RedOx Chemistry Important reactions in the Earth system … Atmosphere is an oxydant medium du to the abundance of O2 (e.g. photo-oxidant pollution, combustions …) Biochemistry depends on it ( source of energy, respiration, photosythesis )

Oxidation state A measure of the degree of oxidation of an atom in a substance. It is defined as the charge an atom might be imagined to have when electrons are counted according to an agreed- upon set of rules: (1) the oxidation state of a free element (uncombined element, eg. H2, O2 …) is zero (2) for a simple (monoatomic) ion, the oxidation state is equal to the net charge on the ion (3) hydrogen has an oxidation state of 1 and oxygen has an oxidation state of −2 when they are present in most compounds. (Exceptions to this are that hydrogen has an oxidation state of −1 in hydrides of active metals, e.g. LiH, and oxygen has an oxidation state of −1 in peroxides, e.g. H2O2) (4) the algebraic sum of oxidation states of all atoms in a neutral molecule must be zero, while in ions the algebraic sum of the oxidation states of the constituent atoms must be equal to the charge on the ion. For example, the oxidation states of sulfur in H2S, S8(elementary sulfur), SO2, SO3, and H2SO4 are, respectively: −2, 0, +4, +6 and +6. The higher the oxidation state of a given atom the greater its degree of oxidation; the lower the oxidation state the greater its degree of reduction. Reduction = reduction of the oxidation state = gain of electrons Oxydation = increase of the oxidation state = loss of electrons

In redox reactions, electrons are transferred from one species to another. A species is oxidized when it loses electrons.  Here, zinc loses two electrons to go from neutral zinc metal to the Zn 2+ ion. A species is reduced when it gains electrons.  Here, each of the H + gains an electron and they combine to form H 2. Variation of oxidation number redox reactions

A redox reaction can be decomposed in 2 redox half reactions : Guide for writing and balancing half-reaction and full redox equations Consider the unbalanced full redox reaction Identify the key element that undergoes an oxidation state change. Write down unbalanced Half reactions Balance the number of atoms of the key element on both sides. Add the appropriate number of electrons to compensate for the change of oxidation state. Add H+ (in acid medium), or OH- (in basic medium), to balance the charge on both sides of the half-reactions; and H2O, if necessary, to balance the equations in oxygen and hydrogen. Combine half reactions checking for balnce of atom and net charge. Examples … H2 + O2 -> H2O

Half - reactions are listed in redox tables. The larger is the electron affinity of the oxidizing agent, the more the equilibrium is shifted to the left and the smaller is the half-reaction equiilibrium constant. Half-reactions equilibrium constant gives a way to compare the electronegativity (oxidizing strength) of compounds. Standard electrochemical potentials (or Standard redox potentials ) E 0 (in V), are more widely used in redox tables. There is a relation between half-reactions equilibrium constants and standard electrochemical potentials ( see also Hobbs textbook chap6): where n is the number of moles of electrons transferred, and F is a constant, the Faraday. (1 F = 96,485 C/mol = 96,485 J/V/mol) Strength of oxidant and reductants

Analogy with acid-base reactions The formalism adopted for redox reactions is quite similar than for acid / base reactions The same way we defined pH for an acid base reaction, we can define pe as : pe = - log [e-] pe : can be thought of as a measure of the oxidizing potential of the solution

Sometime instead of using pe as a master variable one use instead Eh, the electrochemical potential (or redox potential) of the solution. The relation is : Redox potential can be calculated from standard redox potential for any composition using the Nernst equation : Remarks : If you check out the variation of an element oxidation number in an acid base reaction (or in a phase change equilibrium), you will find it is constant. Redox reactions and acid base reactions represent clearly different “kind” of chemistry determining the composition of a system.

pe – pH stability diagrams pe and pH can be thought as master variables of a phase space determining the oxidation and acidity state of a chemical system at thermodynamic equilibrium. As pe and pH vary in the environment, one move from one chemical regime to an other corresponding, for a given element, to a different speciation. In a pe-pH diagram, we want to draw lines materializing transition between different chemical forms. For a same phase (e.g liquid), the lines represents equal activities between the chemical forms. If there is a phase transition the exact location of the boundaries are determined by fixing arbitrarily activities Qualitatively … Quantitatively …

pe – pH stability diagrams : ex 1: Water

The water stability diagram is fundamental for natural systems because it determines the limits of aqueous chemistry

pe – pH stability diagrams : ex 2 : simplified sulfur system