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00:00Hello everyone. Welcome to this exciting lecture series on electrochemistry.
00:07In this series, we will study about an important concept in electrochemistry that is the electrochemical series.
00:19This series is also known as the activity series.
00:22This concept plays a vital role in understanding how different elements and ions behave in redox reactions.
00:29This series helps us understand the relative tendencies of elements or ions to gain or lose electrons.
00:37And it serves as a foundational tool in electrochemistry.
00:41The electrochemical series is also called the activity series.
00:46It is a method of arranging the elements.
00:49This arrangement is not random.
00:51It is based on the electrode potentials of different elements and the ions.
00:56The series is structured by comparing their tendencies to lose electrons or gain electrons.
01:03The electrochemical series tells us how electro-positive or electro-negative the element or ion combination is,
01:10compared to the standard hydrogen electrode.
01:13The standard hydrogen electrode, or SHE, acts as a universal reference point assigned a potential of zero volts.
01:21Other elements are compared to it to understand whether they are stronger reducers or the oxidizers.
01:28This combination is also called a half cell.
01:31A half cell is essentially a single electrode immersed in a solution containing of its own ions.
01:37The behavior of this cell helps us understand and define where the element lies in the electrochemical series.
01:46A more electro-positive element loses electrons more easily than the hydrogen electron or the hydrogen in the standard hydrogen electrode.
01:56Elements above hydrogen in the electrochemical series are more likely to undergo oxidation because they readily lose electrons and act as strong reducing agents.
02:07On the other hand, a more electronegative element gains electrons more easily.
02:13Elements below hydrogen in this series prefer to gain electrons, making them strong oxidizing agents.
02:20This makes them ideal for the reduction half of the redox reactions.
02:25In general, more electronegative element takes electrons from more electro-positive ones.
02:32This explains why redox reactions occur.
02:36The element that wants more electrons, more electronegative elements, they pull them from one that is more willing to give them up,
02:46which are called the electro-positive elements.
02:49The electrochemical series thus can set to be a measure of electronegative nature.
02:55Ultimately, lower an element is in the series, the more electronegative it is.
03:02And the more powerful it is in the oxidizing agent, it provides a handy tool for predicting chemical behavior and spontaneity in the chemical reactions.
03:17So, next we will see a simple standard hydrogen electrode because it is used in electrochemical cells.
03:24And with the help of this reference electrode, we measure the potentials of the other elements.
03:31We see that here is a reservoir or a beaker or a container which has some electrolyte.
03:37This electrolyte has a molar concentration of 1 or it has a molarity of 1 mol per decimeter cube.
03:44Here we have a platinum electrode.
03:47This platinum electrode is connected via the platinum wire to the external circuit which provides electric current to this electrode.
03:55Here we have a hydrogen gas at a pressure of 100 kPa which is introduced into the cell.
04:02We see that there is a glass tube or here is a glass bell which covers our electrode and the wire.
04:11The hydrogen gas bubbles through, goes into this hole and it combined with the platinum electrode gets reduced or oxidized based on the other elements.
04:23And the hydrogen gas in the form of H2, it bubbles out from this.
04:28Okay.
04:29So, this was just the standard hydrogen electrode.
04:31Next, we will look into the standard electrochemical cell which is used to measure the electrode potentials of the other half cells.
04:45So, here we see that the hydrogen electrode that we have studied just now.
04:51It is placed in one container.
04:53So, this was the first half cell which is called the standard hydrogen electrode cell.
04:58The other half cell contains the metal of which we want to measure the electrode potential.
05:05So, between these two electrodes, there is a salt bridge which is used here to migrate ions so that the electrode neutrality of both of the cells is maintained.
05:18And the cell operation goes on continuously.
05:21We also have an external circuit which connects the hydrogen electrode to the other electrode.
05:27So, this is the basic structure and with the help of this cell, we can measure the electrode potential of the other cells.
05:34We see that the hydrogen electrode potential has an electrode potential of zero by definition.
05:40So, because this is a reference electrode, we have assigned it an arbitrary value of the zero volts and all other potentials are defined with respect to these hydrogen electrodes.
05:55In this slide, we will explore the electrochemical series which is also known as the activity series.
06:02It basically ranks elements and ions by their electrode potentials.
06:07This actually helps us understand the oxidizing and the reducing strength and it can predict chemical reactivity in a redox system.
06:16So, we see that the half cells which are the elements and ion pairs with very high negative electrode potentials are actually high in this electrochemical series.
06:26So, we see that the metal reducing activity is increasing from bottom to top.
06:33So, all of the elements which are above the hydrogen value which have the arbitrary value of the zero, they will be the strong reducing agents because they lose electrons more easily.
06:45So, the reducing activity increase from bottom to top above the hydrogen.
06:52Opposite to that, the half cells with positive electrode potentials are oxidizing agents.
06:57So, below zero or below hydrogen electrode, all of the elements or the ions will have strong oxidizing properties.
07:07So, they are strong oxidizing agents.
07:10So, we see that the more the oxidizing power, the more positive the value will be of any ion.
07:16So, the metal oxidizing activity will increase from top to bottom and the metal reducing property will increase from bottom to top.
07:29Now that we have understand what is the electrochemical series and how its values are obtained, let us move to the important part of the video which are the applications of electrochemical series.
07:41First of all, we will look at the reactivity of the metals.
07:46The reactivity of metal depends on the capacity or the ability of an atom to lose electrons to form a cation.
07:55So, if an atom lose electrons more easily to form cations, it will be a more reactive metal.
08:04This ability is measured by what is called the standard reduction potential or SRP for short.
08:12So, a metal that has a highly negative SRP or even slightly positive one, it will easily lose electrons that makes it more reactive chemically.
08:24We can observe a general trend in the activity series that if the metal reactivity decreases as you move down the group.
08:34So, metals at the top are far more reactive than those below.
08:39For example, alkali and alkali metals which are at the top and have a very negative standard reduction potentials are extremely reactive metals.
08:51They react with cold water releasing hydrogen, they dissolve in acids to form salts and react badly with substances that accept electrons.
09:03On the other hand, the metals like iron, lead and tin, nickel for example and also cobalt which are lower in the series, these metals don't react with cold water.
09:16But, they do react with steam to release hydrogen.
09:21Finally, metals like copper, silver and gold which lie below hydrogen in the reactivity series are much less reactive.
09:31They do not react with water at all, even in the steam.
09:36That is why they are also called noble metals.
09:39So, these metals do not react with water in any conditions.
09:43Moderate reactive metals which are a little higher in the series, they react with steam.
09:50And the metals which are at the top of the series, they react with even cold water.
09:54So, it means that our electrochemical series predicts the reactivity of metals based on their position in the series.
10:02Now, let us talk about the second application of electrochemical series.
10:09We will talk about the electropositive character which is another way of describing a metals ability to lose electrons and form positive ions.
10:20So, it also depends on the tendency to lose electrons.
10:25So, we see that just like reactivity, the electropositive nature of metals decreases as we move down the electrochemical series.
10:34This trend is again based on the standard reduction potential of each metal.
10:39We basically divide metals into three main groups based on their SRP values.
10:46First, we have the strongly electropositive metals which have SRP value around negative 2 volts or even more negative.
10:57They include alkali metals and alkalinous metals such as sodium, potassium and calcium.
11:03These metals easily lose electrons and are highly reactive.
11:07Next, there are moderately electropositive metals.
11:12With SRP values between 0 to minus 2 volts, metals like aluminum, zinc, iron, nickel and cobalt fall into this category.
11:21They also lose electrons but not as readily as the first group.
11:27And finally, we have the weakly electropositive metals which are actually below hydrogen in the electrochemical series.
11:34Their SRP values are positive meaning they don't lose electrons easily.
11:39Examples include copper, mercury and silver.
11:42So, in summary, the more negative the standard potential is, the more electropositive and reactive the metal will be.
11:50Next, we will also study about the application of displacement of hydrogen.
12:01This application explains how we can predict whether a metal can liberate hydrogen gas from either dilute acids or the water using the electrochemical series.
12:12First of all, let us talk about dilute acids.
12:17Metals that are above hydrogen in the electrochemical series have a negative reduction potential, meaning they tend to lose electrons easily.
12:26These electrons can be accepted by H plus science in the acid, resulting in the evolution of hydrogen gas.
12:33For example, we see that the oxidation half cell reaction is manganese loses some number of electrons to produce a positive ion.
12:45The reduction half cell is given as H plus science.
12:48They take these electrons and form a hydrogen gas.
12:52So, metals like zinc, magnesium and aluminum can displace hydrogen from HCl or H2SO4 which are citron acid or which can be dilute acid.
13:04However, metals below hydrogen in the series like copper, mercury, gold and platinum cannot displace hydrogen as they do not readily lose electrons as these metals.
13:17Now, regarding water, alkali metals like sodium and potassium and alkaline earth metals like calcium are so reactive they even displace hydrogen from cold water.
13:30Metals like magnesium, zinc and iron can do this only from hot water or steam due to their slightly lower reactivity.
13:42And we see that the metals below hydrogen in the series, for example copper, mercury, gold and platinum, they cannot liberate hydrogen even when they are reacted with steam or steam at very high temperatures.
13:59So, in short, this application helps us identify which metals can produce hydrogen gas and under what conditions simply by looking at their position in the electrochemical series.
14:11Moving toward the next application of electrochemical series, let us discuss the reducing power of metals and oxidizing nature of the non-metals.
14:24First of all, we will talk about the reducing power of metals.
14:27The reducing power of metals depends on how easily it loses electrons.
14:32This is related to the metal's reduction potential value.
14:37The more electronegative this value is, the stronger metal acts as a reducing agent.
14:45This value decreases from top to bottom in the electrochemical series.
14:50For example, we see that the sodium with the potential of minus 2.71 volts is much stronger than the zinc with the value of 0.7 minus 0.76 volts, which in turn much stronger than the iron, which have a less electronegative value than these two elements.
15:09This trend means that the alkali and alkaline metals are some of the strongest reducing agents because they have very high negative reduction potentials.
15:21On the other hand, the oxidizing power of non-metals depend on their ability to accept electrons or gain electrons.
15:30They have the value of the positive reduction potentials.
15:35As we move down the group in the electrochemical series, the oxidizing nature also increases.
15:42So it means that the more positive reduction potential is, the stronger it will be an oxidizing agent.
15:49Ok.
15:51So for example, fluorine is the strongest oxidant followed by the calorine, which is followed by the boramine,
15:59and the iodine.
16:01So we see that the fluorine has most electropositive value.
16:05Then we have a little less value, which is again a positive value for the calorine.
16:11Then the boramine have less value than the calorine,
16:15and iodine has less value than the boramine.
16:18So although their values are decreasing, ok,
16:22but all of these are the oxidizing nature.
16:29Next important application for the electrochemical series is the product of analysis.
16:34So we can actually predict the product of electrolysis happening in a cell.
16:41So we see that when multiple ions are present in an electrolyte,
16:45ions with a higher oxidizing power discharge first at the cathode.
16:50So we see that at the cathode, ions with that are stronger oxidizing agents,
16:56that is with a higher reduction potential values, they get discharge first.
17:00The order of cation deposition starts from potassium, ok,
17:04and calcium, which have the lowest reduction potential,
17:07and they go on increasing to gold and silver, which are the highest in values.
17:12At anode, stronger reducing agents, which have the lower reduction potential value,
17:18they discharge first.
17:20Here is an increasing order for the anion discharge.
17:24The sulfate ions have the lowest value,
17:27and the iodine ions have the highest value.
17:31So it means that iodine is discharged before boramine,
17:36and boramine will be discharged before colorine,
17:40and they in turn from hydroxylene, nitrate ions,
17:44and at the end, sulfate ions will be discharged at the anode.
17:50For example, we will consider the electrolysis example for the NaCl.
17:56At cathode, hydrogen ions are discharged to form hydrogen gas,
18:01since H plus ions have a higher reduction potential value than sodium ions.
18:05At anode, the calorine gas are discharged to form calorine gas,
18:14because the calorines have more reducing power,
18:18or the lower reduction potential than the hydroxylene.
18:22So, calorine gets discharged before the hydroxylene.
18:28Let us look into the second example for the copper sulfate solution.
18:32The copper ions discharge at the cathode and copper metal deposited.
18:37We see that at the anode, the hydroxyl ions discharge at the anode,
18:43releasing the hydro oxygen gas.
18:45This is all because of their values from their oxidizing power or their position in the electrochemical series.
18:55So these predictions help understand and control industrial and laboratory electrolysis processes.
19:01So, moving next into the applications of the electrochemical series,
19:10we will now look into the one of the most important applications of the electrochemical series,
19:17which is the calculation of the cell EMF.
19:20The electrochemical series gives us a scientific and systematic way to determine
19:25how much electrical energy a hydroxyl reaction can produce.
19:29Let us break this down by using the principles of half cell reactions and the standard electrode potentials.
19:38We see that each half cell undergoes a reaction, one oxidation and the other reduction.
19:45In an electrochemical cell, two half cell reactions occur as we have already known that.
19:50One loses electrons and gets oxidized and the other gains electrons and gets reduced.
19:57These reactions are associated with their oxidizing potential for the reduction and their reduction potential for the oxidation.
20:06So, the overall semi-cell EMF is the sum of oxidizing and reducing potentials of the cell.
20:13This total electromotive force or the EMF tells us how much potential difference exists between the two half cells.
20:21The cell works because of the difference in the tendency to gain or lose electrons.
20:27It measures the spontaneity of the overall reaction in the cell.
20:32A positive cell EMF means a redox reaction is spontaneous, meaning it will occur on its own without some external input.
20:41It is also a measure of the work that can be done by the cell.
20:45The EMF corresponds to the maximum electrical work that the electrochemical cell can perform under an ideal condition.
20:53The electrochemical series helps us measure the cell EMF by using standard electrode potentials of the two half cells.
21:03We can look at the E0 value from the electrochemical series and apply them to calculate the cell EMF.
21:10This is how the theoretical voltage of a cell is predicted.
21:14We can see the formula here.
21:17E0 of the cell is equal to E reduction cell minus the E oxidation value.
21:26This concept has been discussed in detail for the standard electrode potential video.
21:32So, you can also watch that video to understand it more.
21:35So, we see that here in this equation, E0RD is the standard reduction potential of the reduction half cell.
21:45And the E0OX is the standard reduction potential of the oxidation half cell.
21:52Even though it is oxidation, we use its reduction potential with subtraction.
21:57Because all values in a series are reduction potential by the convention.
22:01So, that is why we have used this negative sign here.
22:08Moving on to the next slide.
22:10This slide focuses on another major application of the electrochemical series,
22:16which is determining whether a redox reaction will occur on its own or not.
22:21In simple terms, the electrochemical series let us predict the spontaneity of a chemical change based on their cell EMF values.
22:30Let us walk through how this works.
22:33Measuring the spontaneity of a redox reaction is basically directly related to the cell EMF.
22:40The cell EMF, or the electromotive force, acts as an indicator of the energy difference between oxidizing and reducing agents in the redox system.
22:50The greater this value is, in the positive terms, the more thermodynamically favorable or spontaneous the reaction will be.
22:59If the cell EMF is positive, the reaction is spontaneous.
23:04Ok.
23:05A positive EMF means the electrons naturally flow from a node to the cathode and the redox process happens without external input.
23:16This confirms that the reactants are arranged properly to drive the reaction forward.
23:23Secondly, if the cell EMF is negative, a reaction is non-spontaneous.
23:29A negative EMF implies the reaction requires external energy to proceed.
23:35The reactants in this case are not naturally inclined to exchange electrons in the given direction.
23:41Thus, we can tell if a redox reaction can proceed spontaneously by examining the reactants and the products.
23:51By consulting with the electrochemical series, we can identify the standard electrode potentials of the reactants and the products.
23:59Comparing their ENUT values and applying the EMF formula allows us to predict feasibility without some lab trials.
24:11Next, in this slide, let us look at how we can estimate the Gibbs free energy of a redox reaction using cell EMF.
24:21This is another important application of an electrochemical series.
24:26Gibbs free energy is a key thermodynamic concept that tells us whether a reaction is thermodynamically favorable or not.
24:34In other words, whether it will happen on its own or not.
24:39Gibbs free energy is another measure of the spontaneity of a reaction.
24:48Just like the cell EMF tells about the feasibility, Gibbs free energy tells us how much useful work a system can perform.
24:56It is basically directly linked to the cell EMF.
25:00Meaning that as EMF changes, so does the spontaneity predicted by the Gibbs free energy.
25:09It is related to the cell EMF as follows.
25:12Delta G-nut cell will be equal to minus N F multiplied by E-nut of the cell.
25:19Here, in this equation, N is the number of electrons involved in the redox process.
25:27F stands for the Faraday's constant, which is approximately 96.485 coulombs per mole.
25:40So this is the value for the Faraday's constant.
25:43Again, based on the cell sign of the EMF, we have the following results.
25:50If the cell EMF is negative, Gibbs free energy is positive and the reaction is not spontaneous.
26:00A negative E-nut cell means the system cannot perform work naturally.
26:05The corresponding value of the positive delta G cell indicates that the energy must be supplied for the reaction to occur.
26:13As we know that the positive value of delta G means we have to add some energy into the cell or in any system.
26:20So basically this formula also predicts that if the cell value is negative, the delta G will be positive.
26:27And from the both values, we can predict that the reaction will be spontaneous or not.
26:33Next, we see that if the cell EMF is positive and we put it in here, the Gibbs free energy is negative and the reaction will be spontaneous.
26:44A positive E-nut cell results in a negative EG delta G-nut value, which confirms that the reaction is energetically favorable.
26:53This relationship actually gives us a clear thermodynamic picture of whether or not a redox process will proceed on its own.
27:08So, at the end, in this slide, we will see how we can use the electrochemical series to predict the end product of a redox reaction.
27:19Even when only the reactants are given, we can predict the end products.
27:24This is an extremely practical application of the electrochemical series and it helps us to understand what transformations are likely to happen in an electrochemical cell.
27:37So, if we are given only the reactants of a reaction, we can calculate the end products of the reaction as follows.
27:46We use the electrochemical series to write the standard electropotential values of each of the reactants.
27:55By consulting the series, we can assign each species its E-nut value, which is the standard electropotential value.
28:03This step is essential to rank the species in terms of their tendency to gain or lose electrons.
28:09Then, we note which have the highest and the lowest reduction potential.
28:15The species with the highest reduction potential is more eager to gain electrons, that is, it wants to be reduced.
28:22Conversely, the one with the lowest reduction potential is more likely to lose electrons, that is, to oxidize.
28:32Once we have these values, we can predict the end product as follows.
28:39The ions with the highest reduction potential are reduced at the cathode.
28:43The cathode is always the site of the reduction, as we have already known that.
28:49The ions with the lowest reduction potential are oxidized at the cathode.
28:57Anode is the site where oxidation happens.
29:00The species with the lowest or the most negative E-nut values will undergo oxidation here.
29:06The oxidized and the reduced ions gives us the end product of the reaction.
29:12By identifying which species are oxidized and reduced, we directly identify the final product of the redox reaction.
29:21This completes the redox reaction using nothing but the electrochemical series only and logical comparison of the E-nut values.
29:33So, that is all for today's video.
29:36I hope that you have understood and gained some knowledge from this video.
29:41Thank you very much.