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Galvanic and Electrolytic Cells

Galvanic and Electrolytic Cells

Imagine this: you are a 14th century charlatan masquerading as an alchemist, and you've just discovered the formula to create gold. You are proclaimed a hero and assigned by the king to create more gold at once. Unfortunately for you, and for your livelihood, no matter what you combine, gold will not come out. However, instead, let's imagine that you are a chemist from the future. Clever as you are, you set up a solution of gold ions, a block of metal, and a battery. To the king's astonishment, you pull out a gold bar, where the metal once was. What you've just done is used an electrolytic cell to coat the metal bar with a thin layer of gold. The king keeps you on as his personal alchemist, and fortunately for you, you get to keep your head.

So, in this scenario, what kind of magic did you just perform? Well, this was done using a process called electrolysis. Here we will learn all about electrolytic and Galvanic (Voltaic) cells. That way, if you ever get sent 700 years into the past, you'll have a seat at the royal table.

    • First, we will take a look at what Galvanic and electrolytic cells are.
    • Next, we will discuss some examples of Galvanic and electrolytic cells.
    • Finally, we will summarize the similarities and differences between each cell.

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.

In an electrochemical cell, a redox reaction occurs, which produces energy. This may host either spontaneous or nonspontaneous reactions, depending on the cell conditions.

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.

Galvanic and Electrolytic Cells galvanic and electroyltic cell StudySmarterFig. 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, SO42-, 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 the salt bridge, a strong electrolyte, like KNO3, is suspendend in a gel. Negative ions flow into the anode to reduce the build-up of positive charge, and positive ions flow from the bridge into the cathode to neutralize the negative charge.

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.

The oxidation of metals (corrosion) is a spontaneous, and often undesirable, redox reaction which is seen everywhere. A perfect example of this is observed any time you encounter rust. Rust is the oxidation of iron metal. When iron oxidizes, the iron turns into flakes, which then break off and destroy the integrity of the component.

There are many techniques to prevent this from happening, but one popular method is known as galvanizing. Galvanization is the process of coating iron with another metal, such as zinc. Zinc acts as a sacrificial anode, which is oxidized instead of the iron. Zinc metal doesn't flake when it is oxidized, so it forms a protective zinc oxide layer over the iron.

Galvanization is typically done by dipping the iron in molten zinc, but this may also be done with electroplating. Electricity is used to reduce metal ions onto an iron (or steel) surface. Instead of zinc ions being deposited onto a solid zinc electrode, they are deposited onto an iron electrode. The zinc electrolyte solution serves as the anode, and the iron electrode serves as the cathode. Electroplating can be done with many different metals serving as either the anode or cathode.

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.

It is no secret that gasoline vehicles are causing a lot of greenhouse gas emissions. As a result, a lot of money is being invested into vehicles run on renewable energy. One of the more popular types is electric vehicles, which use gigantic (and very heavy) rechargeable batteries to operate. These typically use lithium-ion batteries, which are very effective. However, lithium is not renewable and will soon be in short supply. One of the alternatives to batteries is the hydrogen fuel cell. This cell combusts hydrogen, instead of gasoline, which produces water as the sole byproduct.

Galvanic and Electrolytic Cells hydrogen fuel cell example StudySmarterFig 3. Hydrogen fuel car - A vehicle which uses hydrogen as its fuel source would emit water instead of carbon dioxide and other harmful gases.

While the cartoon is a simple diagram of a hydrogen fuel cell, it perfectly highlights everything we have learned so far about Galvanic and electrolytic cells. Hydrogen fuel is combusted, forming hydrogen cations, H+. These ions can then combine with O2 from the atmosphere. Hydrogen oxidation serves as the anode, and oxygen reduction serves as the cathode. A wire connects each half-cell which carries electrons, and a membrane in the middle separates them and transports hydrogen ions. Now, what do you think? Is this a Galvanic or an electrolytic cell?

The vehicle run by hydrogen fuel is the ultimate green energy goal for the future. It would help reduce global carbon emissions, and would also use a renewable resource. While gasoline is still abundant, global reserves are beginning to deplete, and will not last long at our current rate of consumption. Within our lifetimes, it is possible that the hydrogen fuel cell could replace the gasoline engine.

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 Cell
Electrolytic Cell
Hosts a spontaneous reactionHosts a nonspontaneous reaction
Oxidation occurs at the anodeOxidation occurs at the anode
The anode provides electrons and has a negative chargeThe anode pulls electrons and has a positive charge
Reduction occurs at the cathodeReduction occurs at the cathode
The cathode pulls electrons and has a positive chargeThe cathode provides electrons and has a negative charge

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?

Galvanic and Electrolytic Cells - Key takeaways

  • A Galvanic cell hosts a spontaneous reaction which produces chemical energy.
  • An electrolytic cell hosts a nonspontaneous reaction which requires electrical energy to operate.
  • In all electrochemical cells, oxidation occurs at the anode, and reduction occurs at the cathode.
  • In a Galvanic cell, the anode is negative, and the cathode is positive. In an electrolytic cell, the anode is positive, and the cathode is negative.
  • Rechargeable batteries host a redox reaction which is spontaneous, and then reversed with electricity.
  • Green fuel, such as hydrogen, can be produced by using an electrolytic cell to split water. Then it can be combusted in a Galvanic cell to provide green energy.

References

  1. Fig 1. - Electrochemical Cells (https://commons.wikimedia.org/wiki/File:Two_types_of_cells,_galvanic_and_electrolytic.jpg) by the UC Davis Library is licensed by CC BY-SA 4.0 (https://creativecommons.org/licenses/by-sa/4.0/deed.en).
  2. Nivaldo Tro, Travis Fridgen, Lawton Shaw, Chemistry a Molecular Approach, 3rd ed., 2017.

Frequently Asked Questions about Galvanic and Electrolytic Cells

A Galvanic cell uses a spontaneous reaction to convert chemical energy into electrical energy. An electrolytic cell uses electrical energy to drive a nonspontaneous reaction, creating stored chemical energy.

A Galvanic cell hosts a spontaneous reaction, while an electrolytic cell hosts a nonspontaneous reaction. A Galvanic cell can be imagined as a battery operating, and an electrolytic cell is the reverse of that reaction.

In both Galvanic and electrolytic cells, the electrodes are chemically transformed through a redox reaction. The anode will be oxidized, and the cathode will be reduced.

A rechargeable battery can be visualized as both a Galvanic and an electrolytic cell in one. The battery is the Galvanic cell, which is a spontaneous reaction. The rechargeable part is the electrolytic cell, which reverses the spontaneous reaction. This allows the battery to operate once again.

Galvanic and electrolytic cells are distinct because they operate different directions of a reaction. A Galvanic reaction hosts the spontaneous reaction, and an electrolytic cell hosts the nonspontaneous reaction.

Final Galvanic and Electrolytic Cells Quiz

Question

What is an electrochemical cell?

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Answer

An electrochemical cell hosts either a spontaneous or nonspontaneous reaction.

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Question

What is a Galvanic cell?

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Answer

A Galvanic cell is a cell which hosts a spontaneous reaction.

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Question

In a Galvanic cell, what are the charges on the anode and cathode?

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Answer

The anode has a negative (-) charge, and the cathode has a positive (+) charge.

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Question

A galvanic cell can only function if ______

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Answer

There is a salt bridge connecting the two half cells.

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Question

An electrolytic cell can only function if _____

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Answer

There is sufficient energy provided to force the reaction to occur.

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Question

What are electrolytic cells?

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Answer

Electrochemical cells which host a nonspontaneous reaction.

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Question

To make a nonspontaneous redox reaction take place by exploiting electrical energy (electrolysis), which type of electrochemical cell is used?

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Answer

Electrolytic cell.

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Question

Which type of electrochemical cell will convert chemical energy into electrical energy?

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Answer

A Galvanic cell.

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Question

In an electrolytic cell, the cathode is the ___ pole and the anode is the ___ pole.

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Answer

negative; positive.

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Question

In a galvanic cell, oxidation occurs at the ____ and reduction occurs at the ____.

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Answer

anode; cathode.

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Question

Which type of cell will convert electrical energy into chemical energy?

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Answer

An electrolytic cell.

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Question

What is a potential green source of hydrogen gas?

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Answer

Water splitting.

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Question

Water splitting can be achieved by which type of electrochemical cell?

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Answer

A Galvanic cell.

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Question

In an electrochemical cell, what reaction occurs at the anode?

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Answer

Oxidation occurs at the anode.

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Question

In an electrochemical cell, what type of reaction occurs at the cathode?

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Answer

Reduction occurs at the cathode.

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Question

In an electrolytic cell, oxidation occurs at the ____ and reduction occurs at the ____.

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Answer

anode; cathode.

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