Consider the concentration cell in Fig. 18.10 . If the Fe^2+ concentration in the right compartm

step 1 The concentrations of iron 2 plus in the two compartments are now .01 and 1 times 10 to the nega

Consider the concentration cell in Fig. 18.10 . If the Fe^2+...

Consider the concentration cell in Fig. $18.10 .$ If the $\mathrm{Fe}^{2+}$ concentration in the right compartment is changed from 0.1 $\mathrm{M}$ to $1 \times 10^{-7} \mathrm{M} \mathrm{Fe}^{2+},$ predict the direction of electron flow, and designate the anode and cathode compartments.

Key Concepts

Concentration Cell

A concentration cell is a type of electrochemical cell in which both electrodes are made of the same material, but the ionic concentrations in the two half?cells are different. The difference in concentration creates a potential difference between the electrodes, driving a spontaneous redox reaction as the cell works to equalize the ion concentrations.

Nernst Equation

The Nernst equation relates the electrode potential of a half?cell to the standard electrode potential and the activities (or concentrations) of the chemical species involved in the redox reaction. It is essential for understanding how changes in ion concentration affect the cell potential and thus predict the direction of electron flow in a concentration cell.

Redox Reactions

Redox reactions, which involve oxidation and reduction processes, underpin the functioning of electrochemical cells. In these reactions, oxidation occurs at the anode (loss of electrons) and reduction occurs at the cathode (gain of electrons). Recognizing which side undergoes oxidation or reduction allows us to designate the anode and cathode and determine the direction of electron flow.

Galvanic Cell

A galvanic cell is an electrochemical cell that drives a spontaneous redox reaction to produce electrical energy. In the context of a concentration cell, even though both electrodes are identical, the difference in ion concentration creates a potential that forces electrons to move from the anode (lower concentration electrode where oxidation occurs) to the cathode (higher concentration electrode where reduction occurs).

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