Using electrode potentials, calculate the standard...

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

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Key Concepts

Standard Electrode Potentials

These are values that indicate the intrinsic ability of a substance to gain electrons (be reduced) under standard conditions. They provide a measure of the tendency for redox reactions to occur when electrodes are set up in an electrochemical cell, often referenced relative to the standard hydrogen electrode.

Standard Free Energy Change

This concept refers to the change in Gibbs free energy associated with a reaction occurring under standard conditions. It determines the spontaneity of a reaction and is directly related to the cell potential by the equation ?G° = -nFE°, where n is the electrons transferred and F is the Faraday constant.

Faraday Constant

The Faraday constant represents the total charge per mole of electrons (approximately 96,485 coulombs/mol). It serves as a bridge between the electrochemical potential (voltage) and the energy change in electrical work or chemical processes.

Electron Transfer in Redox Reactions

In redox chemistry, reactions are divided into oxidation and reduction half-reactions that involve the transfer of electrons. Determining the number of electrons exchanged is crucial since it directly influences the calculation of free energy change associated with a redox process.

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