How do we find the potential for the entire cell? To find the overall reaction, we add together our reduction half-reaction and our oxidation half-reaction. That gave us our overall reaction. To find our standard cell potential, we just need to add together our reduction potential for the half-reaction and the oxidation potential for the oxidation half-reaction.
To find the potential for the cell, we add the reduction potential and the oxidation potential. That gives us our standard cell potential. I talked about the fact that you can use a voltmeter to measure the potential difference, to measure the voltage of a voltaic cell. That's one of the nice things about the standard reduction potential table. We can calculate the voltage of our voltaic cells this way.
Let's look in more detail at our half-reactions. Let's start with the oxidation half-reaction. We know that zinc is being oxidized, right? Zinc is losing 2 electrons, and those two electrons that zinc loses are the same two electrons that caused the reduction of copper.
Zinc is the agent for the reduction of copper. We say that zinc is the reducing agent. Sometimes students find this confusing because zinc is being oxidized, so why is it the reducing agent?
So zinc is the reducing agent. Let's look at our standard reduction potential table and let's see if that can help us understand oxidizing agents and reducing agents. We've been comparing these two half-reactions, right?
These two half-reactions right here. It has the higher, has the more positive value, I should say, for the standard reduction potential. As you go up on your standard reduction potential, you're increasing in the tendency for something to be reduced, and therefore, you're increasing the strength as an oxidizing agent.
As you move up on your standard reduction potential, increased strength as an oxidizing agent. All right, let's think about the opposite.
Since the standard electrode potentials are given in their ability to be reduced, the bigger the standard reduction potentials, the easier they are to be reduced; in other words, they are simply better oxidizing agents. For example, F 2 has a potential of 2.
F 2 reduces easily and is therefore a good oxidizing agent. In contrast, Li s would rather undergo oxidation, so it is a good reducing agent.
In a galvanic cell, where a spontaneous redox reaction drives the cell to produce an electric potential, the change in Gibbs free energy must be negative. This is the opposite of the cell potential, which is positive when electrons flow spontaneously through the electrochemical cell.
As such, the following rules apply:. The other approach, which we will use, requires information like that given in Appendix L. Using those data, the cell potential can be determined.
If the cell potential is positive, the process is spontaneous. Collecting information from Appendix L and the problem,. The process is not spontaneous under standard conditions.
Check Your Learning What is the cell potential for the following reaction at room temperature? What are the values of n and Q for the overall reaction? Is the reaction spontaneous under these conditions? Finally, we will take a brief look at a special type of cell called a concentration cell. In a concentration cell, the electrodes are the same material and the half-cells differ only in concentration. Since one or both compartments is not standard, the cell potentials will be unequal; therefore, there will be a potential difference, which can be determined with the aid of the Nernst equation.
Concentration Cells What is the cell potential of the concentration cell described by. Substituting into the Nernst equation,. Check your answer: In a concentration cell, the standard cell potential will always be zero. Check Your Learning What value of Q for the previous concentration cell would result in a voltage of 0.
If the concentration of zinc ion at the cathode was 0. Electrical work w ele is the negative of the product of the total charge Q and the cell potential E cell. Electrical work is the maximum work that the system can produce and so is equal to the change in free energy.
Thus, anything that can be done with or to a free energy change can also be done to or with a cell potential. The Nernst equation relates the cell potential at nonstandard conditions to the logarithm of the reaction quotient. Concentration cells exploit this relationship and produce a positive cell potential using half-cells that differ only in the concentration of their solutes.
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