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CHMB31H3 (8)
Chapter 5

CHMB31 Chapter 5

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Alen Hadzovic

Chapter 5: Oxidation and Reduction Reduction Potentials • Reduction: electron gain, decrease in the oxidation number of the element. • Oxidation: electron loss, increase in the oxidation number of an element. • Reducing Agent: the species that supplies electrons. • Oxidizing Agent: the species that removes electrons. • A redox reaction can be expressed as the difference of two reduction half- reactions. o Half-Reactions: the electron loss (oxidation) and gain (reduction) are displayed. o Oxidation Half-Reaction: a substance loses electrons (Zn(s)  Zn^2+ + 2e-). o Reduction Half-Reaction: a substance gains electrons (2H+ + 2e-  H2(g)). o Redox Couple: the oxidized and reduced species in a half-reduction. • A reaction is thermodynamically favorable or spontaneous in the sense K > 1, if E°> 0, where E°is the difference of the standard potentials corresponding to the half-reactions into which the overall reaction may be divided. o Δ G < 0 if the reaction is spontaneous. o Δ G°= -RTlnK o Galvanic Cell: an electrochemical cell in which a chemical reaction is used to generate an electric current, in which the reaction driving the electric current through the external circuit is the reaction of interest. o Cathode: the electrode at which reduction occurs. o Anode: the side of oxidation. o Δ G = -vFE, where F is the Faraday’s constant at 96.48 kC/mol. o Δ G°= -vFE°, where E°is the standard potential. o Fuel Cell: converts a chemical fuel directly into electrical power. • The atomization and ionization of a metal and the hydration enthalpy of its ions all contribute to the value of the standard potential. • The oxidized member of a couple is a strong oxidizing agent if E° is positive and large; the reduced member is a strong reducing agent if E° is negative and large. o A negative standard potential signifies a couple in which the reduced species is a reducing agent for H+ ions under standard conditions in aqueous solution. o The reduced member of a couple has a thermodynamic tendency to reduce the oxidized member of any couple that lies above it in the series. • The cell potential at an arbitrary composition of the reaction mixture is given by the Nernst Equation. o Δ Gr = Δ Gr°+ RTlnQ, where Q = ([C]^c [D]^d)/([A]^a [B]^b). o Nernst Equation: Ecell = Ecell° - (RT/vF) lnQ, where the reaction is spontaneous if Ecell > 0 and Δ Gr < 0; at equilibrium, Ecell = 0 and Q = K. o lnK = vFEcell°/RT o Δ Sr°= vF{Ecell°(T2) – Ecell°(T1)}/T2-T1 Redox Stability • Many redox reactions in aqueous solution involve transfer of H+ as well as electrons and the electrode potential therefore depends on the pH. o The potential decreases (becoming more negative) as the pH increases and the solution becomes more basic. • For metals with large, negative standard potentials, reaction with aqueous acids leads to the production of H2 unless a passivating oxide layer is formed. o M(s) + H2O(l)  M+(aq) + ½ H2(g) + OH-(aq); favorable when M is an s-block metal, a 3d-series metal from Group 3 to at least Group 8 or 9 and beyond or a lanthanoid. o M(s) + H+ (aq)  M+(aq) + ½ H2(g) o Passivated: protected against reaction (magnesium and aluminum can be sued for years in the presence of water and oxygen). • Water can act as a reducing agent, that is be oxidized by other species. • The stability field of water shows the region of pH and reduction potential where couples are neither oxidized by nor reduce hydrogen ions. o A reducing agent that can reduce water to H2 rapidly, or an oxidizing agent that can oxidize water to O2 rapidly, cannot survive in aqueous solution. o Stability Field: the range of values of potential and pH for which water is thermodynamically stable towards both oxidation and reduction. o Species that are oxidized by water have potentials lying below a H2 production line.
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