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What is oxidant and reductant? What is the half-reactions and redox pairs?

A quite general reaction:


– n e–

Ared + Box Aox + Bred  
In this reaction + n e–  

Ared is oxidize because it loses electrons; it is a reductant (reducing agent) because it acts as donor of electrons and causes another species to be reduced.

Box is reduced because it gains electrons; it is an oxidant (oxidizing agent) because it acts as acceptor of electrons and causes another species to be oxidized.

In the reaction


– 2 e–

Fe(s) + Cu2+(aq) Fe2+(aq) + Cu(s)

+ 2 e–


Fe is oxidized, it acts as reductant of Cu2+,


Cu2+ ion is reduced, it acts as oxidant of Fe.



Every redox reaction can be formally separated into two parts called half-reactions (half-equations, half-cells) that represent either oxidation only or reduction only; they do not occur without the other half-reaction taking place at the same time:


oxidation Fe Fe2+ + 2 e–
reduction Cu2+ + 2 e– Cu


Pairs of the oxidized and reduced species Aox / Ared and

Box / Bred that appear in half-equations are called

redox pairs (or redox couples).


Components of a particular redox pair can differ not only in the number of electrons but also in the number of hydrogen, oxygen, as well as other atoms.

Hydrogenation and dehydrogenation, Oxygenation and deoxygenation. Hydrogenation and hydratation, dehydrogenation and dehydratation are same phenomenon? Why?

Hydrogenation and dehydrogenation are redox reactions, the products of which contain more or less hydrogen atoms


(as well as less or more multiple covalent bonds – the terms saturation or desaturation are also used).


Oxygenation and deoxygenation are redox reactions, the products of which contain more or less oxygen atoms.


A special type of oxygenation is hydroxylation.

Different types of redox reactions – examples:


– Loss and gain of electrons


Zn + Cu2+ Zn2+ + Cu oxidation of zinc
Cu2+ + Fe Cu + Fe2+ reduction of cupric ion to copper



– Oxygenation and deoxygenation


C(s) + O2 →CO2 oxidation (combustion) of carbon


CO2→ CO + ½O2 reduction of carbon dioxide by deoxygenation

– Dehydrogenation and hydrogenation



CH3CH2-OH – 2H CH3CH=O dehydrogenation of ethanol to acetaldehyde  
  CH3–C –COOH + 2H    
CH3–CH–COOH hydrogenation (reduction)  
        OH of pyruvate to lactate  



Do not confuse the terms

hydrogenation dehydrogenation






hydratation, dehydratation !

In organic chemistry,


hydrogenated products are sometimes named by adding the prefix dihydro– to the name of a original compound, and dehydrogenated products by adding the prefix dehydro– to the name of a original compound.



Hydratation and dehydratation are not redox reactions; there is no change in the sum of the both carbon oxidation numbers (one of them is oxidized and another is reduced in addition or elimination of water).


Which kind of redox reactions type do you know? What is Standard electrode potentials E0, electromotive force (electric driving force)? Nernst equation?

Well-known strong oxidants and reductants – examples:


Oxidizing agents – H2O2, KMnO4, K2Cr2O7, Cl2, I2 Reducing agents – H2, C, Fe, Zn, SnCl2


Oxidants and reductants differ in their ability to react with other agents considerably.


The strength of oxidants and reductants (their tendency to gain or lose electrons) is expressed for particular redox pairs by standard electrode potentials E0.

Standard electrode potential E0

is the potential for an electrochemical half-cell

(both oxidized and reduced form at c = 1 mol/l) established relative to the potential of 0.000 V for the standard hydrogen electrode

(H+/H couple under standard state conditions).



The equilibrium electromotive force Ecell (the potential of the galvanic cell) that is the potential difference between the two

half-cells is measured:

E0cell = E0 = E0red – E0ox

Any other reference electrode (which is stable and the potential known) may be used for measurement of electrode potentials, e.g. silver chloride

or saturated calomel electrode (E0 = + 0.246 V).

Electrode potentials E under non-standard conditions


Nernst equation


E = E0 +   RT   ln [Aox]a  



E, E0 el. potentials in volts


R = 8.314 l kPa K–1 mol–1

F = 96 500 C mol–1


n = number of moles of electrons transferred


[Aox] and [Ared ] relevant concentrations of reactants

After expressing R, T (298 K), and F in numbers and transposing natural logarithm into decadic (ln x = 2.3 log x), the equation will take the form

  0.059   [A ]a    
E = E0 +   log ox    

The electrode potential of half-cells at various concentrations of redox pair components can be calculated. On the contrary, the ratio of both redox pair components can be estimated from the measured values of electrode potentials.


Date: 2015-01-29; view: 3117

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