Consider the concentration cell shown below. Calculate the...

Consider the concentration cell shown below. Calculate the cell potential at $25^{\circ} \mathrm{C}$ when the concentration of $\mathrm{Ag}^{+}$ in the compartment on the right is the following. a. $1.0 M$ b. $2.0 M$ c. $0.10 \mathrm{M}$ d. $4.0 \times 10^{-5} M$ e. Calculate the potential when both solutions are $0.10 \mathrm{M}$ in $\mathrm{Ag}^{+}$ For each case, also identify the cathode, the anode, and the direction in which electrons flow.

Consider the concentration cell shown below. Calculate the cell potential at $25^{\circ} \mathrm{C}$ when the concentration of $\mathrm{Ag}^{+}$ in the compartment on the right is the following.
a. $1.0 M$
b. $2.0 M$
c. $0.10 \mathrm{M}$
d. $4.0 \times 10^{-5} M$
e. Calculate the potential when both solutions are $0.10 \mathrm{M}$ in $\mathrm{Ag}^{+}$
For each case, also identify the cathode, the anode, and the direction in which electrons flow.

Key Concepts

Concentration Cells

A concentration cell is a type of galvanic cell where both electrodes are made from the same material but are immersed in solutions with different ion concentrations. The cell potential arises solely due to this concentration gradient, making it unique in that the materials themselves do not create a potential difference under standard conditions.

Nernst Equation

The Nernst equation provides a way to calculate the cell potential under non-standard conditions by relating the standard electrode potential to the concentrations of the reacting species. It accounts for temperature and variations in ion concentration, allowing for precise determination of the cell's electromotive force when different concentrations are involved.

Electrode Identification (Anode and Cathode)

In any electrochemical cell, the anode is where oxidation occurs, releasing electrons, while the cathode is where reduction occurs, accepting electrons. In the case of a concentration cell, the electrode immersed in the lower concentration solution acts as the anode, and the electrode in the higher concentration solution acts as the cathode, driving the spontaneous flow of electrons.

Electron Flow Direction

Electrons in an electrochemical cell flow from the anode to the cathode. This movement is driven by the potential difference created by the concentration gradient in a concentration cell, with electrons naturally moving from the site of oxidation (lower ion concentration) towards the site of reduction (higher ion concentration).

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