Chapter 21
Electrochemistry
I. Oxidation and reduction. These represent a type of chemical reaction in which one
substance transfers an electron to another substance. These are also called Redox
reactions because they involve reduction and oxidation.
A.
Oxidation – is a chemical reaction in which a substance loses electrons
and the ionic charge on the atom increases.
B.
Reduction – is a chemical reaction in which a substance gains electrons
and the ionic charge on the atom decreases.
C.
Reducing agent (reductant) – the substance that gets oxidized, and
therefore causes the other substance to become reduced – hence, the term
“reducing agent”.
D.
Oxidizing agent (oxidant) – the substance that gets reduced, and
therefore causes the other substance to become oxidized – hence, the term
“oxidizing agent”.
E.
Oxidation always accompanies reduction and reduction always
accompanies oxidation. You cannot have one without the other.
F.
“LEO” = loss of electrons is oxidation”.
G.
“GRE” = gain of electrons is reduction”.
Example
Oxidation
Reduction
Zn(s) 
Cu2+(aq) +
Zn2+(aq)
2e
+
Cu(s)
2e-
II. Oxidation numbers revisited.
A.
One way to express the number of electrons lost or gained by a substance
in a redox reaction is assigning an oxidation number to an atom.
1.
Oxidation number can be zero, positive, or negative.
2.
neutral oxidation number or zero charge indicates that there are
the same number of protons as electrons. This is the state for
uncombined neutral atoms.
3.
positive oxidation number has fewer electrons than protons.
4.
negative oxidation number has more electrons than protons.
5.
A list of oxidation number rules is on p. 639. Know them – but, it
should be a review.
III. Electrochemistry
A.
Electrochemistry is the study between electrical and chemical energy.
1.
use of electricity to cause a chemical reaction
2.
the production of electricity by a chemical reaction.
 Both involve oxidation and reduction reactions.
B.
Voltaic cells are electrochemical cells used to convert chemical energy
into electrical energy.
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1. Half-cell is one part f a voltaic cell in which either oxidation or
reduction occur.
2. Salt bridge a tube containing strong electrolytes such as potassium
sulfate, K2SO4. It allows ions to pass from one half-cell to another but
prevents the solutions from mixing.
3. Electrode is a conductor in a circuit that carries electrons to of from a
substance.
a. The electrode where oxidation occurs is called the anode.
b. The electrode where reduction occurs is called the cathode
4. Battery is a group of cells linked together.
Car Battery
Oxidation
Reduction
C.
Pb(s) +
PbO2(s) +
SO42-(aq)
PbSO4 (s) + 2e4H+(aq) + SO42-(aq) + 2e- PbSO4 (s) + H2O(l)
How a Voltaic cell works.
1. Electrons are produced at the zinc rod.
Zn(s)  Zn2+(aq) + 2eThe zinc rod is oxidized and is called the anode.
2. Electrons leave the zinc anode and pass through the circuit to the
copper rod.
3. Electrons enter the copper rod and the following reaction takes place.
Cu2+(aq) + 2e-  Cu(s)
4. To complete the circuit both negative and positive move through the
solutions via the salt bridge.
Zn(s)  Zn2+(aq) + 2eCu2+(aq) + 2e-  Cu(s)
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 Zn2+(aq) + Cu(s)
Zn(s) + Cu2+(aq)
Net Reaction
This reaction can be represented by the following shorthand.
Zn(s) | ZnSO4(aq) || CuSO4(aq) | Cu(s)
The anode is written first, to the left of the double lines.
The cathode is written last, to the right of the double lines.
Vertical lines indicate boundaries of phases that are in contact. The double vertical
lines represent the salt bridge.
D.
Dry cell is a voltaic cell in which the electrolyte is a paste. These are
flashlight batteries.
Oxidation:
Zn(s)  Zn2+(aq) + 2e-
Reduction:
2MnO2(s) + 2NH4+(aq) + 2e-  Mn2O3(s) + 2NH3(aq) + H2O(l)
IV. Cell Potentials
A. Electrical potential of a voltaic cell is a measure of the cell’s ability to
produce electric current. Electrical potential is measured in Volts (V).
B. Cell potential is equal to the difference between the reduction potentials
between the two half-cells.
Eocell = Eored - Eooxid
C. The standard reduction potential of the hydrogen electrode has been assigned
the value of 0.00 V. See page 674 for a table of reduction potentials.
D. When calculating cell potentials if the cell potential is positive then the
reaction is spontaneous. If the cell potential is negative then the reaction is
nonspontaneous.
E. Sample calculation
Zn(s) + 2Ag+(aq)  Zn2+ + 2Ag(s)
1. First identify the half-reactions
Oxidation
Zn(s)  Zn2+ + 2e-
EoZn2+ = -0.76 V
Reduction
Ag+(aq) + e-  Ag(s)
EoAg+ = +0.80 V
Eocell = Eored
- Eooxid
Eocell = 0.80 V –(-0.76 V) = +1.56V
Eocell is greater than zero so the reaction is spontaneous.
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