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Ch.1 - Intro to General Chemistry
Ch.2 - Atoms & Elements
Ch.3 - Chemical Reactions
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Ch.4 - Chemical Quantities & Aqueous Reactions
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Redox Reaction
Balancing Redox Reaction
The Nernst Equation
Faraday's Constant
Galvanic Cell
Batteries and Electricity
Additional Guides
Nernst Equation (IGNORE)
Galvanic (Voltaic Cells) & Electrolytic Cells

Concept #1: Galvanic versus Electrolytic Cells

Transcript

Hey guys, in this new video, we're going to take a look at the difference between a galvanic cell versus a voltaic cell. The first thing we're going to talk about before we talk about the differences between these two cells is first we're going to say with redox reactions, we will now deal with a new variable. It's called the reaction’s cell potential. It uses the variable E. We have E without the circle that means it's under non-standard conditions. Then we have E with a circle, which means we are under standard conditions.
Standard conditions usually refer to us having a molarity of 1 or having a pressure of 1, or having a pH that's close to neutrality, so pH equals 7. When we have these conditions, we are talking about cell potential under standard conditions. We're going to say here, the greater the variable, then the more likely reduction will occur. The smaller this variable, then the more likely oxidation will occur. Remember also that if we know about reduction and oxidation, then we know which electrode we're talking about. Remember, reduction predominantly happens at the cathode and oxidation happens at the anode.
Now, let's first talk about the galvanic cell. In a galvanic cell, we have the creation of electricity. So galvanic cells basically, they either produce electricity or they discharge electricity over time. Basically, a galvanic cell can also be called a voltaic cell. Both of them because they produce electricity, they are just batteries. They are a spontaneous cell because they create this electricity on their own. The opposite of a galvanic or voltaic cell is called an electrolytic cell. An electrolytic cell does not produce or discharge electricity, instead it uses electricity to happen. Here it consumes electricity. We’d say here that an electrolytic cell, it is non-spontaneous. It requires an outside energy source in order for it to occur.
What we have here is an example of a basic galvanic cell, galvanic or voltaic cell. It's based off of these two equations here. We need room to work this out guys so let me take myself out of the image.
Now we're going to say generally speaking, the larger your E value, again reduction is more likely to occur so it's the cathode. Right here, this is the cathode. The smaller your cell potential, the more likely oxidation occurs so it's the anode. Based on these two equations, we're saying here that the nickel solid is our cathode. When it comes to a galvanic or voltaic cell, the positive electrode is the cathode. So we have nickel here, nickel solid, and then here cadmium solid represents my anode. Based on these two half-reactions, the nickel has and Ni2+ ions floating around in solution. Then here the cadmium has Cadmium two plus ions floating around in solution.
We're going to say here in a galvanic or voltaic cell, the anode is negative. Now what's happening here is oxidation remember means we're losing electrons. Over time, electron start to leave the anode so they're heading this way towards this electrode here. Here it's producing electricity, so right here we have this something called a voltmeter which we'll talk about how much voltage is being produced later. Here there are two periodic trends you need to keep in mind when it comes to this process. We have ionization energy and we have electron affinity. Remember, cadmium is the anode. It’s undergoing oxidation. Here, ionization energy is the energy to remove an electron. We have the cadmium solid losing electrons to become cadmium 2+. Here go the two electrons that's lost.
This is a spontaneous reaction which means the electrons come off the anode very easily. The amount of energy required is very low. On the opposite end, electron affinity means the energy involved in adding an electron. In this case, we have nickel 2+ ion gaining two electrons to become nickel solid. Here, again this is a spontaneous reaction, so this happens naturally. We're going to say here the propensity or the likelihood of nickel gaining those electrons is very high. The electron affinity here would be high. Nickel wants to gain those electrons.
What's happening here over time what we should realize is this cathode metal here is gaining electrons. Over time, the surface of the cathode electrode is becoming more and more negative. Imagine if the surface is becoming more negative, what do you think those positive ions dissolved in the solution are going to do? They're going to start to connect themselves to this negative surface. Positive ions and negative surface combine together. What's going to start to happen is they're going to in crossed themselves onto this cathode. Over time, the cathode is going to get fatter. It's going to get bigger. We're going to say here we're going to have the cathode plates out. Meaning, it's going to get bigger because those positive ions that are dissolved in solution will connect to the surface of the cathode metal.
At the same time, the anode is losing more and more electrons. Over time, the anode piece of metal gets smaller and smaller. More and more of it is going to dissolve away. The anode overtime is going to shrink that's why we say the anode dissolves away. But this process can't happen without the use of this structure here. This is called your salt bridge. Your salt bridge contains inert ions. It contains usually chloride ions and nitrate ions. They're usually connected to maybe sodium or potassium. Remember, we say that chloride ion and nitrate ion are inert negative ions. Why are they inert? Because they come from strong acids. Remember, if you come from a strong acid, you're a very, very weak conjugate base so weak that you're neutral.
What's happening here is we have these negative ions basically traveling from the cathode solution towards the anode solution. What's happening here is those negative ions are connecting with the ions in the solution and neutralizing those cadmium2+ ions. We have cadmium maybe combining with the chloride ions if they're in there to form cadmium chloride. Or we have the cadmium ions combining with the nitrate ions to become cadmium nitrate. This is important because the salt bridge, it neutralizes the positive ions in the anode chamber. But more importantly, they help to complete the circuit. For those of you who have taken physics, remember we have to have the movement of light charges in opposite directions to complete an electrical circuit.
We have negative electrons traveling this way from the anode to the cathode. But at the same time from the cathode side to the anode side, we have those negative ions. I'm just putting it as NO3 minus but it could also be chloride ions. You need these same charges, negative charges, to flow in opposite directions to complete the circuit. The bridge has two things that it's doing. It’s neutralizing the positive ions in the anode chamber and at the same time helping to complete the circuit so that the galvanic or voltaic cell helps to produce electricity.
Here talking about producing electricity, how could we help to make more electricity happen? Remember, I told you that we want to neutralize these positive ions here in the anode chamber because if they get too high, too large, it's going to disrupt the electrical flow. You want to make sure that the concentration of the anode ions is low. At the same time, we need electrons to flow from the anode to the cathode. The electrons are more likely to flow towards the cathode if there's a buildup of positive ions within the solution. You want to make sure that the cathode’s concentration is high, thereby attracting more electrons to that side. This helps with the like charges to move in opposite directions. This is the major points that come to dealing with the galvanic or voltaic cell.
What's the difference between this and electrolytic cell? Pretty much a lot of things. We're going to say here for example, remember an electrolytic cell is non-spontaneous. An electrolytic cell we're going to say here we still have the movement of electrons from the anode to the cathode, but in an electrolytic cell, the anode is actually positive and the cathode is actually negative. Think about it. Why is it non-spontaneous? Negative electrons are traveling from the anode to the cathode still, but think about it. Why would negative electrons want to travel to a negative electrode? They wouldn't want to go that way that's why we need the use of a battery. Instead of having a voltmeter up here, it would have actually a battery, so you’d place a battery on top. That battery would help to force the electrons to leave the positive anode and head towards the negative cathode. That's one huge difference between a spontaneous cell versus a non-spontaneous cell.
How do we figure out the amount of voltage being produced in this galvanic or voltaic cell? Remember, cell potential which is E cell equals cathode minus anode. Here, we’re going to say that your cathode is negative 0.25 volts minus a minus 0.40 volts. Remember, minus of a minus is really a positive. That's point 0.15 volts. You’re going to place 0.15 volts in here. Here we just found out that a galvanic or voltaic cell is spontaneous and it has a positive cell potential.
What kind of effect does that have? That's where we move down to this list right here We have four variables we’re looking at. We’re looking at the entropy of the universe, which deals with the second law of thermodynamics. We're looking at Gibbs free energy. We're looking at your equilibrium constant and we have your cell potential. When we have them all in these signs, positive and negative and greater than 1 and positive, it is a spontaneous reaction. If you're a spontaneous reaction, what kind of cell are you? You’re again a galvanic or voltaic cell. Then here we're going to say if you're the complete opposite, you’re non-spontaneous and so you are an electrolytic cell.
Then let's say you're at 0, 0, 1, and 0, here you're not spontaneous, you're not non-spontaneous. You are at equilibrium. What do you represent then? You represent a dead battery in this case. Going through all the different things that we've seen, remember a galvanic and voltaic cell is spontaneous, ionization energy is low. For the anode, electron affinity is high. For an electrolytic cell, it’s the exact opposite. For an electrolytic cell, ionization energy would be high because electrons don’t want to leave a positive anode. Electron affinity would be low for an electrolytic cell. Why? Because electrons don't want to go to a negative electrode, such as the cathode when we're dealing with an electrolytic cell.
Basically, if you can remember the steps for a galvanic/voltaic cell, the electrolytic is almost the exact opposite. The only place that they agree is that anodes are always oxidized. They always lose electrons to help reduce the cathode. That's the only thing they really have in common with one another. Remember, this whole process is only possible with the help of the salt bridge which has negative ions flowing from the cathode portion to the anode portion to help neutralize those ions. Also, we do have some leakage and some seeping out of these positive ions that can help to go towards the cathode side as this surface becomes more and more negative over time. Remember that the anode dissolves away as the cathode builds up or plates out. Of course, remember cell potential is cathode minus anode to help you find your cell potential. 

Concept #2: An electrolytic cell presents an electrochemical cell that is nonspontaneous. 

Example #1: A certain electrochemical cell involves a five electron change and has a value of Keq = 3.0 x 1016 at 298 K. The value of ΔHo for the reaction is -68.3 kJ/mol. Calculate the values of ΔGo, ΔEo, for a standard electrochemical cell constructed based on this reaction and also ΔSo for the reaction.

Practice: Given the following standard reduction potentials,

Hg22+(aq)  +  2 e    2 Hg (l)                                                          E° = +0.789 V

 

Hg2Cl2(s)  +  2 e    2 Hg (l)  +  2 Cl-(aq)                                        E° = +0.271 V

determine Ksp for Hg2Cl2(s) at 25 °C.

Example #2: The cell notation for a redox reaction is given as the following at (T = 298 K):

Zn (s) Ι Zn2+ (aq, 0.37 M) ΙΙ Ni2+ (aq, 0.059 M) Ι Ni (s)

a)  Write the balanced half-reactions occurring at the anode and the cathode.

b)  Write out the complete balanced redox reaction.

c)  Determine the Ecell.

d)  Calculate the maximum electrical work that can be produced by this cell.

e)  Calculate the reactant quotient, Q, for this cell and the cell potential under non-standard conditions.

 

Example #3: Answer each of the following questions based on the following half reactions:

HALF REACTIONS                                Eo (V)

Cl2 (g)  +  2 e            2 Cl (aq)          + 1.36

l2 (g)  +  2 e              2 l (aq)             + 0.535

Pb2+ (aq)   +  2 e      Pb (s)                  - 0.126

V2+ (aq)  +  2 e         V (s)                    - 1.18

a)  Which is the strongest oxidizing agent?

b)  Which is the strongest reducing agent?

c)  Will I (aq) reduce Cl2 (g) to Cl (g)?