Oxidation-reduction reactions predicting spontaneous

Predicting the Spontaneous Direction of an Oxidation-Reduction Reaction... 

Because reaction (19.3) is a spontaneous reaction, the displacement of Zn " (aq) by Cu(s)—the reverse of reaction (19.3)—does not occur spontaneously. This is the observation made in Figure 19-1. In Section 19-3, we will discuss how to predict the direction of spontaneous change for oxidation-reduction reactions. 

We can use the electrochemical series to predict the thermodynamic tendency for a reaction to take place under standard conditions. A cell reaction that is spontaneous under standard conditions (that is, has K > 1) has AG° < 0 and therefore the corresponding cell has E° > 0. The standard emf is positive when ER° > Et that is, when the standard potential for the reduction half-reaction is more positive than that for the oxidation half-reaction. 

When predicting whether a / V / reaction is spontaneous using Table 18.1, some students find it useful to circle the two potential reactants and connect them with a line. If the line has a negative slope, the reaction is spontaneous if the line has a positive slope, the reaction is nonspontaneous. If both potential reactants are on the same side of the arrows, no reaction can occur because an oxidation requires a simultaneous reduction. 

To predict whether a redox reaction is spontaneous, remember that an oxidizing agent can oxidize any reducing agent that lies below it in the table but can t oxidize one that lies above it. To calculate E° for a redox reaction, sum the E° values for the reduction and oxidation half-reactions. 

Redox reactions are more conveniently described in terms of relative electrical potentials instead of the equivalent changes in Gibbs free energy. The electrons in Equation 6.8 come from or go to some other redox couple, and whether or not the reaction proceeds in the forward direction depends on the relative electrical potentials of these two couples. Therefore, a specific electrical potential is assigned to a couple accepting or donating electrons, a value known as its oxidation-reduction or redox potential. This redox potential is compared with that of another couple to predict the direction for spontaneous electron flow when the two couples interact—electrons spontaneously move toward higher redox potentials. The redox potential of species /, ), is defined as... 

We saw earlier that one can predict the thermodynamic open circuit (zero current) potential of an electrochemical cell by combining two half-cell reactions, one for the anode and the second for the cathode. The half-cell reaction with the lower (i.e., more negative) equilibrium potential will proceed spontaneously in the anodic direction, where the electrode acts as an electron sink for the anodic de-electronation (oxidation) reaction and the higher equilibrium potential reaction will occur spontaneously at the cathode, where the electrode acts as an electron source for the electronation (reduction) reaction. A cell with spontaneous reactions at the anode and cathode is called a self-... 

Oxidation and Reduction Some Definitions 578 16.5 The Activity Series Predicting Spontaneous Redox Reactions 589 16.7 Electrolysis Using Electricity to Do Chemistry 597... 

The potential measured with an electrode in contact with a solution of its ions. Electrode potential values will predict whether a substance will be reduced or oxidized. Values are usually expressed as a reduction potential (M + —> M). A positive electrode potential would indicate that reduction is spontaneous. A negative potential for this reaction would suggest that the oxidation reaction (M —> M +) would be spontaneous. Electronegativity The tendency of an atom to pull an electron toward it in a chemical bond ... 

The redox couples are often located on a scale of increasing potentials according to their standard potential values (as acids are located on the acidity scale according to their pKa values). The origin of the scale is fixed at the value E = 0.00 V, which is the standard potential of the couple H+/H2(g) (Fig. 14.1). The strongest oxidants are located farthest on the right. The strongest reductants are located farthest on the left. For example, if the Zn +/Zn(s) and Cu +/Cu(s) couples are placed face to face, we can, with the rule above, predict the spontaneous reaction... 

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