Oxidation / Reduction Reactions

Oxidation / Reduction Reactions

Loss of electrons or hydrogen = oxidation

Gain of electrons or hydrogen = reduction

               

                .Fe                  Fe²⁺ + 2e^–

                Na                   Na⁺ + e^–

                Reductant – a species that loses electrons.

                Oxidant – a species that gains electrons.

Free electrons do not exist in solution. Any electron lost from one species in solution must be immediately gained by another.

Ox₁ + Red₂ « Red₁ + Ox₂

Redox Reactions

The redox reactions are those involving changes of oxidation state of reactants.

Oxidation-reduction reactions involve electron (e^–) exchange.

The process is described by paired (or coupled) half-reactions involving oxidation and reduction together.

 

Example:

Fe³⁺ + Cu⁺ <-> Fe²⁺ + Cu²⁺

This reaction can be broken down into two half reactions:

Fe³⁺ + e^– <-> Fe²⁺

Cu⁺ <-> Cu²⁺ + e^–

Simple Reaction

2Na  +  Cl₂  >  2Na⁺Cl^–

The sodium starts out with an oxidation number of zero (0) and ends up having an oxidation number of +I. It has been oxidized from a sodium atom to a positive sodium ion.

The chlorine also starts out with an oxidation number of zero (0), but it ends up with an oxidation number of -I. It, therefore, has been reduced from chlorine atoms to negative chloride ions.

Half Reaction

When an oxidation or reduction reaction is written independently; for example, the reduction of CO₂

CO₂ + 4e^– + 4H⁺ = CH₂O + H₂O

or for the oxidation of H₂O

2H₂O = O₂ + 4e^– + 4H⁺

a ‘free’ electron (e^–) is written in the equation.

An overall redox reaction will never have an (e^–) shown:

CO₂ + H₂O = CH₂O + O₂

Redox Potential

 

Redox potential (also known as Reduction potential, oxidation / reduction potential, ORP, pE) is a measure of the tendency of a chemical species to acquire electrons and thereby be reduced. Reduction potential is measured in volts (V), or millivolts (mV).

Each chemical species has its own intrinsic reduction potential, the more positive the potential, the greater the species affinity for electron and tendency to be reduced.

Redox potential defined as the negative logarithm of electron activity.

The pe indicates the tendency of a solution to donate or accept a proton.

If pe is low (high electron activity), there is a strong tendency for the solution to donate protons – the solution is reducing.

If pe is high (low electron activity), there is a strong tendency for the solution to accept protons – the solution is oxidizing.

THE pe OF A HALF REACTION

Consider the half reaction

MnO₂(s) + 4H⁺ + 2e^– « Mn²⁺ + 2H₂O(l)

The equilibrium constant is

Solving for the electron activity

For equilibrium reactions:

Ox + n e^– = Red

K = [Red] / [Ox][e^–]ⁿ

                                                                                                [K is the equilibrium constant]

For the electron transfer reaction:

pE = (1/n) ( log K – log [Red] / [Ox] )

pE is analogous to pH such that pE = -log[e^–]

If reactions are always written such that n=1:

pE = log K – log [Red] / [Ox]

pE = pE^o – log [Red] / [Ox] (Nernst Equ.)

Expressions for pE

For example:     ½O₂ + 2H⁺ + 2e^– = H₂O

pE = ½ ( log K – log 1 / PO₂^1/2[H⁺]² )

pE = ½ ( log K + log PO₂^1/2[H⁺]² )

pE = ½ ( log K + log PO₂^1/2 + 2 log [H⁺] )

pE = ½ ( log 10^41.55 + ½ log PO₂ – 2 pH )

pE = 20.78 + 1/4 log PO₂ – pH

For example:     Fe²⁺ = Fe³⁺ + e^–

pE = log K + log [Fe³⁺] / [Fe²⁺]

pE = log 10^12.53 + log [Fe³⁺] / [Fe²⁺]

pE = 12.53 + log [Fe³⁺] / [Fe²⁺]

The limit of pE in water

Water may be either oxidized as

2H2O                       O2 + 4H+ +4e-

Or reduce

Electron Donors

• By far the most prevalent electron donor in the shallow subsurface is organic carbon.

CH₂O + H₂O = CO₂ + 4e^– + 4H⁺

• This half reaction is what supplies energy to microorganisms within soils.
• The electron donor (organic carbon) is the reducing agent and is oxidized to CO₂.
• Inorganic Electron Donors
• Common electron donors that participate in chemical redox couples include:
• Mn(II) = Mn(IV) + 2e^–
• Fe(II) = Fe(III) + e^–
• S^2- = SO₄^2- + 8e^–
• As(III) = As(V) + 2e-
• Environmentally-important redox systems
• These pE values refer to typical environmental conditions with pH=7 and oxygen partial pressure of 0.21 atm. The scale on the right shows the free energy of a mole of electrons relative to their level in H₂
• The two conjugate forms of any redox pair are present in equal concentrations when the pE is at the level at which the pair is shown. At pE’s above or below this level, the reduced or oxidized form will predominate.
• The sugar glucose, denoted by the general formula for carbohydrates {CH₂O}, is the source of chemical energy for most organisms. Note that it is thermodynamically stable (and thus capable of being formed by photosynthesis) only under highly reducing conditions.
• Organisms derive their metabolic free energy when electrons fall from glucose to a lower-lying acceptor on the right.
• Delivery of electrons from glucose to O₂ (8) is the source of metabolic free energy for all aerobic organisms, yielding 125 kJ per mole of electrons transferred.