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EngrSajjad
  • 6/9/2025
Transcript
00:00Hello everyone. Welcome to this exciting lecture series on electrochemistry. In today's lecture,
00:09we are going to study about electrochemical and electrolytic cells.
00:15So, let us dive into the fascinating topic of electrolytic and electrochemical cells,
00:22which are key devices in the field of electrochemistry. First of all,
00:27electrochemistry is a branch of chemistry that specifically studies how electrical energy and chemical reactions interact.
00:35Two fundamental devices here are electrolytic cells and electrochemical or galvanic cells.
00:42In this slide, we are focusing on electrolytic cells.
00:47Electrolytic cells are used when we want to force a non-spontaneous reaction to occur.
00:54This means that the reaction would not happen on its own without help.
00:59So, we supply that help using an external power source such as a battery.
01:05We can see from here.
01:08The cell includes two electrodes, the anode where oxidation happens and the cathode where reduction takes place.
01:17And these electrodes are placed in an electrolytic solution that contains free-moving ions.
01:24When we apply an electric current, the ions in an electrolyte starts to migrate.
01:31Cations, which are the positive ions, move toward the cathode to generate electrons and they are reduced.
01:40And ions, which are negatively charged, move toward the anode to lose electrons, that is the oxidation process.
01:49This setup is commonly used in processes like electroplating, electrolysis of water and metal refining, where we want to break down compounds or deposit metals using the electricity.
02:04So, electrolytic cells are powerful tools to carry out chemical changes that would not happen on their own.
02:12Let us further explore the process of electrolytic cells with the help of an example.
02:22In this example, we will study the electrolysis of a molten sodium chloride, which is a great example of how electrical energy can be used to drive a non-spontaneous chemical reaction.
02:34We begin with a source of direct current connected to a pair of inert electrodes, which are immersed in molten sodium chloride.
02:45The electrodes are inert, meaning they do not take part in the reaction themselves.
02:50They are just there to provide a surface for a reduction to happen and that the sodium chloride must be in its molten state, not the solid state, so that the ions are free to move in the cell.
03:05Now, here is what happens when the current is applied.
03:09In molten form, the sodium mines flow towards the negatively charged electrodes and the chloride ions flow towards the positively charged electrodes.
03:21We can simply see from here, sodium mines moving towards the cathode and chloride ions moving towards the anode.
03:28At cathode or the negative electrode, the following important steps take place.
03:36When sodium mines collide with a negative electrode, the battery carries a large enough potential to force these ions to pick up electrons to form a sodium metal.
03:47That is, a reduction reaction happens when sodium ions gain electrons to become neutral sodium ions.
03:55So, here we can see that this reaction can be represented as Na plus the gain electron to form sodium metal.
04:06So, at the negative electrode, which is a cathode, sodium metal is formed by the gain of electrons.
04:14This is how pure metallic sodium is produced from the molten sodium chloride using the process of electrolysis.
04:22Now, now that we have seen the negative electrode, let us move to the positive electrode known as the anode.
04:35As the electrolysis continues,
04:39These chloride ions that collide with a positive electrode are oxidized to give a chlorine gas, Cl2 gas, which bubbles off at this electrode.
04:50This is a classic oxidation reaction where each chloride ion loses an electron.
04:56The electrons are pulled away by the power source and the two chloride ions combine to form a single molecule of chlorine gas.
05:05The half equation for this oxidation reaction is that at the positive electrode, two chloride ions, they give up their two electrons and form a Cl2 gas, which bubbles out of the cell.
05:20So, we are seeing the loss of electrons here and due to the loss of electrons, oxidation happens at the anode.
05:32And now, putting both half reactions together, let us consider what is happening overall.
05:39The net effect of passing an electron through the molten sodium chloride in the cell is to decompose sodium chloride into its elements, which are the sodium metal and the chlorine gas.
05:53It is a clear demonstration of how electrolysis can break down a compound into its elements by using electrical energy.
06:09Now that we have seen what happens in the cell, let us talk briefly about the word itself, which is the electrolysis.
06:18The example given above explains why the process is called electrolysis.
06:23The name electrolysis is not just a random name.
06:27It actually tells us something about the process.
06:31The suffix lysis comes from the Greek stem meaning to loosen or split up.
06:39And that is exactly what we are doing here, splitting up a compound, as we have seen in the previous example.
06:48So, when we say electrolysis, we literally mean splitting with the electricity.
06:55That is the beauty of it.
06:57Electrolysis uses an electric current to split a compound into its respective element.
07:04As we can see here, in this case, we are taking molten sodium chloride and using electrical energy to break it down into liquid sodium metal and gaseous chlorine.
07:17Here we can see that we obtained liquid sodium metal here and the chlorine gas here.
07:22This process is a clean, elegant process that transforms a stable compound into its pure elemental compounds.
07:36Now, let us take a quick look at the dotted vertical line in the center of this cell.
07:44This represents the diaphragm.
07:47Now, what is the purpose of this diagram?
07:50Well, it keeps the Cl2 gas produced at the anode from coming into contact with the sodium metal generated at the cathode.
08:00This separation is really important because if the two products were to mix, they could react with each other and undo all the hard work of electrolysis.
08:10The function of this diaphragm can be understood better if we look at a more detailed and realistic drawing of the cell.
08:26But, for now, just remember that it acts like a barrier or divider to maintain the product's purity and prevent unwanted side directions inside the electrolysis.
08:39Okay.
08:40It is one of the small components that plays a big role in making the process efficient and safe.
08:47Now, here we can see that there are two cells.
08:54This is the real cell that is the actual process for the molten sodium chloride to produce sodium and the chlorine gas.
09:04So, in this loud slide, let us talk about the down cell which is a real world application of the electrolysis of molten sodium chloride.
09:14Okay.
09:15First of all, we see that the concept that we have studied in the previous slide is that inside the cell, there is a diaphragm made of iron gas.
09:26Here it is in the divided line, but in the cell, we can say here in the dotted line iso also.
09:34So, this at the two sides is the diaphragm.
09:39It is basically in the round form, but because this is a 2D picture, we can show here by only the dotted lines.
09:48So, this is the iron screen which is basically iron gas and it serves the purpose for this diaphragm.
09:55Okay.
09:56It is actually a fine mesh screen which separates the two electrodes basically and helps prevent the product from the mixing.
10:10Next, we can see that anode is at the center and it is made of the graphite.
10:17Also, at the sides, we can see these two things.
10:23These are the cathodes.
10:24Okay.
10:25Both of these.
10:26And these cathodes are separated from this anode using this diaphragm.
10:31Next, we see that in this reaction, the feedstock for the down cell is not just pure sodium chloride.
10:42Instead, it is a 3 to 2 mixture by mass of calcium chloride and sodium chloride.
10:49This combination is used for practical reasons.
10:55Okay.
10:56You see, the pure sodium chloride has to be heated to more than 800 degrees centigrade before it melts
11:04because that is the melting point of sodium chloride, which is a pretty energy intensive process.
11:11But, by adding calcium chloride, we can lower the melting point to about 580 degrees centigrade,
11:18making the process more efficient and less costly.
11:23Let us now complete our understanding of how the down cell operates during the electrolysis.
11:34Inside the cell, calorine gas forms on the graphite anode, which is inserted into the bottom of the cell.
11:41As the reaction proceeds, the calorine gas bubbles up through the molten sodium chloride,
11:46okay, and it is collected in a funnel at the top.
11:50It is a clever way to safely collect the gaseous products as it forms.
11:54So, the calorine leaves from this top pipe and from the system.
12:01At the same time, sodium metal forms at the cathode.
12:05So, these two are the cathodes.
12:07These are placed along the inner side of the cell.
12:10Because molten sodium is less dense than the molten sodium or molten salt,
12:15it floats up through the molten sodium chloride.
12:18So, it floats up through the molten sodium chloride solution.
12:22And it is gathered in a sodium chloride collecting ring.
12:26So, due to its low density, it rises up and we can simply collect the sodium metal from here.
12:32Direction is at the electrodes are simple but powerful.
12:39At anode, we have, as we have seen earlier,
12:42calorine gives office its two electrons to provide us Cl2 gas.
12:48And at cathode, Na plus signs gain an electron to produce sodium metal.
12:53So, overall, the reaction is NaCl.
12:57This will give two Na atoms and a calorine gas.
13:02So, we have seen that this is a great example of electrolysis in action.
13:07Which efficiently separates molten sodium chloride into its two elements, sodium, metal and calorine,
13:13through the use of electric current and a clever cell design.
13:17So, we have studied and understand the electrolytic cell.
13:27Now, we are going to study about the electrochemical cell.
13:31Which is the second type of our cell.
13:34So, what is an electrochemical cell?
13:40Electrochemical cells, such as galvanic cells or voltage cells,
13:45are devices that generate electrical energy from spontaneous chemical reactions.
13:50These are the reverse of what we see in the electrolysis and electrolytic cells.
13:56They basically convert chemical energy into electrical energy by using a redox reaction.
14:02In the process, electrons flow from the anode where oxidation takes place to the cathode where reduction takes place.
14:10And this flow of electrons is what gives us the electricity.
14:15It is this clever manipulation of redox reaction that powers everything from simple batteries to large-scale energy systems that we see around our cells.
14:27So, in the coming slides, we will study in detail about the electrochemical cells.
14:40Now, we will talk about the components of an electrochemical cell.
14:44We can see a diagram of the electrochemical cell here.
14:48But, we will see the components and some of their functions one by one in the coming slides.
14:55Let us begin with the anode, which is a key part of any chemical cell.
15:00The anode is a site where oxidation takes place.
15:05Meaning, it is where electrons are losing from a substance during a redox reaction.
15:12Now, the sign of the anode depends upon the type of cell.
15:20In galvanic cell, the anode is negative.
15:23That is because the substance at the anode, like zinc in a typical cell, gets oxidized and release electrons.
15:31For example, Z then gives off two electrons to produce 2 plus sign.
15:36The electrons accumulate the electrode, making it a negative.
15:42So, the site where oxidation occurs is the site called an anode.
15:49On the other hand, in an electrolytic cell, the anode is positive.
15:54This is because the external power supply is pushing electrons out of the anode, forcing a non-spontaneous reaction to happen.
16:05The electrode is being drained of electrons.
16:08So, it takes on a positive charge.
16:11Also, a reduction happens on the anode in an electrolytic cell.
16:21Okay.
16:22So, despite these differences in sign, one rule is always holds true.
16:29Oxidation always happens at the anode, whether the cell is spontaneous or not.
16:34Next, we will see about the cathode.
16:41The cathode is a site of reduction, which means it is where the gain of electron happens.
16:48In other words, chemical species such as metal lines accept electron at this electrode and undergo reduction.
16:56In galvanic cell, the cathode is positive because it receives electrons that flow spontaneously through the external circuit from the anode.
17:05This positive charge attracts the ions that gain electrons.
17:09In contrast, in an electrolytic cell, the cathode is negative.
17:15This is because the reaction is non-spontaneous and is driven by an external power source that forces electrons toward the cathode.
17:24A good example of the reduction reaction at the cathode is given here.
17:28CO2 plus signs, they gain two electrons to form solid copper metal and this reaction happens at the cathode.
17:41Next, we will study about the electrodes.
17:44Electrodes are just the metal parts which are immersed in the solution and they provide a platform for the reduction reactions or the oxidation reactions to happen.
17:55Okay.
17:56So, electrodes can be of two types.
17:58They can be active or inert electrodes.
18:02Active electrodes participate directly in the redox reaction during a chemical reaction in the cell.
18:09They must be solid and electrically conductive metals like zinc or copper.
18:14For example, in a daniel cell, the zinc electrode loses mass as it oxidizes while the copper electrode gains mass due to reduction.
18:24Secondly, inert electrodes, on the other hand, they do not participate chemically in the redox reaction.
18:38Instead, they provide a surface for the electron transfer when no solid redox material is present in the solution.
18:45Okay.
18:46Common examples of inert electrodes include platinum and graphite.
18:58These are often used in systems such as hydrogen electrodes, H2 gas or iron redox couples like Fe3 plus and Fe2 minus where the redox species are in the solution.
19:10The choice between inert and active electrodes depends on the specific electrochemical system and whether the electrode material itself undergoes chemical change during the reaction or not.
19:27The next component that we are going to study is the electrolyte.
19:33The electrolyte is a conductive ionic solution that allows the flow of charge between the electrodes by enabling the movement of ions within the cell.
19:44Okay.
19:45It plays a key role in maintaining electrical neutrality by allowing positive and negative ions to migrate and balance the electron flow happening during the cell.
19:57The electrolyte must be carefully chosen to be chemically compatible with the electrode materials.
20:05This helps avoid unwanted side reactions such as precipitation or electrode corrosion.
20:12For example, in copper electrochemical system, sodium sulfate is often used as electrolyte with copper sulfate solution because it prevents the precipitation of unwanted compounds and ensures smooth flow of ions.
20:32Next, we will study about the external circuit.
20:35Next, we will study about the external circuit.
20:36So, external circuit is just a potentiometer which receives the power from the cell.
20:42It can be any resistance or our light source.
20:47Okay.
20:48Nothing that we want to run.
20:50It can be anything.
20:51But also, the external circuit provides a pathway for the electrons to move from the anode to the cathode.
21:01Okay.
21:02The electrons travel from here to here through this potentiometer or the external circuit which completes the electrical connection between two electrodes also.
21:13In a galvanic cell, electrons flow spontaneously through the external circuit as the cell generates electrical energy from a spontaneous redox reaction.
21:25Conversely, in an electrolytic cell, the external circuit is connected to the external energy source that drives the non-spontaneous reactions by forcing electrons to move in a required direction.
21:39This external circuit is essential for allowing the redox reaction token by enabling continuous electron flow and is typically connected by wires and sometimes includes a load or measuring devices.
21:56Next, we will study about the salt bridge which is the basic component of an electrochemical cell.
22:05The salt bridge is typically a tube which is filled with an inert electrolyte solution such as potassium nitrate, KNO3 or sodium sulphate which is Na2SO4.
22:17Its main function is to complete the electrical circuit internally by allowing ions to flow between two half cells which prevent the bolt up of charges that would otherwise stop the reaction.
22:34So, if there were no salt bridge, some of the ions, for example, ZN2 plus ions, they will accumulate here and the other charges, for example, K plus ion or NO3 minus charges, they will build up here.
22:49So, at the cathode, negative charges will build, and at the anode, positive charges will build.
22:56So, to keep the neutrality, the salt bridge provides a path for the charges to travel.
23:03The salt bridge also prevents the direct mixing of the oxidizing and reducing agents in the separate solution, which force electrons to travel through the external wire instead of directly through the solution.
23:21Okay.
23:22Acting like a semi-permeable membrane, it allows the migration of electrolyte ions to maintain charge balance while minimizing the direct crossover of redox ions.
23:38This ensures the redox reactions continue smoothly and the current keeps flowing in the cell.
23:44And now, we will look into a typical example of electrochemical cell, which is a Daniel cell.
23:58So, the Daniel cell is a typical galvanic cell that harnesses the spontaneous redox reaction between zinc metal and cupric ions to produce an electric current.
24:11Okay.
24:12The cell consists of a copper vessel, which is filled with a saturated copper sulfate solution, which acts as a depolarizer by providing CO2 plus ion for the reduction reactions.
24:26Additionally, a dilute sulfuric acid is also present, and it acts as an electrolyte, which facilitates the ion flow and maintains electrical neutrality in the system.
24:40Also, an amalgamated zinc rod is immersed in a zinc sulfate solution, where the zinc metal undergoes oxidation, releasing electrons to the external circuit.
24:52So, in the anode, oxidation happens, and at the cathode, reduction happens.
24:58The overall cell is also shown by the standard nitrogen here.
25:03Here, we can see that Zden and Zden SO4, they are in the one half cell, and Cu SO4 and copper, they are in the other half cell.
25:12They are separated by these two lines, which means that this is an electrochemical cell, and these two electrolytes are the separate systems.
25:27Moving further into the electrochemical cells, we will now study the steps involved in the Daniel cell.
25:34So, the Daniel cell's overall chemical reaction is conventionally written as Zden plus, Cu 2 plus.
25:43They give rise to Zden plus and Cu solid metal.
25:48Here, solid zinc is oxidized to zinc ions, while copper ion in solutions are reduced to solid copper.
25:59In this cell, electrons flow from the zinc electrode to the copper electrode through the external circuit.
26:06This electron flow generates an electric current that can be used to provide any work.
26:13Meanwhile, the metal ions travel between the two half cells via the salt bridge.
26:19So, electrons flow from here, while the ions flow from here through the salt bridge,
26:27which maintain the electrical neutrality.
26:30Zden 2 plus signs enter the zinc half cell reaction, while the Cu 2 plus signs in the copper half cell gain electrons
26:37and deposits it as a copper metal with the electrode.
26:45The current, defined as the flow of positive charges, flows from copper electrode to the zinc electrode,
26:52which means it flows from cathode to anode externally.
26:56The Daniel cell is considered a reversible cell, meaning under certain conditions,
27:05the reaction can be reversed by applying external voltage.
27:09However, voltage cells in general can be reversible or irreversible,
27:15depending on their chemical makeup and the construction.
27:19Next, we will look into the chemical reactions.
27:29The electrochemical processes in the Daniel cell occur at the two electrodes through separate redox half reactions.
27:37At a node, zinc undergoes oxidation, losing electrons and entering the solution as Zden 2 plus sign.
27:45This is represented by the half cell reaction, which is shown here.
27:49The anode is thus the site of electron release and is the source of electrons for the external circuit.
27:57At the cathode, copper ions from the CuSO4 solution gain electrons and are reduced to solid copper.
28:05The reduction half cell reaction is given here.
28:09It is important to note that the positively charged copper ions are attracted to the cathode,
28:17which is negatively charged in terms of electron accumulation, not electrostatic charges.
28:22This movement results from the overall reduction in the chemical potential energy during the redox reaction,
28:31but not the electrostatic attraction alone.
28:34So, copper ions move to the copper electrode not just because of electrostatic attraction alone,
28:43but also due to the reduction in the potential energy.
28:46The overall reaction can be written as Zden plus Cu2 plus.
28:52They give rise to Zden 2 plus and copper solid metal.
28:57So, now we will study about some of the results of the Daniel cell or the process that happens in the cell.
29:10As the Daniel cell operates, zinc atoms at the anode lose electrons and dissolve into the solution as Zden 2 plus ions,
29:20which leads to the gradual corrosion of the zinc electrode.
29:25Simultaneously, copper ions in the copper sulfate solution gain electrons at the cathode
29:33and are deposited as solid copper, resulting in the accumulation of copper material at the positive electrode.
29:40The overall chemical transformation not only produces a visible change in the electrodes,
29:49but also releases electrical energy.
29:51The Daniel cell yields about approximately 213 kilojoule per mole of zinc ionized.
30:02So, this one energy is relieved with one mole of ionization or oxidation of the zinc.
30:09This significant energy output is attributed to the weaker metallic bonding in the zinc compared to copper.
30:21Because zinc atoms require less energy to release electrons and enter the solution,
30:27the reaction proceeds spontaneously and efficiently in terms of energy conservation.
30:33So, as we have studied in detail about electrolytic and electrochemical cells,
30:41we will now try to compare both of these cells and will look into their comparison.
30:47First of all, we will define both of the cells.
30:51Electrochemical cell, it is a cell which produces electricity as a result of chemical reaction
30:57and it will be called an electrochemical cell.
31:00While the electrolytic cell is an electrochemical cell in which electrical energy is converted into
31:07chemical energy.
31:08Okay.
31:09So, this is the basic difference.
31:11It uses, this cell uses electrical energy to produce chemical energy.
31:18This cell, which is an electrochemical cell, it produces electricity with the help of a chemical reaction.
31:25Moving towards the anode.
31:27Anode of the chemical, electrochemical cell is the negative electrode.
31:31While the anode of the electrolytic cell is the positive electrode.
31:35The cathode of the electrochemical cell is positive, while the cathode of the electrolytic cell is the negative.
31:44Next, the direction of the current.
31:48In an electrochemical cell, the current flows from cell to the external circuit, while in the electrolytic cell, the current flows from external circuit to the cell.
31:59Moving toward the chemical reaction, in an electrochemical cell, a spontaneous chemical reaction occurs to produce electricity, while in the electrolytic cell, a non-spontaneous chemical reaction takes place.
32:11Studying the construction of the cell, the construction of an electrochemical cell is in the form of two half-cells, while in the electrochemical cell, the complete cell is assembled in a single cell container.
32:27The energy conservation.
32:29In electrochemical cell, chemical energy is converted into the electrical energy, while in the electrolytic cell, electrical energy is converted into the chemical energy.
32:41Also, electrochemical cells do not require external source to supply power, while the external power supply is required in an electrolytic cell.
32:55In electrochemical cells, the electrons flow from anode to cathode through the external circuit, while in an electrolytic cell, the external battery supplies the electrons that go in through cathode and come out through the anode.
33:15Electrochemical cells are used as a source of electricity, such as battery.
33:20Electrolytic cells are used in electrolysis of the compounds.
33:23Now that we have understood both the functions of the electrolytic and electrochemical cells, we will now move toward the applications of the cell.
33:37One of the major applications of the electrochemical cell is as a battery.
33:42Okay, so in here, we will study only about the applications of the electrolytic cell, which are more important applications.
33:51Okay, so now moving toward the applications of the electrolytic cells.
33:55First, one of the most important applications of the electrolytic cell is the electrolysis of water.
34:01Using an electrolytic cell, we can break down water into hydrogen and oxygen gases.
34:08This process is vital for producing clean hydrogen fuel.
34:12Second, electrolytic cells are used in the extraction of aluminium from booxide ore.
34:19This process would not be feasible without electrolytic technology.
34:23Thirdly, electrolytic cells play a key role in electroplating also.
34:29It is a method where a thin layer of metal, like chromium, silver, or gold, is deposited on another metal's surface.
34:37This helps enhance the corrosion resistance and improve appearance.
34:43Fourthly, they are also employed in electrorefining, which is the purification of non-ferrous metals, such as copper, nickel, and lead.
35:01The impure metal is refined by depositing pure metal onto the cathode.
35:06Lastly, processes like electrovining, where metals are recovered from ores or solutions using the electricity, they rely on electrolytic cells.
35:19It is vital used in mining and metal recovery industries.
35:24These work style applications make electrolytic cells not just an academic topic, but a key player in the industrial chemistry and sustainable technologies.
35:36So, that is the end of our lecture.
35:41Thank you very much.