Oxidizing Agents - Online Tutor, Practice Problems & Exam Prep

The permanganate ion (MnO₄^-) is a strong oxidizing agent used in redox titrations, often requiring standardization due to its instability and tendency to form manganese(IV) oxide precipitate. In acidic solutions, it reduces to manganese(II) ion, which is colorless, allowing it to serve as its own indicator with a light pink endpoint. Other oxidizing agents include cerium(IV) ion, dichromate ion, and triiodide ion, each with unique properties and applications in organic chemistry, particularly in the oxidation of alcohols to aldehydes and carboxylic acids.

Oxidizing Agent - Permanganate Ion

Here we're taking a look at the oxidizing agent of permanganate ion. We're going to say the most commonly used oxidizing agent titrants include our permanganate ion, which is MnO₄^-, cerium(IV) ion, dichromate ion, as well as our triiodide ion. Now, we are going to say here that an analyte, also sometimes called a titrant, that is a weak reducing agent requires one of these strong oxidizing agents during titrations. That's because it goes in terms of your chemical thermodynamics. By using a strong oxidizing agent, we're going to basically push the equilibrium towards the product side to ensure that our analyte or titrant is successfully oxidized by these titrants. Now, we're going to say that our permanganate ion itself is difficult to isolate because it can easily oxidize its aqueous solvent, water, to form manganese(IV) precipitate.

For those of you who remember the lab where you deal with permanganate ion, remember it has a dark purplish color. If you get it on your hands or skin, it's going to stain them, and it's incredibly difficult to get that color out. Also, remember that depending on the labs that you've taken, sometimes you have to prepare your permanganate solution. And remember that you have to keep it in a place that's dark, absent of any light because light itself could somehow cause some of this permanganate ion to be converted into manganese(IV) oxide. And that's the reason why, although permanganate ion is a strong and great oxidizing agent, the whole storage of it makes it incredibly difficult to use, and usually, we'll go with other types of oxidizing agents instead.

Now, we're going to say here in order to isolate the oxidizing agent of permanganate ion, we must catalyze it with either acids, bases, manganese oxide, or manganese(II) ion. These help to push the equilibrium towards the side that helps to generate the permanganate ion. Here below, we have a typical half-cell reduction reaction within an acidic environment for permanganate ion.

MnO 4 - 1 + 8 H⁺ + 5 e^- → Mn²⁺ + 4 H₂O

If we were to increase the amount of permanganate ion, we would thereby force the equilibrium to go in the reverse direction. By this process, we'd help to make more permanganate ion. If we were to expose it to light, light would cause a catalytic response which would help to make more manganese(IV) oxide precipitate. If we used a base, adding a base would help to reduce the amount of H⁺ ion. We need to replenish that, so again the reaction would shift in the reverse direction to replenish it and as a result create more permanganate.

On the other hand, if we were to increase the amount of acid, we would increase the amount of H⁺ ion and following Le Chatelier's principle, we have to move in the forward direction to get rid of the excess acid added. This would only help to diminish the amount of your permanganate ion. So just remember, these are the 4 main strong types of oxidizing agents. The first two, permanganate ion and cerium(IV) ion, are actually the strongest out of the 4. The other two are weaker but with them being weaker, that gives them an advantage because they become more stable.

So just remember when it comes to your permanganate ion, we're going to say it's incredibly difficult to isolate. It has to be handled with extreme care. Go to the next video and see the steps necessary to standardize and prepare the permanganate ion.

Oxidizing Agent - Permanganate Ion

So recall that the natural color of our permanganate ion is a dark purple violet color. Now, depending on the pH of the solution, it can influence the type of products that we form and the color transitions that will result. Here we're going to say in a highly basic solution, so they have pHs that are equal to or greater than 12. They have concentrations that tend to be 1 molar or greater. We're going to say here that our permanganate ion is reduced to the manganate ion. This manganate ion has a greenish color to it. So here, we have a transition in color. However, violet and green, it's hard to see exactly where the endpoint is reached in terms of going from a purplish violet color to that greenish color.

Next, we're going to say in highly acidic solutions, so these have pHs that are equal to or less than 1. We're going to say permanganate ion is reduced to the manganese(II) ion here. Manganese(II) ion is a colorless ion. We're going to say here in acidic environments, we can say that our permanganate ion can serve as its own indicator because here we go from a colorless solution and as we're titrating with permanganate ion, it would transition towards a light pinkish color once we reach our endpoint. Here in a highly acidic environment, we can use permanganate ion as our oxidizing agent as well as our own indicator.

Finally, we can say that in a neutral or basic solution, the permanganate ion is reduced to manganese(IV) oxide precipitate. It's reduced to this solid here. This solid here has a brownish color. Realize that our permanganate ion is one of the strongest oxidizing agents that we have. But in terms of standardizing and preparation as well as storage, it's incredibly difficult to use. So we usually rely on the other strong oxidizing agents as other ways of dealing with redox titrations. Now of course at some point, you will come into contact with using permanganate for some type of titration process if you haven't already. Realize that it has its usefulness but again, it's incredibly difficult to use. Therefore, we tend to stay away from it unless absolutely necessary.

Oxidizing Agent - Cerium (IV) Ion

So now we take a look at the cerium 4 ion as our next strong oxidizing agent. Like other strong oxidizing agents, it needs to be standardized in order to be stored or to take part in additional reactions. It is less commonly used because of the cost in preparation, storage, and utilization. Although the permanganate ion is hard to maintain because it tends to go to its precipitate form, it would be a better option than the cerium 4 ion. The process for the cerium 4 ion is just too time-consuming and costly. Remember, we also have the other two strong oxidizing agents that we can later talk about in terms of their utilization. We have the cerium 4 ion being reduced to produce our cerium 3 ion. The cerium 3 ion is more stable in terms of its oxidative state out of the two ions. This means that it drives the equilibrium towards the product side. That means that the cerium 4 ion greatly wants to accept an electron to become this form here as a product. That's what makes it such a good oxidizing agent. By putting cerium 4 ion next to an analyte, it will quickly try to remove an electron from it to become this more stable ion.

In terms of the standardization process for cerium, there are multiple ways to do it. Two common ways, one is to use it from a primary standard of cerium ammonium nitrate which is given by our formula here and it's dissolved within an acidic media. Here, we'd have 1 molar of sulfuric acid solution. In addition to this, we could have cerium 4 hydroxide being standardized against an iron 2 or oxalic acid solution and then using ferroin as the indicator to determine what the endpoint would be. Here, their equations can be given as these two bottom ones here. By using our cerium ammonium nitrate, we tend to use that particular compound to help us determine the amount of iron within a given unknown sample and that's why we have iron ion within our first equation. Here, iron 2 goes from being plus 2 to being plus 3. It has lost an electron. That electron has been taken by the cerium 4 ion. Again, it does this because it wants to form the more stable cerium 3 ion.

In terms of our cerium 4 hydroxide, we could use our oxalic acid here and we'd say that it produces cerium 3 ion, the more stable ion but now two moles of it. We'd produce two moles of carbon dioxide as well as two moles of aqueous hydrogen ion. In terms of oxidation reactions, a common reaction involves the reaction of ferrous ammonium sulfate hexahydrate, which is given by this compound, and cerium ammonium nitrate, which we saw above. We're going to say the net equation is the equation that we saw just above here. The other common type of reaction that deals with cerium 4 ions involves malonic acid. Here, our malonic acid will be oxidized in the process. Cerium is reduced to cerium 3 ion. Malonic acid is oxidized here to formic acid. During this oxidation process where malonic acid becomes formic acid, we also have the generation of carbon dioxide gas and 6 moles of hydrogen ions as products.

Now, remember, permanganate ion and cerium 4 ion represent the two stronger oxidizing agents out of the common four that are typically used. They both have their strengths but also their weaknesses. Cerium 4 ion is more stable if it can be created in terms of a solution. However, the process of making it is costly and time-consuming. Permanganate ion is not as expensive in terms of making the solution but again it's hard to maintain permanganate ion solution because it always wants to create some of that manganese oxide precipitate. You always have to handle it with great care. When we look at these two oxidizing agents, they both really are trying to remove an electron from our analyte species to become a more reduced, more stable form. So, we'll continue on with our discussion of other strong oxidizing agents and keep in mind that all of these here are just helping to oxidize a given analyte in order to make a more favorable ion in the process.

Oxidizing Agent - Dichromate Ion

So now, we take a look at our next oxidizing agent, which is our dichromate ion. It is normally found as a form of potassium dichromate, with the potassium ion being more of a spectator ion. Now, we're going to say here, that the dichromate ion is not as strong as an oxidizing agent like cerium (Ce⁴⁺) or permanganate. But its advantages are because it's not as strong, it's more stable. Therefore, it can be used as a primary standard. The disadvantage of it not being as strong as cerium (Ce⁴⁺) ion or our permanganate ion is that it wouldn't work with weaker reducing agents. Now, we're going to say that its reduction half-cell reaction within an acidic solution can be seen as follows:

CrO72- +14 H+ +6 e- → 2 Cr3+ +7 HO

In terms of our dichromate and our chromium (Cr³⁺), we can say that dichromate has an orange tinge to it, while chromium (Cr³⁺) is green. These colors are not different enough for us to really see the difference, so we can't use dichromate as its own indicator like we would with permanganate. Now, we're going to say here that Diphenylaminesulfonic acid or Diphenylbenzadinesulfonic acid are the preferred indicators to discover the endpoint. The overall equation is going to be monitored by our calomel and platinum electrodes, our 2 types of reference electrodes. Here, the Diphenylenesulfonic acid represents the reduced form. It is colorless. Then the oxidized form would actually appear as red-violet. We'd see that color change in terms of the indicator, and once that color change happens, we'd be able to spot the endpoint in terms of our redox titration.

Now, once placed into a basic solution, dichromate is converted to chromate ion. Here, the chromate ion can react with 4 moles of H₂O liquid to give us 3 electrons involved. This produces chromium (Cr³⁺) hydroxide plus 5 hydroxide ions. This process is in terms of the standardization of our dichromate ion. We're creating a form that's more stable and easier to store for later use. So, remember, although the dichromate ion is not as strong as cerium (Ce⁴⁺) ion or permanganate, it does have its advantages. By being weak, it can be more safely stored for longer periods of time, and it can be used as a primary standard.

Click on to the next video to see the different and common types of oxidation reactions that occur with our dichromate ion.

Oxidizing Agent - Dichromate Ion

So many of the oxidizing reactions that happen with our dichromate ion occur within the field of organic chemistry. Now, a lot of reactions dealing with this particular oxidizing agent have to do with alcohols being oxidized. Here, we're going to say that primary alcohols can be oxidized into aldehydes. So a primary alcohol we could have is this type of primary alcohol, ethanol. Now when we use our dichromate ion, it can be oxidized into an aldehyde. Remember, an aldehyde is basically a carbonyl group which is C double bond O, single bonded to a hydrogen. So that is what makes this an aldehyde. Here, we can say that this aldehyde would be ethanol or acetaldehyde. Now, under more stringent conditions, we can change this into a carboxylic acid. So we could increase the heat, increase the concentration of our dichromate ion. That would force this to continue its oxidation process and become a carboxylic acid. Here, this would be called ethanoic acid, or by its more common name, acetic acid. Now, secondary alcohols can be oxidized to ketones. So a secondary alcohol that we can have here is 2-propanol. We use this oxidizing agent of dichromate, and that changes it into a ketone. Remember the feature of a ketone? It is a carbonyl, so C double bond O, with a carbon on each side. So here, 2-propanol becomes 2-propanone or just propanone, or by its common name, acetone. Here, tertiary alcohols, like any tertiary alcohols, they cannot be oxidized by any form of oxidizing agent because it would require cutting a carbon-carbon bond, which energetically is highly unfavored because there's too much of a cost in energy needed to do it. So just realize here again, dichromate ions are not as strong as some of the other oxidizing agents such as cerium IV ion or permanganate ion, but it has a lot of use in terms of organic chemistry when we talk about the oxidation of alcohols as well as aldehydes under more stringent conditions.

Oxidizing Agent - Triiodide Ion

So the triiodide ion represents the weakest of the 4 strong oxidizing agents. Because it's the weakest, it only really works when we're using a stronger reducing agent. Now, we're going to say its half-cell reduction reaction can be seen as your iodide ion absorbing 2 electrons to be reduced into 2 iodide ions. So here we're going to say molecular iodine, so I2, is slightly soluble in an aqueous solution. Now we have to make it more soluble, so what we do is we combine the iodide of the molecular iodine with 1 iodide ion. They combine together to give me my triiodide ion here which is much more polar and therefore much more soluble within an aqueous solvent like water.

Now this creates the new reduction reaction as I3− + 2e− gives 3I−. That's my new half-reduction reaction.

In terms of standardization, we would say a solution of triiodide is normally going to be normalized by Na2S2O3, and here we're using starch as our indicator for the triiodine ion. Here with the presence of iodine starch would go black when you use it as the indicator to spot the triiodide ion.

Now we're going to say the newly created iodine solution is colorless, but with exposure to air, the iodide ion will undergo oxidation to the yellow triiodide ion. So basically, if we expose this reaction to air, it will move in the reverse direction to create more of this triiodide ion here, and that is yellow in color. That's a way of testing to see, do we have a pure sample of just your iodide solution or did it move in reverse to create more of the triiodide ion?

In terms of oxidation, the most common type of oxidation we cover deals with ascorbic acid otherwise known as vitamin C. We call it the biological antioxidant because its role is just as a reducing agent. Now we're going to say here that iodine quickly oxidizes ascorbic acid and it's going to generate dehydroascorbic acid. Here we can see our equation as this here. So as a result of this, we oxidize our ascorbic acid which has this formula because of molecular iodine, and we create the new form here of dehydroascorbic acid. If we look at the changes done, we can see that we've lost these hydrogens here. As a result of losing those 2 hydrogens, you've created 2 new carbonyl groups. Remember, within biological systems, oxidation can mean 2 things: oxidation can mean the loss of hydrogens, or it can mean the forming of bonds with phosphorus, oxygen, nitrogen, sulfur, or a halogen. So, the oxygens have lost hydrogens and these carbons here, they've helped to make new bonds with the oxygens. That's what represents an oxidation. So remember, your 4 different types of oxidizing agents vary in strength, stability, as well as creation. Based on the situation in the lab, different ones will be required. Here, remember, we have our cerium ion as well as our permanganate ion being the strongest, followed by our dichromate ion, and finally, we have our triiodide ion being the weakest of the 4 oxidizing agents.

Here’s what students ask on this topic:

What are the common oxidizing agents used in redox titrations?

Common oxidizing agents used in redox titrations include permanganate ion (MnO₄^-), cerium(IV) ion (Ce⁴⁺), dichromate ion (Cr₂O₇^2-), and triiodide ion (I₃^-). Permanganate ion is one of the strongest oxidizing agents but is difficult to maintain due to its tendency to form manganese(IV) oxide precipitate. Cerium(IV) ion is also strong but costly to prepare and store. Dichromate ion is less strong but more stable, making it suitable as a primary standard. Triiodide ion is the weakest and is used with strong reducing agents.

Why is permanganate ion difficult to use as an oxidizing agent?

Permanganate ion (MnO₄^-) is difficult to use as an oxidizing agent because it tends to form manganese(IV) oxide (MnO₂) precipitate, especially in aqueous solutions. This makes it challenging to maintain a pure permanganate solution. Additionally, permanganate ion is sensitive to light, which can catalyze its decomposition. Therefore, it must be stored in dark conditions and handled with care. Despite these challenges, its strong oxidizing properties make it useful in specific titrations where it can also serve as its own indicator due to its distinct color change.

How does the pH of a solution affect the products formed by permanganate ion?

The pH of a solution significantly affects the products formed by permanganate ion (MnO₄^-). In highly basic solutions (pH ≥ 12), permanganate ion is reduced to manganate ion (MnO₄^2-), which is green. In highly acidic solutions (pH ≤ 1), it is reduced to manganese(II) ion (Mn²⁺), which is colorless. In neutral or mildly basic solutions, permanganate ion is reduced to manganese(IV) oxide (MnO₂), a brown precipitate. These pH-dependent reactions are crucial for selecting the appropriate conditions for redox titrations involving permanganate ion.

What are the advantages and disadvantages of using dichromate ion as an oxidizing agent?

Dichromate ion (Cr₂O₇^2-) has several advantages and disadvantages as an oxidizing agent. Advantages include its stability, allowing it to be used as a primary standard, and its suitability for long-term storage. It is also useful in organic chemistry for oxidizing alcohols to aldehydes and carboxylic acids. However, its disadvantages include being less strong than other oxidizing agents like permanganate ion and cerium(IV) ion, making it ineffective with weaker reducing agents. Additionally, the color change from orange dichromate to green chromium(III) ion is not distinct enough to serve as its own indicator.

How is triiodide ion used in redox titrations?

Triiodide ion (I₃^-) is used in redox titrations primarily with strong reducing agents due to its relatively weak oxidizing strength. It is formed by combining molecular iodine (I₂) with iodide ion (I^-), making it more soluble in aqueous solutions. The standardization of triiodide solutions is typically done using sodium thiosulfate (Na₂S₂O₃) with starch as an indicator, which turns black in the presence of iodine. Triiodide ion is commonly used to oxidize ascorbic acid (vitamin C) to dehydroascorbic acid, making it useful in biological and chemical analyses.