Varying the concentrations of half-cells

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In summary, varying the concentrations of half-cells involves changing the concentrations of reactants or products in a half-cell reaction. It is important in order to study the effect of concentration on reaction rate and extent. This can be done by adding or removing solutions, manually or through automated equipment. Factors such as temperature, pressure, and electrolyte type can affect the concentrations. Varying concentrations can also impact the overall reaction rate and equilibrium.
  • #1
~angel~
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I'm not sure ow to do this question.

Calculate the reduction potential of a half-cell consisting of a platinum electrode immersed in a 2.0M Fe2+ and 0.2M Fe3+ solution 25c.

Thanks.
 
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  • #2
The Nernst Equation is given by:

[itex]
E= E_o - \frac{0.059}{n} log \frac{N_1}{N_2}
[/itex]

The above half cell reaction can be written as:

[itex]
Fe^2^+ ------> Fe^3^+ e^_
[/itex]


here n=1 as you can see in the cell reaction.
 
  • #3


To calculate the reduction potential of this half-cell, we can use the Nernst equation: Ecell = E°cell - (RT/nF)lnQ, where Ecell is the cell potential, E°cell is the standard cell potential, R is the gas constant, T is the temperature in Kelvin, n is the number of electrons transferred in the reaction, F is Faraday's constant, and Q is the reaction quotient.

First, we need to determine the standard cell potential (E°cell) for the half-cell. This can be found by looking up the standard reduction potentials for the Fe2+/Fe3+ redox couple. According to the NIST Chemistry WebBook, the standard reduction potential for this couple is 0.771 V at 25°C.

Next, we need to calculate the reaction quotient (Q) using the concentrations of Fe2+ and Fe3+ given in the question. Q = [Fe3+]/[Fe2+] = 0.2/2.0 = 0.1.

Now we can plug in all the values into the Nernst equation and solve for Ecell:

Ecell = 0.771 V - (8.314 J/mol·K)(298 K)/(1 mol)(96,485 C/mol)ln(0.1) = 0.771 V - 0.059 V = 0.712 V

Therefore, the reduction potential for this half-cell is 0.712 V at 25°C. This means that at standard conditions, the reaction is spontaneous and favors the reduction of Fe3+ to Fe2+. Varying the concentrations of the half-cell can affect the reaction quotient and therefore change the cell potential. This can be seen by plugging in different values for Q in the Nernst equation and observing the resulting changes in Ecell.
 

What is meant by "varying the concentrations of half-cells"?

Varying the concentrations of half-cells refers to changing the concentrations of the reactants or products in a half-cell reaction. This can be done by adding or removing a solution containing the specific ions involved in the reaction.

Why is it important to vary the concentrations of half-cells?

Varying the concentrations of half-cells allows for the study of the effect of concentration on the rate and extent of a reaction. This information is crucial in understanding the underlying mechanisms of reactions and can help in the design of more efficient processes.

How do you vary the concentrations of half-cells in an experiment?

To vary the concentrations of half-cells, one can add or remove a solution with the desired concentration. This can be done manually or through the use of automated equipment such as burettes or pumps. The concentrations can also be varied by changing the volume of the solutions.

What factors can affect the concentrations of half-cells?

The concentrations of half-cells can be affected by various factors such as temperature, pressure, and the presence of other substances in the solution. Additionally, the type and concentration of electrolytes used in the solution can also affect the final concentrations of the half-cells.

How do varying concentrations of half-cells impact the overall reaction rate?

Varying the concentrations of half-cells can have a significant impact on the overall reaction rate. Higher concentrations of reactants typically lead to faster reaction rates, while lower concentrations can slow down the reaction. Changes in concentrations can also affect the equilibrium of the reaction, leading to shifts in the overall rate of the reaction.

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