Basic Redox Concepts Francisco Javier Cervigon Ruckauer

Basic Redox Concepts

REDOX CHEMISTRY INTRODUCTION

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BALANCING REDOX EQUATIONS BY OXIDATION NUMBER METHOD

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BALANCING REDOX EQUATIONS BY ION-ELECTRON METHOD

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QUANTITATIVE AND THERMODYNAMIC ASPECTS OF REDOX REACTIONS

Splitting an overall reaction into two half-reactions provides an easy way to understand redox reactions but it is important to note that this division is purely conceptual. Redox reactions occur by a more complicated mechanism in which electrons are never free. According Moeller[1] it is more correct to consider oxidation-reduction as an increase or decrease in the oxidation number or state. In order to predict a reaction it would be really useful to determine the tendency toward an increase or decrease in oxidation state.

Just as some acids or bases are stronger than others, the relative oxidizing or reducing power of a half reaction also have a range of values. These values can be measured quantitatively and in order to help organize the oxidation and reduction trends of species they are expressed and tabulated by convention as the half-cell potential of reduction. The potential is relative to the potential of a half reaction in which a hydrogen ion (I M) is reduced to hydrogen gas, to which is assigned a standard potential, Eo of zero.

2H+(aq)+2e−→H2(g)Eo=0,00V

A complete list of potential can be found here. As shown there are positive and negative values, where the more positive voltage in the standard reduction potential the stronger the oxidizing power of the species. Conversely if a reduction potential has a relative large negative reduction potential the reaction proceeds readily in reverse, and the more negative the more reducing power of the species.

The overall standard potential for a redox reaction may be calculated from the standard reduction potentials for the constituent half-reactions. The half-reaction with the more positive reduction potential will be the oxidizing agent (the cathode in a galvanic cell) while the species with a more negative reduction potential will be the reducing agent (the anode in a galvanic cell). The cell potential can be easily calculated by the difference between both reduction potentials, the oxidant minus the reductant.

Eoreaction=Eocathode−Eoanode

When the overall potential is positive the reaction will be spontaneous in the written direction while if the value is negative the reactive will be spontaneous in the opposite direction. The potential cell difference is the driving force of the reaction for a specified oxidation-reduction couple. The value of this potential is related to the change in Gibbs energy for the overall reaction.

ΔGo=−nFEo

Where F is the Faraday constant (96485 C/mol), and n the number of moles of transferred electrons. So although a positive value of Eoindicates spontaneous reaction is more appropriate to consider that spontaneity according to the negative value of ΔGo.

As we know from thermodynamics, variation of free energy is related with the equilibrium constant by the equation:

ΔGo=−RTlnKeq

thus, we can establish the following relationship between the electrode potential and the equilibrium constant:

lnK=nFE0RT

However, we have to keep in mind that half-cell potential is concentration dependent. Under non-standard conditions the electrode potential can be calculated from the Nernst Equation

E=Eo−RTnFlnQ

Where Q is the quotient of the reaction,R is the ideal gas constant (8.314V⋅C/mol⋅K), T is the temperature in kelvin, Fthe faraday constant, n transferred electrons and Eois the potential under standard conditions (I M for solution or 1 bar for gases). We can express the Nernst equation employing decimal logarithm as follows:

E=Eo−2,303∗RTnFlogQ

The factor 2,303RT/nF has a value of 0,0591 V at 25oC so we finally write that as:

E=Eo−0,0591nlogQ

One final consideration that is important to note is that the value E is an intensive property so it is independent of the amount of matter. Thus, if we combine two half-cell reactions in order to get another half-reaction is not possible to deduce the potential of this new reaction by adding or subtracting the potentials of the combined half-reactions. However, we can use the value of variations of free energy to determine that value.

For example, if we don’t know the potential of the following half-reaction of reduction of iron(3+) to iron metal

Fe3+(aq)+3e−→Fe(s)Eo=?V

And we have the potentials for the reduction of iron(3+) to iron(2+) and from iron(II) to iron metal according the following equations:

Fe3+(aq)+e−→Fe2+(aq)Eo=+0,77V

Fe2+(aq)+2e−→Fe(aq)Eo=−0,44V

It is obvious that by adding these two equations we obtain the desired unknown reduction of iron(3+) to iron metal, although the potential won’t be the sum of the potentials of both combined reactions. However, if we obtain the variation of free energy for both equations, due to free energy being an extensive property we can obtain the free energy of the combined reaction by simply adding the values of ΔGo=−nFEo for both reactions.

Fe3+(aq)+e−→Fe2+(aq)ΔGo=−1(F)(+0.77V)

Fe2+(aq)+2e−→Fe(aq)ΔGo=−2(F)(−0,44V)

Fe3+(aq)+3e−→Fe(aq)ΔGo=(+0,11FV)

We can now calculate Eoemploying its relationship with the variation of free energy:

E0=−ΔG0nF=−0,11FV3F=−0,04V

[1] “Inorganic Chemistry”, New York, John Wiley & Sons. Inc. 1952

EXAMPLE OF REDOX REACTIONS

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Francisco Javier Cervigon Ruckauer