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 Rules1. 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.

2. 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.

3. Hydrogen has an oxidation state of +1 when bonded to non-metals and -1 when bonded to metals (hydrides).

4. 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.

5. O, unless it is a peroxide (O22-), has an oxidation state of -2 in which case

it is a -1.6. The sum of the oxidation numbers in a neutral compound is always 0.

Redox ReactionsIn 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 half-reactions:

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= Eo

SHE + 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, Eo

cell, 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 half-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• 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.