Using electrode potentials, calculate the standard...

Using electrode potentials, calculate the standard freeenergy change for the reaction
$$
\mathrm{Na}(s)+\frac{1}{2} \mathrm{Br}_{2}(g) \stackrel{\mathrm{H}_{2} \mathrm{O}}{\longrightarrow} \mathrm{Na}^{+}(a q)+\mathrm{Br}^{-}(a q)
$$

Key Concepts

Standard Electrode Potential

This is the measure of the intrinsic tendency of a chemical species to be reduced or oxidized under standard conditions (1 M concentration, 1 atm pressure, 25°C), expressed in volts relative to the standard hydrogen electrode. These potentials are fundamental in determining the direction of electron flow in an electrochemical cell and, consequently, the feasibility of redox reactions.

Gibbs Free Energy Change

The standard free energy change (?G°) indicates the spontaneity of a reaction under standard conditions. In electrochemistry, it is directly related to the cell potential by the equation ?G° = -nFE°cell, where a more positive cell potential results in a more negative free energy change, signifying a spontaneous reaction.

Electron Transfer Number

This concept denotes the number of electrons exchanged in the redox reaction, which is critical for linking the stoichiometry of the reaction to its thermodynamic properties. The integer value of electrons transferred (n) is used in calculations involving the conversion of electrical potential to free energy.

Faraday Constant

The Faraday constant (F) represents the total charge of one mole of electrons, approximately 96485 coulombs per mole. It serves as a crucial conversion factor in the energy equation ?G° = -nFE°cell, bridging the gap between electrical measurements (volts) and chemical energy (joules).

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