Galvanic Cell (Simplified) - Online Tutor, Practice Problems & Exam Prep

A galvanic cell, or voltaic cell, converts stored chemical energy into electrical energy through spontaneous reactions, featuring an anode (negative, where oxidation occurs) and a cathode (positive, where reduction occurs). Electrons flow from the anode to the cathode, aided by a salt bridge that maintains charge balance. In contrast, an electrolytic cell requires external electrical energy to drive non-spontaneous reactions, reversing the charge of electrodes. Both cells share the fundamental roles of oxidation and reduction, essential for understanding electrochemical processes and applications in batteries.

Galvanic Cell (Simplified) Concept 1

Now, a galvanic cell, also known as a voltaic cell, is a spontaneous cell that produces or discharges electricity, therefore making it a battery. Now, we're going to say that it uses stored chemical energy and converts it into electrical energy. With our galvanic or voltaic cell, we have 2 types of electrodes. We have our anode and our cathode. The anode itself represents our negatively charged metal electrode, and this is the compartment where oxidation occurs. So the anode is oxidation, and remember, with oxidation, we have the loss of electrons. So here with oxidation, we say we lose electrons. The cathode represents our positive metal electrode, and it's the department where reduction occurs. And remember, with reduction, we say we gain electrons, so we're gaining electrons here. Now, if we were to look here at this image of our galvanic cell, remember we said that our anode here is negative and our cathode here is positive. So here, this represents our anode, our metal electrode, and it looks like the metal electrode here is zinc, and our cathode, which is positive, is this metal electrode here, which is copper. Remember, we said that the anode loses electrons, so electrons are literally leaving this electrode and they're heading towards our cathode. So, we have the flow of electrons going from anode to cathode. So here's our anode here, and here is our cathode here. Now, next, we're going to say we have a salt bridge. Now, what's the whole point of a salt bridge? Well, the whole point of a salt bridge, if you've taken physics, if you haven't taken physics, is to help close the circuit of our galvanic or our voltaic cell. So here we have negatively charged ions heading the opposite direction. So here, the flow of like charges flowing in opposite directions helps to close itself. And we're going to say here that our salt bridge is just a tube that connects both half-cells to one another and allows for the flow of neutral ions. When you're an ion, you can be acidic, basic, or neutral. For the galvanic or voltaic cell to work properly, we need the ions to be neutral. They cannot be acidic or basic in nature. So if we take a look here within our salt bridge, which is this tube here, which connects the 2 together, we have our negative bromide ions heading the opposite direction of the electrons. The electrons are heading this way, and the negative ions have to head this way. So here, they're moving towards the anode compartment. Sodium ions also represent, neutral ions, they're heading towards the cathode compartment. And we're going to say here that this movement of electrons from the anode to the cathode is what's going to help generate electricity. And we read this by using a voltmeter. So the voltmeter is just a device that records the amount of electricity generated by the galvanic cell. Okay. So that's what we can say in terms of this image of our galvanic cell.

Galvanic Cell (Simplified) Example 1

The purpose of a galvanic cell is to a, purify solids, b, allow for only oxidation, c, generate electricity, or d, to consume electricity. Alright. So, we did not talk about purifying solids through the process of a galvanic or voltaic cell, so this would be out. Now, a galvanic cell utilizes a redox reaction, a spontaneous redox reaction, in order to generate or create electricity. Remember, redox reactions involve both reduction and oxidation. It wouldn't only be oxidation. And in defining and describing a galvanic/voltaic cell, I said that they produce or create electricity. So, they're generating electricity. To say they're consuming electricity is the opposite. Remember, a galvanic, which is the same thing as a voltaic cell, they're just batteries. They're making or discharging electricity. Okay? So, it's being used. They're just using it. So, they're a battery at the end of the day. So, the answer here will be option c, generate electricity.

Galvanic Cell (Simplified) Example 2

Now remember in our discussion of the galvanic cell we talked about the two metal electrodes involved. Here we had our anode and our cathode. Remember, at the anode we have the loss of electrons, and if we take a look here at the zinc compartment we know that it represented the anode, and we know what happened here was that zinc was oxidized to zinc ion. So what does that entail? Well, we know that the zinc metal electrode is losing electrons over time and as it loses the electrons it's going to produce zinc 2+ ions. So here we see the electron traveling away from the metal zinc electrode and producing as a result a zinc 2+ ion.

Now, what does this do over time? Well, you might hear from your professor that the anode dissolves away. Given enough time, over time our zinc is going to become smaller and smaller, so a portion of it will disappear. Pieces of it will go away because we're losing electrons and over time that adds up to mass loss.

What would the anode reaction look like? Well, here if our zinc solid is losing electrons, the electrons' loss would have to be on the products side, it's producing 2+ ions, ions would be aqueous in solution, and here are the two electrons it's lost:

Zn(s) → Zn (2+)(aq) + 2 e- Zn(s) → Zn (2+)(aq) + 2 e-

On the other side, the cathode is gaining electrons, so here the surface of the cathode is gaining electrons so over time the surface becomes more negative. Because the surface is getting a slight negative charge it's going to attract the positive copper 2+ ions floating within the solution. They're going to be attracted to the electrons and travel towards the surface and when they come into contact with the electrons on the surface they're neutralized, thereby becoming copper solid. Thus we're getting an encrusting effect on top of the copper metal electrode of more copper. You may hear that cathodes plate out, and that's because of the partial negative charges gain. You can have these metal cations in the solution adhering to the surface, thereby making the electrode itself bigger and bulkier over time.

Now, what does that look like in terms of a half reaction?

Cu (2+)(aq) + 2 e- → Cu(s) Cu (2+)(aq) + 2 e- → Cu(s)

Now, we can give the overall reaction from this, remember your electrons are always cancelling out, and then you bring down everything else:

Zn(s) + Cu (2+)(aq) → Zn (2+)(aq) + Cu(s) Zn(s) + Cu (2+)(aq) → Zn (2+)(aq) + Cu(s)

Now if we look at this example from what we just covered we can say, how many electrons are transferred between the zinc and copper electrodes in the Galvanic cell? Well, we can say here that two electrons are being canceled from both half reactions, that's because two electrons are being transferred. So here my answer would have to be option B.

So remember, the anode is where oxidation occurs, we're losing electrons, and given enough time the anode itself could start diminishing in size. The cathode over time gains electrons giving itself a negative surface, which thereby attracts the metal cations already dissolved within the solution. This causes the cathode over time to get a little bit bigger.

Okay, so just keep this in mind when looking at a typical galvanic cell with two metal electrodes.

Galvanic Cell (Simplified) Concept 2

An electrolytic cell is a non-spontaneous cell that utilizes electrolysis in order to operate. Now, electrolysis is when we have chemical reactions that consume external chemical, electrical energy in order to occur. No matter the cell, whether it be an electrolytic cell or a galvanic/voltaic cell, it doesn't matter. The cathode is always the site of reduction, and the anode is always the site of oxidation. However, because electrolytic cells are non-spontaneous, their signs will be different. Here, since the process is non-spontaneous, the cathode is now negatively charged and the anode is now positively charged. So, basically, our understanding here is that when it comes to electrolytic cells, they are the opposites of galvanic cells. So, just remember this fundamental idea that they are opposites of one another.

Galvanic Cell (Simplified) Example 3

It states, identify the location within an electrolytic cell where the loss of electrons will occur. Alright. So, loss of electrons means that we are undergoing oxidation. Now remember, it doesn't matter what type of cell we're dealing with, whether it be an electrolytic cell, a galvanic cell, or a voltaic cell, oxidation always occurs at the anode. Reduction always occurs at the cathode. So again, the loss of electrons will occur at the anode. That means that option b is the correct answer.

Galvanic Cell (Simplified) Concept 3

An electrolytic cell has to consume electrical energy in order to convert it into chemical energy. Now, in order to do this, it requires a battery. So, one major difference between an electrolytic cell and a galvanic cell is that a galvanic cell itself is a battery that generates electricity, and an electrolytic cell requires a battery in order to work. Now, when it comes to an electrolytic cell, we still say that the cathode is the place of reduction. So, electrons are still moving from the anode towards the cathode, and the atom itself, since it's losing electrons, is a site of oxidation. Now, one difference here with an electrolytic cell is that we have our cathode being negative and our anode being positive. Because of this, we can say that an electrolytic cell is non-spontaneous. It's doing something that it really shouldn't be doing. Negative electrons really don't want to head towards a place that's also negative. So those negative electrons are being forced to go to the negative cathode. That's why the battery is also required. It's the energy we need in order to force the electrons to go there.

Now, we're also going to say that our electrons are heading towards the cathode, and as before, when we talked about a galvanic cell, the same is true here. The negative ion is heading towards the anode compartment and the positive ion is heading towards the cathode compartment. Now, if we look at the cathode compartment, yes, it's negatively charged, but again the cathode is where reduction occurs. So here, our Tin (II) ion is being reduced to Tin solid. This would mean that our reaction would be: Sn2++2e−→Sn (solid). And then, for the anode, it's still where the site of oxidation occurs, so our copper solid is being oxidized to copper (II) ion. So here we have Cu(s)→Cu2+(aq)+2e−.

Now, remember our electrons will cancel out and we'll get our overall equation as Sn2+(aq)+Cu(s)→Sn(s)+Cu2+(aq). Now, what's the application of an electrolytic cell? Well, we can say that in an electrolytic cell, we can find examples of batteries such as double-A batteries, triple-A batteries, etc., and also, it's a defining feature of rechargeable lithium batteries. So, although we said that electrolytic cells can be thought of as the opposite of a galvanic cell, and that's true for a lot of reasons, we can say that what binds them all together is that the anode is the site of oxidation, and the cathode is the site of reduction. So, that's what our different electrochemical cells have in common. Other than that, there's a lot different between a galvanic cell and an electrolytic cell. So, keep in mind these features that we've talked about for this particular type of electrochemical cell.

Galvanic Cell (Simplified) Example 4

Here it says, which of the following is true about an electrolytic cell? It changes chemical energy into electrical energy. No. It does the exact opposite. It harnesses the electrical energy from batteries and uses it to make chemical energy. It uses a positive cathode. So remember, in this case, the cathode is negatively charged, not positive. It uses an electrical current to make a nonspontaneous reaction go. Electrolytic cells are nonspontaneous; they can't do the process naturally without the use of outside energy. So, they're going to have to siphon off this electrical energy from batteries, use that energy to make themselves go and occur. So this here is true. Here, it can't be all of the above because options a and b we found was not true. So here we're going to say electric cells are nonspontaneous; as a result, they require an outside energy source. Here it happens in the form of batteries. We harness that electrical energy and use it to convert it into chemical energy.

Which of the following statements is true for a salt bridge?

• A. Contains neutral atoms that interact with the ions in both half-cell compartments.
• B. Serves as a route through which ions can flow freely.
• C. Serves as the site of oxidation.
• D. Serves as the site of reduction.

Which of the following statements is TRUE for a voltaic cell, but FALSE for an electrolytic cell?

I. The flow of electrons is spontaneous.
II. Oxidation occurs at the anode.
III. Electrons flow from the anode to the cathode.

• A. Only I
• B. I and II
• C. II and III
• D. I, II, and III
• E. Only II

What is the balanced half reaction that occurs at the anode in the overall cell reaction of a voltaic cell?

3 MnO₄^–(aq) + 5 Fe(s) → 3 Mn²⁺(aq) + 5 Fe³⁺(aq)

• A. MnO₄^–(aq) + 8 H⁺(aq) + 5 e^– → Mn²⁺(aq) + 4 H₂O(l)
• B. 2 MnO₄^–(aq) + 12 H⁺(aq) + 6 e^– → 2 Mn²⁺(aq) + 3 H₂O(l)
• C. Fe(s) →Fe³⁺(aq) + 3 e^–
• D. Fe(s) →Fe²⁺(aq) + 2 e^–
• E. Fe²⁺(aq) → Fe³⁺(aq) + e^–

Here’s what students ask on this topic:

What is a galvanic cell and how does it work?

A galvanic cell, also known as a voltaic cell, is a type of electrochemical cell that converts stored chemical energy into electrical energy through spontaneous redox reactions. It consists of two electrodes: the anode and the cathode. The anode is the negative electrode where oxidation occurs (loss of electrons), and the cathode is the positive electrode where reduction occurs (gain of electrons). Electrons flow from the anode to the cathode through an external circuit, generating electricity. A salt bridge connects the two half-cells, allowing the flow of neutral ions to maintain charge balance. This setup ensures the continuous flow of electrons, enabling the cell to produce electrical energy.

What is the role of the salt bridge in a galvanic cell?

The salt bridge in a galvanic cell serves to maintain electrical neutrality within the internal circuit, which is essential for the continuous flow of electrons. It is typically a tube filled with a salt solution that connects the two half-cells. The salt bridge allows the movement of neutral ions between the half-cells: negative ions move towards the anode, and positive ions move towards the cathode. This movement counteracts the buildup of charge that would otherwise stop the flow of electrons, ensuring the cell can continue to generate electricity.

How do the anode and cathode differ in a galvanic cell?

In a galvanic cell, the anode and cathode have distinct roles and characteristics. The anode is the negative electrode where oxidation occurs, meaning it loses electrons. For example, in a zinc-copper galvanic cell, zinc acts as the anode and undergoes oxidation: Zn → Zn²⁺ + 2e^-. The cathode is the positive electrode where reduction occurs, meaning it gains electrons. In the same cell, copper acts as the cathode and undergoes reduction: Cu²⁺ + 2e^- → Cu. Electrons flow from the anode to the cathode through an external circuit, generating electrical energy.

What is the difference between a galvanic cell and an electrolytic cell?

The primary difference between a galvanic cell and an electrolytic cell lies in their energy requirements and electrode charges. A galvanic cell generates electrical energy from spontaneous chemical reactions, with the anode being negative and the cathode positive. In contrast, an electrolytic cell requires external electrical energy to drive non-spontaneous reactions, with the anode being positive and the cathode negative. While both types of cells involve oxidation at the anode and reduction at the cathode, the direction of electron flow and the nature of the reactions differ significantly.

How do you measure the electrical output of a galvanic cell?

The electrical output of a galvanic cell is measured using a device called a voltmeter. The voltmeter is connected to the external circuit of the galvanic cell, with one lead attached to the anode and the other to the cathode. It measures the potential difference (voltage) between the two electrodes, which is a direct indication of the electrical energy generated by the cell. The voltage depends on the specific materials used for the electrodes and the concentration of the electrolyte solutions.