Transcript Slide 1

Redox Reactions
Oxidation-reduction reactions, also called redox
reactions, involve the transfer of electrons
from one species to another. That electron
transfer causes a change in oxidation state for
both reactive partners.
Oxidation
Numbers
(State)
Oxidation Number Rules
1.
2.
3.
4.
5.
6.
Atoms in elemental form have oxidation states of zero. For
example, each H atom in the H2 molecule has an oxidation
number of 0 and each P atom in the P4 molecule has an oxidation
number of 0.
The charge on a monoatomic ion is equivalent to its charge. For
example, K+ has an oxidation number of +1 and S2- has an oxidtion
number of -2. In ionic compounds the alkali metal ions always
have a 1+ charge and therefore has an oxidation number of +1.
The alkaline earth metals are always +2 and aluminum is always
+3 in ionic compounds.
Hydrogen has an oxidation state of +1 when bonded to nonmetals and -1 when bonded to metals (hydrides).
F, because it only forms one bond and is the most electronegative
element, has an oxidation state of -1. Additionally, the other
halogens USUALLY form a -1 charge.
O, unless it is a peroxide (O22-), has an oxidation state of -2 in
which case it is a -1.
The sum of the oxidation numbers in a neutral compound is
always 0.
Redox Reactions
In a redox reaction, the reducing agent is
oxidized, meaning that its oxidation number
increases due to the loss of one or more
electrons (it provides electrons for the other
to be reduced). The oxidizing agent is reduced,
meaning that its oxidation number has
decreased due to the gain of one or more
electrons (it takes electrons away for the other
to be oxidized).
Balancing Redox Reactions
A half-reaction is an equation that
shows either a reduction or an
oxidation. Aqueuous redox
equations are conveiently balanced
by the method of half-reactions.
Steps to Balancing Redox Reactions
1. Separate oxidation and reduction halfreactions:
2. Balance all atoms except for hydrogen and
oxygen in each half-reaction.
3. Balance oxygen by adding H2O as needed:
Steps to Balancing Redox Reactions
4. To balance hydrogen, add H+ as needed:
5. Balance the charge of each reaction by
adding electrons to side with the greater
charge:
Steps to Balancing Redox Reactions
6. Multiply each half-reaction by the least
integer factor that equalizes the number of
electrons in each half-reaction. Then, add the
half-reactions to obtain the overall balanced
reaction in acidic solution:
Galvanic or Voltaic Cells
Galvanic cells harness the electrical energy
available from the electron transfer in a redox
reaction to perform useful electrical work. In
other words, they spontaneously transform
chemical energy into electrical energy. The
transfer of electrons of a redox reaction takes
place through an external pathway.
Galvanic or Voltaic Cells
Galvanic or Voltaic Cells
"The Red Cat ate An Ox" meaning reduction takes place
at the cathode and oxidation takes place at the anode.
The anode, as the source of the negatively charged
electrons is usually marked with a minus sign (-) and
the cathode is marked with a plus sign (+).
Because the reactants are separated, the reaction can
only occur when the transfer of electrons takes place
through an external circuit. Electrons always flow
spontaneously from the anode to the cathode
(alphabetical order).
Galvanic or Voltaic Cells
Galvanic or Voltaic Cells
The voltaic cell uses a salt bridge to complete the
electrical circuit. As oxidation and reduction take
place, ions form the half-cell compartments
migrate through the salt bridge to maintain the
electrical neutrality of the respective solutions.
A potential difference exists between anode and
cathode of a voltaic cell. This cell potential
pushes electrons through the external circuit and
is measured in volts.
Galvanic or Voltaic Cells
Pt (s) | Cu2+ (aq), H+ (aq)
Mg (s) | Mg2+ (aq) || Al3+ (aq) | Al (s)
Standard Reduction Potential
For example, when the following cell is constructed,
an Eocell of 0.34 V is observed (note the setup of
the SHE as the anode because Cu2+ has a greater
reduction potential than H+):
Pt (s) | H2 (g) | H+ (aq) || Cu2+ (aq) | Cu (s)
Because the SHE has a potential of exactly zero
volts, as defined above, the reaction:
has a value of 0.34 V for its Eo (recall that Eocell=
EoSHE + Eo).
Adding Standard Reduction Potential
If Eocell is positive, then the reaction is
spontaneous. Conversely, if Eocell is negative,
then the reaction is non-spontaneous as
written but spontaneous in the reverse
direction.
To compute the cell potential of a reaction at
standard conditions, Eocell, you do not need to
balance the equation of your redox reaction.
Adding Standard Reduction Potential
Adding Standard Reduction Potential
In other words, to determine the spontaneous
reaction for any cell made up of two half cells,
reverse the half-reactions with the less
positive E°red and add it to the half-reaction
with the more positive voltage. Similarly, to
determine the cell potential of any two
coupled half-reactions, change the sign of the
potential for the reversed half-reaction and
add it to the potential of the reduction halfreaction.
Key Terms
• Electrochemistry - The study of the exchange between
electrical and chemical energy.
• Battery - A galvanic cell or cells connected in series
with a constant amount of reagents. A battery stores
energy in the form of electrical potential energy.
• Electrorefining - Process by which materials, usually
metals, are purified by means of an electrolytic cell.
The anode is the impure metal and the cathode is a
very pure sample of the metal.
• Oxidation - The loss of an electron from a species (an
increase in its oxidation number).
• Oxidation Number - A conceptual bookkeeping
numbering system that allows us to track the number
electrons transferred during a redox reaction.
Key Terms
• Electrochemistry - The study of the exchange between
electrical and chemical energy.
• Battery - A galvanic cell or cells connected in series
with a constant amount of reagents. A battery stores
energy in the form of electrical potential energy.
• Electrorefining - Process by which materials, usually
metals, are purified by means of an electrolytic cell.
The anode is the impure metal and the cathode is a
very pure sample of the metal.
• Oxidation - The loss of an electron from a species (an
increase in its oxidation number).
• Oxidation Number - A conceptual bookkeeping
numbering system that allows us to track the number
electrons transferred during a redox reaction.
Key Terms
• Reduction - The gain of an electron by a species (a decrease
in oxidation number).
• Redox - A reaction involving the transfer of one or more
electrons from the reducing agent to the oxidizing agent.
• Reducing Agent - A reactant in a redox reaction that
donates an electron to the reduced species. The reducing
agent is oxidized.
• Oxidizing Agent - A reactant in a redox reaction that
accepts an electron from the oxidized species. The oxidizing
agent is reduced.
• Galvanic Cell - An electrochemical cell with a positive cell
potential that allows chemical energy to be converted into
electrical energy.
• Cell Potential - The overall electrical potential of an
electrochemical cell. It is the sum of the reduction potential
of the cathode and the oxidation potential of the anode.