In an electrochemical cell, we physically separate the oxidation and reduction chemistry in different "compartments". The electrons from the oxidation are then run through an external circuit before being used in the reduction reaction. As this is moving negative charge from one location to another, we need to compensate for this by moving other charges to balance this displacement of charge. This is accomplished by using a salt bridge that allows the migration of spectator ions to balance the flow of electrons.
Each half of the electrochemical cell has an electrode to which the wire for our external circuit is connected. The chemistry takes place at the surface of this electrode. The electrode on the oxidation side is called the anode. The electrode on the reduction side is called the cathode. As a short hand, we simply refer to the two sides of the cell as the anode side and the cathode side.
For any given set of conditions, the chemistry is spontaneous in only one direction. When the cell is set up such that this is the direction we desire, we call the cell a voltaic cell (or a galvanic cell or a battery). When we want the chemistry to go in the non-spontaneous direction, the cell is an electrolytic cell and requires an external power source to drive the chemistry in a non-spontaneous direction. It is important to note that this is one of the amazing things about electrochemistry. We can use an external voltage to change the direction of a chemical reaction by adjusting the free energy (we essentially get to chose which is lower in free energy the products or reactants). So with an electrochemical cell, we can get a chemical reaction to move in both the spontaneous and the non-spontaneous direction. You are all familiar with this concept but you might not have thought about it as chemistry. When you are using a battery you are running a voltaic electrochemical cell in the spontaneous direction. When it gets to equilibrium the voltage on the battery is zero and no more current flows. However, you can "recharge" the battery by attaching it to an external power supply and forcing the chemistry to go in the reverse direction until you have regenerated the original chemical contents.
For more practice with electrochemical cell problems, see the worksheet and key below.
In an electrochemical cell, we are running a redox reaction, but we have physically separated the oxidation and reduction reactions into different locations. These take place in different "halves" of the cell.
We then connect the two halves of the cell with an external wire to allow for electron flow from one side to the other. We also connect the two halves with a "salt bridge" to allow spectator ions to flow to counter balance the electron flow.
The oxidation and reduction reactions are now in different locations. To allow the electrons to flow from the oxidation reaction (that is producing electrons) to the reduction reaction (that is consuming the electrons), we need to connect the two reactions electrically. This is accomplished by using metal electrodes. These electrodes serve a the place where the chemistry is taking place. Sometimes they are part of the reaction. Sometimes they are simply used as a means to "deliver/collect" the electrons. The electrode in the oxidation reaction is call the anode. The electrode for the reduction reaction is called the cathode. The electrons flow from the anode to the cathode.
Let's look at a picture of the cell below which contains zinc metal and zinc ions on the oxidation side and copper metal and copper ions on the reduction side.
The left hand side is a picture of the anode half of the cell where the oxidation is taking place. The solid piece of zinc is the anode. The righthand side is the reduction half where the copper 2+ ions are being reduced to copper metal. The solid piece of copper is the cathode. The two sides are connected by a wire that allows electrons to flow form the anode to the cathode. They are also connected by a salt bridge. Since negative charge is flowing from the anode side to the cathode side, some spectator ions need to compensate this flow (or the cathode side will end up with lots of extra negative charge). So in the salt bridge either sodium + ions could flow from the anode side to the cathode side or sulfate 2- ions could flow from the cathode side to the anode side. Or both could happen.
For some reactions, neither the reactants nor the products are a metal. In this case, an inert (non-reactive) metal electrode is required to deliver/collect the electrons for the reaction.
For example, in the cell below at the anode, Fe2+ is oxidized to Fe3+. This reaction occurs at the surface of a Pt electrode.
It becomes very cumbersome to draw a picture of an electrochemical cell each time you would like to discuss one.
As such, we have developed a short hand notation for a cell.
For the cell above we have the anode on the left and the cathode on the right. The oxidation is occurring at the anode
\[\rm{Ni(s) \rightarrow Ni^{2+}(aq) + 2e^-}\]
The reduction is occurring at the cathode
\[\rm{2H^+(aq) + 2e^- \rightarrow H_2(g)}\]
The overall reaction is
\[\rm{Ni(s) + 2H^+(aq) \rightarrow Ni^{2+}(aq) + H_2(g)}\]
Rather than having this picture of the cell, we have a shorthand that is the same moving from left to right in our diagram. We start with the anode which we write to the right of a "" symbol. In total, this will look like
< Ni | Ni2+ || H+ | H2 | Pt >
This is the most generic of notations. To be more specific we might include the concentrations (or pressures) along with the phases.
< Ni(s) | Ni2+ (aq,1M) || H+ (aq, 1M) | H2 (g, 1 atm) | Pt(s) >
Reading this from left to right, we have a nickel anode that is oxidizing to Ni2+ connected by a salt bridge to a solution of H+ that is reducing to H2 gas at a Pt cathode. Note: on both the anode and cathode sides of the notation the reactants are on the left and the products are on the right. (Ni -> Ni2+) (H+ -> H2). Also, in the cell notation the half reaction might not be balanced with each other or themselves. For example you write H+ not 2H+ because if you have a concentration, it is for H+.
For any half reaction, we can measure the potential (free energy) by comparing it to a standard reaction. The reaction we choose is the following.
\[\rm{2H^+(aq) + 2e^- \rightarrow H_2(g)}\]
Where the concentration of the H+ is 1M and the pressure of the H2 gas is 1 atm. Such an electrode is called a "standard hydrogen electrode" or SHE.
We can then make a cell with any other half reaction (under standard conditions where concentrations are 1M) and measure the potential (voltage).
For example here is a cell to compare the potential of the reaction of Ni to Ni2+ to the SHE.
We can then measure all such combinations and tabulate the data. Note: for some reactions, we are measuring the oxidation and for others the reduction, but we will always tabulate the reduction potential for the half reaction.
Having compared many reactions to the standard hydrogen potential, we can now make a table of reduction potentials for all half-reaction, (or oxidation potentials but we need to pick one and stick to it).
Below is a table for reduction potentials.
At the top of the table are the reactions that are "easiest" to oxidize and thus the hardest to run as reductions. At the bottom are the easiest reductions, and thus the most difficult to run as oxidations.
When looking at the table, we need to be careful since everything is written as a reduction. For example, from this table we can find the substance that is easiest to oxidize. That is from the top of the table. But the substance that is being oxidized appears as a product in this table. Li(s) is the easiest to oxidize. The easiest to reduce is a reactant in the table. F2 gas is the easiest to reduce.
Since Li is easy to oxidize, it is an excellent reducing agent (it reduces something else when it is oxidized). F2 is a great oxidizing agent (it oxidizes something else when it is reduced).
From this table, we can now figure out what reactions will be spontaneous. For example, if something is higher in the table (lower potential) it will be oxidized while the reactions with higher (more positive potentials) will be reduced.
Will Ag+ oxidize Fe? Yes. How do we know? The reduction potential for Ag+ is more positive than that for Fe2+. So Ag+ is a strong enough oxidizing agent to oxidize Fe to Fe2+. On the other hand it could not oxidize H2O to H+ and O2.
Below is as more extensive table of standard reduction potentials. An even larger data set can be found on the wikipedia link below. A pdf document of the Standard Reduction Potential Table used in class can be found on the attachment below.
We can calculate the standard potential for any electrochemical cell from the standard potentials of the two half reactions.
For example, imagine we had the following cell.
< Ni | Ni2+ || Fe3+ , Fe2+ | Pt >
We have at the anode side the reaction
\[\rm{Ni \rightarrow Ni^{2+} + 2e^-}\]
and at the cathode side the reaction
\[\rm{Fe^{3+} + 1 e^- \rightarrow Fe^{2+}}\]
The standard potential for this cell is simply
\[\rm{\mathcal{E}^{\circ} = \mathcal{E}^{\circ}_{cathode} - \mathcal{E}^{\circ}_{anode}}\]
Where \(\mathcal{E}^{\circ}\) for each electrode is the standard reduction potential for the half-reaction. Using our table of standard reduction potentials we find for
\[\rm{Ni^{2+} + 2e^- \rightarrow Ni(s) \mathcal{E}^{\circ} = -0.25 V}\]
\[\rm{Fe^{3+} + e^- \rightarrow Fe^{2+} \mathcal{E}^{\circ} = +0.77 V}\]
So the standard potential is
\[\rm{\mathcal{E}^{\circ} = \mathcal{E}^{\circ}_{cathode} - \mathcal{E}^{\circ}_{anode} = 0.77 - (-0.25) = +1.02 V }\]
VERY IMPORTANT. Note: the potential is simply the energy difference between the two half reactions. Do not try to multiple the potentials by the number of electrons. The number of electrons simply relates how many electrons there are per reaction. How many Fe3+ will be reduced per Ni atom that is oxidized. The potential difference (free energy difference) between the two half-reactions is not dependent on the number of electrons.
An electrochemical cell in which the chemistry is spontaneous is called a voltaic cell. This means that the oxidation will occur spontaneously at the anode and the reduction spontaneously at the cathode. We should note that the notation of something as a voltaic cell is a choice. Just as we choose what we want to call the "reactants" and "products" in a chemical reaction, our choice of anode and cathode is based upon what chemistry we would like to see occur. For a voltaic cell our choice ends up being the spontaneous choice.
The standard potential for a voltaic cell is positive. The standard free energy for a reaction ΔrG0 is related to the standard potential, E0, such that negative free energy (spontaneous) corresponds to positive potential. The beauty of electrochemistry and electrochemical cells is that we can now directly measure the free energy difference by measuring electrical potential.
Voltaic cells can be referred to as Galvanic Cells. This is a different word for the identical concept. They are also batteries. Since a battery is an electrochemical cell that produces a voltage (and current) spontaneously, it is a voltaic cell. You will find these three terms used interchangeably.
Below is a picture that summarize the ideas for a voltaic cell.
For a voltaic cell E > 0, ΔG < 0.
Electrons flow from anode to cathode.
The one odd detail is the "sign" of the electrodes. For a voltaic cell the cathode is assigned the "+" sign.
An electrochemical cell in which the chemistry is non-spontaneous is called a electrolytic cell. This means that the oxidation will not occur spontaneously at the anode and the reduction will not be spontaneous at the cathode. This means we would like chemistry to occur that is uphill in free energy. This means that it can't happen unless we make it happen. For most chemical reactions, this can only be accomplished by coupling it to another reaction in which the free energy change is sufficiently negative that the two reactions together are spontaneous. In the case of electrochemistry, it is significantly easier as we can simply drive the reaction by attaching an external power supply.
The standard potential for a voltaic cell is negative. The standard free energy for a reaction ΔrG0 is related to the standard potential, E0, such that positive free energy (non-spontaneous) corresponds to negative potential. To make the reaction "go," we must apply a voltage that is significant enough to overcome the negative potential of the cell itself.
Electrolytic cells thankfully don't have many names. However processes such as electroplating, or electrodeposition are typically electrolytic cells.
Below is a picture that summarizes the cell nomenclature for a electrolytic cell.
For an electrolytic cell E < 0, ΔG > 0.
Electrons flow from anode to cathode (this is always the case). For an electrolytic cell however, this flow is not spontaneous but must be driven by an external power source.
In an electrolytic cell, the anode has the "+" sign.
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