Galvanic and Electrolytic Cells

Anode and Cathode in Galvanic and Electrolytic Cell

Before we get into how to electroplate gold onto metal, first we have to tackle the anode and cathode in galvanic and electrolytic cells. But first we should discuss what an electrochemical cell is.

Now, wait a second, what does spontaneity have to do with this? Well, remember that a spontaneous reaction will occur on its own and is thermodynamically favoured. Conversely, a nonspontaneous reaction will never occur on its own and is thermodynamically forbidden.

Okay, so we know that a redox reaction is occurring in some device. Remember that a redox reaction can be split into two half-reactions, the reduction reaction, and the oxidation reaction. An electrochemical cell mirrors a redox reaction because there are two half-cells, one holding an oxidation reaction, and the other holding a reduction reaction.

In a Galvanic cell, a spontaneous reaction converts chemical energy into electrical energy. Recall that a spontaneous reaction occurs without any external help. The anode holds an oxidation reaction, while the cathode holds a reduction reaction. Electrons will always flow from the higher potential cell to the lower potential cell (i.e., a high amount of electrons will flow to a place with few electrons). Therefore, when setting up a Galvanic cell, you need only understand the standard reduction potentials of half-reactions.

It should be understood that electrons, and energy in general, will move from a place of high density to a place of low density. To do so otherwise would violate the second law of thermodynamics. Now, although this could never occur spontaneously, that doesn't mean it's impossible. By applying a voltage, or essentially by using electricity, this can be forced to occur.

With enough energy, any reaction in the universe can be forced to occur. This is what happens in an electrolytic cell. Take the spontaneous reaction of a Galvanic (Voltaic) cell, and reverse it. This can only happen by applying an amount of energy unique to the redox reaction.

To sum up:

• Electrons will move towards areas with fewer electrons
• Nonspontaneous reactions are operated by using electricity

Galvanic and Electrolytic Cell Diagrams

To dive further into Galvanic and electrolytic cells, we should take a look at a cell diagram. In a spontaneous reaction, electrons flow from the negatively charged anode to the positively charged cathode.

Fig. 1 - Galvanic vs. Electrolytic Cell - Electrons always move from the anode to the cathode. In each cell, the anodic and cathodic species will change.

On the left, it is clear that there is a spontaneous reaction occurring. Electrons move from the negative electrode to the positive electrode. The power generated from the redox reaction powers the light bulb. So, chemical energy is converted into electrical energy.

Now, if we examine the electrolytic cell, we can see that electrons still flow from the anode to the cathode, but now there are two major differences. There is no more light bulb, and most importantly, the anode is now positive and the cathode is negative.

Let's try illustrating this concept with an example. Say we set up a Galvanic cell with a zinc anode and a copper cathode. The redox reaction would look something like this:

\begin{align}&E ^ {\circ} _ {cell} = E ^ {\circ} _ {cathode} - E ^ {\circ} _ {anode} \\&E ^ {\circ} _ {cell} = E ^ {\circ} _ { Cu ^{2+} / Cu } - E ^ {\circ} _ { Zn ^{2+} / Zn } \\&E ^ {\circ} _ {cell} = 0.34 ~ V - ( -0.76 ~ V ) \\&E ^ {\circ} _ {cell} = + 1.10 ~ V\end{align}

Copper (cathode) is being reduced and zinc (anode) is being oxidized. So, zinc is giving electrons to copper.

\begin{align}&E ^ {\circ} _ {cell} = E ^ {\circ} _ {Reduction} - E ^ {\circ} _ {Oxidation} \\&E ^ {\circ} _ {cell} = E ^ {\circ} _ {cathode} - E ^ {\circ} _ {anode} \\&E ^ {\circ} _ {cell} = E ^ {\circ} _ { Cu ^{2+} / Cu } - E ^ {\circ} _ { Zn ^{2+} / Zn } \\\end{align}

Say we let this cell react for a while. We originally had two strips of zinc and copper metal of equal size. After reacting for a while, the zinc strip is slightly smaller, and the copper strip is slightly larger. Their masses have changed because of the redox reaction. But what if we want to reverse the process; to oxidize copper and reduce zinc? Well, the reaction would look something like this:

\begin{align}&E ^ {\circ} _ {cell} = E ^ {\circ} _ {Reduction} - E ^ {\circ} _ {Oxidation} \\&E ^ {\circ} _ {cell} = E ^ {\circ} _ {cathode} - E ^ {\circ} _ {anode} \\&E ^ {\circ} _ {cell} = E ^ {\circ} _ { Zn ^{2+} / Zn } - E ^ {\circ} _ { Cu ^{2+} / Cu } \\\end{align}

Reduction will still occur at the cathode, and oxidation will still occur at the anode. However, now the anodic and cathodic species have changed. Zinc is now our cathode and will receive electrons. To force this reaction, some form of energy is required. But how much?

\begin{align}&E ^ {\circ} _ {cell} = E ^ {\circ} _ {cathode} - E ^ {\circ} _ {anode} \\&E ^ {\circ} _ {cell} = E ^ {\circ} _ { Zn ^{2+} / Zn } - E ^ {\circ} _ { Cu ^{2+} / Cu } \\&E ^ {\circ} _ {cell} = -0.76 ~ V - ( 0.34 ~ V ) \\&E ^ {\circ} _ {cell} = - 1.10 ~ V\end{align}

Since the spontaneous reaction generates +1.10 V of energy, it requires at least that much energy to drive the reverse reaction. By using electricity, we can force any reaction to occur, including depositing gold atoms onto a metal surface.

Cell Design

We learned that a Galvanic cell houses a spontaneous reaction, and electrons will spontaneously flow from the negative anode to the positive cathode. However, if we set up an electrochemical cell based on the previous diagram, this flow would stop almost instantaneously. Electrons move extremely fast, so an imbalance in charge would be reached almost immediately, and there would be no more flow.²

In this diagram, there is a salt bridge between each half-cell. The salt bridge is a tube filled with a strong electrolyte, with semi-permeable membranes on both ends. This allows ions to flow in and out of the tube. Let's break this down a bit.

In the anodic cell, Zn metal is oxidized to Zn²⁺ + 2e^-. The two electrons are flowing through the wire and into the copper electrode. The positively charged Zn²⁺ ions stay in the cell and combine with the ZnSO₄ solution. This generates an excess positive charge in the zinc solution.

In the copper cell, electrons are flowing into the electrode. Cu²⁺ will flow from solution and into the copper electrode to combine with the 2e^-. This leaves sulfate, SO₄^2-, in solution, which creates a negatively charged solution.

As the reaction proceeds, there is an immediate charge imbalance, which stops electron flow. However, with a bridge in the middle that contains ions, the imbalance can be reverted. Ions may flow into either of the electrolyte solutions to help restore balance. Without the salt bridge connecting the two cells, the reaction could not occur for any sustained period of time. Each piece of the Galvanic cell is crucial in completing the puzzle.

In an electrolytic cell, the salt bridge is not necessary. The electrodes can even be kept in the same electrolyte. Since electricity is being used to power the reaction, the cell conditions necessary for reaction are less strict. An electrolytic cell can take many forms, but the most important component is energy. Without sufficient energy, the nonspontaneous reaction will not occur, and nothing will happen.

In a Galvanic cell, a Voltmeter measures the amount of current travelling through the wire. However, in an electrolytic cell, there is no Voltmeter, but rather a source of energy. This could be electrical energy, solar energy as is the case with solar panels, or any other source of energy.

Example of Galvanic Cell and Electrolytic Cell

Let's now look at an example of a Galvanic cell and an electrolytic cell. Can you think of any examples of a Galvanic cell in everyday life? Some process that you utilize to provide usable energy. If you thought of a battery, you are absolutely right!

Batteries employ spontaneous chemical reactions to provide energy. Batteries come in all different shapes and sizes, but they are all just mini Galvanic cells. If you think of a small battery, like a AA or a AAA, it does not operate unless it is plugged into a device. Essentially, when it touches two metal plates, it allows the circuit to complete, which allows the spontaneous redox reaction to occur within the battery.

Now, how about an example of an electrolytic cell? Well, in the same vein of thought, you could say a rechargeable battery. This is a redox reaction that is reversed by applying a certain voltage of electricity.

One of the most common examples of electrolysis is the splitting of water into its molecular components.

$$ 2 H_2O \rightleftharpoons 2 H_2 + O_2 \qquad E^{\circ} = -1.23~V $$

1.23 V of energy is required to split water into oxygen and hydrogen gas. When > 1.23 V is applied to an electrolytic cell containing water, gas will bubble out of solution. This gas is either H₂ or O₂, depending on where it bubbles from. The splitting of water is a redox reaction because hydrogen is being reduced and oxygen is being oxidized. Since hydrogen is being reduced, it will bubble at the cathode. Conversely, as oxygen is being oxidized, it will bubble at the anode.

Okay, so that's all fine, but why do we care about this reaction? Who cares about splitting water?

Well, there are a few reasons, some of them beneficial, some of them not. To start with the negatives, this reaction means that electrolytic cells run in water have limits. If a redox reaction requires more than 1.23 V of energy to work, the reaction cannot be run in a solution containing water. If it is, the water will be split instead, and the preferred redox reaction will not likely occur. Therefore, for electrolytic cells to operate properly, we must carefully choose our electrolytes.

Now, the other, more exciting, reason that we care about this reaction, is that water can be split to provide hydrogen gas.

Similarities and Differences Between Galvanic and Electrolytic Cells

Hopefully, by now, we have established the theoretical and operational aspects of the electrochemical cell. It may be useful now to look at the similarities and differences between Galvanic and electrolytic cells. Let's try summarizing this in a table to help you visualize it.

Galvanic and electrolytic cells are very similar, they just move in opposite directions. This is why a rechargeable battery is so effective. It is one redox reaction that is allowed to move in one direction, and then forced to move in the other direction.

Now that we have discussed how to set up and how to operate electrochemical cells, imagine you are sent into the past to a time of alchemy and charlatans posing as mystics. Since you know how electrolytic cells are operated, the only question remains: who are you going to create gold for?