Chapter 20 Redox Reactions – Notes
Redox reaction – a reaction in which electrons are transferred from one atom to another
Oxidation – loss of electrons from atoms of a substance
Ex – Na  Na+ + eSodium is oxidized
Reduction – gain of electrons by atoms of a substance
Ex – Cl2 + 2e-  2ClChlorine is reduced
*Memory technique* LEO GER (Loss of Electrons is Oxidation, Gain of Electrons is
Reduction)
oxidized
Ex – 2K (s) + Br2 (g)  2KBr (s)
reduced
Oxidizing agent – substance that oxidizes another substance
Reducing agent – substance that reduces another substance by losing electrons
Examples of redox chemistry:
 Silver tarnishing
 Hydrogen peroxide on a wound
 photography
 Chlorine bleach to whiten laundry
Oxidation number - number of electrons lost or gained by the atom when it forms ions;
written with a “+” or “-“ sign before the number
Determining oxidation number – rules on page 655;
Practice determining oxidation numbers:
1. KBr
2. MgCl2
3. HClO3
4. Na2Cr2O7
Steps for balancing redox reactions:
1. Write a chemical equation.
2. Determine oxidation number of each element.
3. Draw a line to connect the atoms involved in oxidation and another line connecting
the atoms involved in reduction. Write the half reaction for each on the line.
4. Balance the charges for redox reaction.
5. Balance equation.
Ex – HNO3 + HI  NO + I2 + H2O
Ch. 21 Electrochemistry
Electrochemistry – study of process by which chemical energy is converted to electrical
energy & vice versa
Recall – redox reactions involve a transfer of electrons from one atom to another
Electrochemical cell – an apparatus that uses a redox reaction to produce electrical
energy or uses electrical energy to cause a chemical reaction



Voltaic cell – type of electrochemical cell that converts chemical energy to
electrical energy by a spontaneous redox reaction
Made up of 2 parts ( half – cells):
1. Anode: where oxidation takes place (-)
2. Cathode: where reduction takes place (+)
Also need a salt bridge and a wire
Ex. Zn + Cu+2  Zn+2 + Cu
Write the half-reactions:
Zn  Zn+2 + 2e- oxidation ANODE
Cu+2 + 2e-  Cu reduction CATHODE
 Look at the activity series on page 301 or remember that the most active
elements are in the upper left corner of the periodic table. The more active
the metal, the more easily the metal loses electrons (oxidizes). Electrons
flow from the oxidized metal to the reduced metal.
Ex. Use an arrow to indicate the direction of electron flow in these pairs:
1. Fe/Cu
2. Hg/Ca
3. Zn/Na
Reduction is the gain of electrons.
Reduction potential - the tendency of an object to gain electrons.
Chart on p. 667 shows the standard reduction potential (Eo) of half-reactions.
Use this equation to determine the Eo of the cell:
Eo cell = Eo reduction – Eo oxidation
If Eo cell is +, the reaction is spontaneous. If Eo is -, the reaction is not spontaneous.
Ex. Ag+ + Al  Ag + Al+3
Eo cell = 0.7996 – (-1.662) = 2.4616, the reaction is spontaneous
Batteries – simplest form is a single voltaic cell
 Zinc–carbon dry cell – the electrolyte is a moist paste; zinc shell is the cell’s
anode [Zn (s) Zn+2 + 2 e-]; carbon (graphite) rod in the center of the dry
cell is the cathode
Electroplating – requires a battery
 Cathode – object to be plated; where reduction occurs [Ag+ + e-  Ag]
 Anode – whatever metal you want the object plated with; in this example, the
anode is silver; where oxidation occurs [Ag  Ag+ + e-]
 Silver ions in electrolyte solution are reduced to silver metal and deposited on
the object
fork
Ag
+
battery
Also review with this animation:
http://www.ausetute.com.au/redox.html
Review the redox rules with this game http://www.quia.com/jg/476236.html
Video Assignment: (CTRL-click to follow the link)
http://www.unitedstreaming.com/studentCenter/index.cfm?cdCode=TD797-3D79