
 The oxidation – reduction reaction:
Cu2+ (aq) + Zn (s)  Zn2+ (aq) + Cu (s)
 Draw the setup for the voltaic cell that can be created in order to
do work. Assume KNO3 is utilized for the salt bridge. Identify, the
anode, cathode, salt bridge, the charge of the two electrodes,
movement of elections, and write the voltaic cell in cell notation.
Table Talk
 The oxidation – reduction reaction:
ClO3- (aq) + 6 H+ (aq) + 3 Zn (s)  3Zn2+ (aq) + Cl- (aq) + 3 H2O (l)
 Draw the setup for the voltaic cell that can be created in order to
do work. Assume KNO3 is utilized for the salt bridge. Identify, the
anode, cathode, salt bridge, the charge of the two electrodes,
movement of elections, and write the voltaic cell in cell notation.
Cell Potentials
 In a voltaic cell, electrons move from an area of high potential
energy to an area of low potential energy
 Differences in potential energy are measured in volts
J
1V = 1
C
 The potential difference between two cells is the cell potential
(Ecell)
 When the cell is run under standard conditions, it is called the
cell potential (Eocell)
standard
Standard Reduction Potentials
 We measure numerous half cells under standard conditions (25
oC, 1
M, and 1 atm)
 We can calculate the standard reduction potential (Eocell) for the
electrochemical cell by utilizing:
o
o
E ocell  E cathode
 E anode
 Therefore, we can calculate the electrochemical potential for a
voltaic cell simply by using a reference table

Insert Table 20.1
Class Example
 A voltaic cell is set up for the following reaction:
Zn (s) + Cu2+ (aq, 1 M)  Cu (s) + Zn2+ (aq, 1 M)
 Using Table 1, calculate the Eocell for this voltaic cell
Table Talk
 You set up the following voltaic cell:
Cr2O72- (aq) + 14 H+ (aq) + 6 I- (aq)  Cr3+ (aq) + 3 I2 (s) + 7 H2O (l)
 Calculate the Eocell for the voltaic cell
Oxidizing and Reducing Agents
 The substance that causes another object to be oxidized is known
as the oxidizing agent (it itself it reduced)
 The substance that causes another object to be reduced is known
as the reducing agent (it itself it oxidized)
 The higher you are on the reduction potential table, the better an
oxidizing agent you are
 The lower you are, the better a reducing agent you are when the
reaction is reversed
Better reducing
agent!
Better oxidizing
agent!
Insert Table 20.1
Class Example
 Using your reduction potential table, rank the following ions in
order of increasing strength as oxidizing agents:
NO3-, Ag+, Cr2O72-
Table Talk
 Using your reduction potential table, rank the following ions in
order of increasing strength as reducing agents:
I-, Fe, Al
White Board Problems
White Board Problems
 You create a voltaic cell with the following reaction:






Fe (s) + 2 Ag+ (aq)  Fe2+ (aq) + 2 Ag (s)
Create a diagram of the two half cells. Label the cathode, anode,
movement of electrons, salt bridge, and write in cell notation.
Calculate the standard reduction potential for the reaction:
Cl2 + 2 I- (aq)  2 Cl- (aq) + I2 (s)
Calculate the standard reduction potential for the reaction:
Ni (s) + 2 Ce4+ (aq)  Ni2+ (aq) + 2 Ce3+
Calculate the standard reduction potential for the reaction:
Cu2+ (aq) + Ca (s)  Cu (s) + Ca2+ (aq)
Which of the following is a stronger oxidizing agent: Fe (s) or Mg (s)
Which of the following is a stronger reducing agent: Cl2 (g) or Br2 (l)
Closing Time
 Read 20.2, 20.3, and 20.4
 Homework: 20.6, 20.8, 20.22, 20.23, 20.27, 20.37, 20.38,
20.39, and 20.43
 Quiz Thursday/Friday
 Saturday School this weekend!