Analyte Oxidation State: Videos & Practice Problems

In redox titrations, analytes often exist in multiple oxidation states, necessitating conversion to a single state for accurate analysis. Auxiliary reducing agents, like zinc and silver, facilitate this pre-reduction step, while oxidizing agents, such as peroxy disulfate and hydrogen peroxide, perform pre-oxidation. These agents ensure the analyte reaches a consistent oxidation state, allowing for effective titration. Understanding the roles of these agents is crucial for mastering oxidation-reduction reactions and achieving precise results in quantitative analysis.

Analyte Oxidation State

So within a given redox titration, we have the analyte and we have our titrant. We're going to say here that it is common for the analyte to exist in multiple oxidation states while in solution. So it possesses multiple charges. And we're going to say, when we do any type of quantitative analysis, it has to be converted to just one oxidative state. Now, we're going to say for example, iron itself can exist as 2 different forms.

It can exist as iron 2 or iron 3. Now, if we have a solution that contains both of these ions, that won't help us in terms of our redox titration. So what we would do is we could basically add either a reducing agent or an oxidizing agent to this solution of these 2 ions and in that way get rid of one of them. Once we've gotten down to 1 oxidation state for our iron, we can utilize that ion and either do an oxidation with it or a reduction. Now, a common type of oxidizing agent tends to be used with these ions is your Cerium 4 ion.

Cerium 4 ion is one of the strongest oxidizing agents that exist. As an oxidizing agent, it would basically remove an electron from iron. So if we're going to remove an electron from iron, that means iron is going to become more positive. And since these are the two forms that it can exist in, it'd have to react with Fe₂₊. Removing an electron from it means that iron now becomes Fe₃₊ This here, it oxidizes iron 2 by removing its electron so it accepts that electron.

So now it becomes Cerium₃₊. And that's how these oxidizing agents and reducing agents behave. They either help to oxidize one of the forms of a particular ion. So at the end, our solution possesses only one particular oxidation state for that given ion. From there, we then continue with our redox titration.

Now here if we're talking about reducing agents, we call them auxiliary reducing agents. We're going to say an auxiliary reducing agent represents an easily oxidized metal. Typically, we have zinc and silver but also less commonly used, tin and cadmium. Here, there are reducing agents. They're going to help reduce one of the ions for a given solution and it's part of what we call the pre-reduction step.

We're going to reduce one of these ions so that at the end, we have only one of the ions remaining that can then undergo some type of redox titration. Now, we're going to say that these auxiliary reducing agents, they're present as either solids, powders, or part of what we call a reduction column. Now, they reduce the analyte to 1 oxidation state. Once the analyte has attained a single oxidation state, the reducing agent is then removed and then we do our redox titration. Now, we're going to say that the 2 types of reduction columns that exist, in the first one, we have what's called the Jones reducer.

Here, this is just your reduction column. It's filled with a zinc amalgam. That's basically zinc bonded to mercury. We're going to say in the process, the zinc metal reduces the analyte while it is oxidized. Here are zinc amalgam.

It's going to help to reduce the analyte. In the process, we have our zinc here. That's ²⁺ and remember, if it's helping to reduce the other species, that means it itself is being oxidized. That's how it loses these 2 electrons. Here we have our mercury liquid that has basically decoupled itself from the zinc amalgam.

Our other common reduction column is called the Walden reducer. Here, this one is filled with basically little pebbles or granules of solid silver metal. We're going to say while the analyte solution is infused with acid. The solution itself is acidified. Here we have our silver solid.

It's going to react with the chloride ion. It gets oxidized and becomes plus 1 and that's how it's able to combine with the chloride ion that's -one. Here this represents the electron that it's lost. That electron that it's lost will then go to the analyte and reduce it so that we have only one oxidation state for that particular analyte. Remember, the whole point of an auxiliary reducing agent is to donate an electron or more to an analyte in order to have only one oxidation state at the end.

Once we get to that one ion, we can do a redox titration with that given analyte. Now that we've talked about auxiliary reducing agents, come back and see what we say in terms of auxiliary oxidizing agents.

Analyte Oxidation State

Alright, everyone. So now we're going to take a look at our auxiliary oxidizing agents. We're going to say the auxiliary oxidizing agents basically just oxidize the analyte and in the process of this, they themselves are reduced. Because they're oxidizing the analyte to just one oxidative state, we say that this is part of our preoxidation stage. So, we're going to create or keep the analyte that has the more positive charge, which therefore corresponds to the greater oxidation number.

Now, we're going to say here that peroxodisulfate, which is this, represents a powerful oxidizing agent that works in conjunction with the silver catalyst. So, here in this reaction, we have peroxodisulfate reacting with silver ion. Here, it's going to be reduced which is why we have the electrons here as reactants. In the process, we create our sulfate ion and then here this represents our new oxidizing agent that can then react with another species in order to oxidize it. Now, we're going to say here that this oxidizing mixture can oxidize for example manganese 2+ ion to permanganate ion.

It could oxidize cerium 3+ to cerium 4+, and that shows you the power of this oxidizing agent mixture because cerium represents one of the strongest oxidizing agents out there. But this auxiliary oxidizing agent is so strong that it's even able to oxidize cerium. From there, we also have our dichromate. Well, here we're going to have chromium being chromium III ion being oxidized to dichromate ion. We have here, this is our vanadium cation.

And then we have VO₂⁺. And this is just our vanadium 4 oxide cation. Now, beside peroxodisulfate, we also have hydrogen peroxide, which is H₂O₂. It represents another powerful oxidizing agent. This one, though, can work within an acidic or basic environment.

Here in this general equation, we have our peroxyacid, hydrogen peroxide actually hydrogen peroxide operating within an acidic environment. Again, it's being reduced so that it can oxidize something else. In the process of its reduction, it's going to become 2 moles of water as a liquid. Now in basic solutions, it oxidizes. Okay.

So here in basic solutions, it could oxidize cobalt 2 ion to cobalt 3 ion. It could oxidize iron 2 to iron 3. It could also oxidize manganese 2 ion to manganese 4 oxide. Now, if we switch it from a basic solution to an acidic solution, it now acts like an auxiliary reducing agent. So now it's going to reduce.

Here it could reduce dichromate ion into chromium 3 ion or it could reduce permanganate ion into manganese 2 ion. It's more versatile than peroxodisulfate because depending on the solution that it's in, it can act either as an auxiliary oxidizing agent and oxidize, or it can act as an auxiliary reducing agent and reduce. Now, other commonly used oxidizing agents also include silver oxide. In organic chemistry, we tend to talk about silver compounds that act as oxidizing agents like Tollens' reagent. For those of you who haven't taken organic, don't worry about it.

For those of you who have taken organic, you might have forgotten. Just a little quick bit of outside information. And we also have sodium bismuthate, which is NaBiO₃, as another special type of oxidizing agent. But out of these four, peroxydisulfate and hydrogen peroxide are the two most commonly referred to auxiliary oxidizing agents. Remember, when it comes to redox titrations, it's not uncommon to have your analyte exist in multiple oxidation states.

It's your responsibility to use either an auxiliary reducing agent or an auxiliary oxidizing agent to get that analyte down to just one single charge. Once we do that, we can then carry on with our redox titrations. Just keep in mind the jobs of your auxiliary oxidizing agents versus your auxiliary reducing agents.

Analyte Oxidation State Calculations

So continuing with this idea of oxidizing agents and reducing agents, let's answer each of the following questions based on the following half reactions. From the half reactions that are given, we can see that they're all written as reductions, where our electrons are written as reactants. Remember, in a reduction, your electrons are written as reactants. What we have here are all these cell potentials for each one of these half reactions. Now, for the first one, it's asking which is the strongest oxidizing agent.

So let's think about what an oxidizing agent is telling us. Remember, if you are the oxidizing agent, that means that you have been reduced. If you've been reduced, that means you represent the cathode. And remember, with the cathode, we have the gaining of an electron and if you're gaining an electron, that means you must be on the same side as the electrons. And if you are the cathode, you should have the largest cell potential present. Since all of them are written as reductions, we can compare them all.

If some were written as reductions and some as oxidations, you'd have to reverse all of the reactions so they are either all written as reductions or written as oxidations. It's more customary to have them all written as reductions where your electrons are reactants. Now we can see here that the largest cell potential is with the first equation where the cell potential is 1.36 volts. We know that our answers will rely on the first half-cell reaction and we're saying here you're gaining electrons here on the same side with the electrons. If we look at our two choices, we have chloride ion or chlorine gas.

It's the chlorine gas that's on the same side with the electrons, so that would be our answer. The strongest oxidizing agent presented here would be the chlorine gas. Now, what's asking is basically for the opposite. We're looking for the strongest reducing agent. So, if you're the reducing agent, that means you've been oxidized.

If you've been oxidized, you represent the anode. Now remember, the anode should have the smallest cell potential. Remember, if you're the anode, you're undergoing oxidation, that means you're losing electrons, which means you won't be on the same side with the electrons because you've lost them. So if we take a look, the smallest one is the bottom reaction or talking about vanadium and our choices are either vanadium solid or vanadium 2 plus ion.

Here again, we want to be away from the electrons. So the strongest reducing agent here would have to be the vanadium solid as my answer. So V solid would be our answer to this question. Now finally it says, will iodine reduce chlorine gas to chloride ion? Alright.

So let's think about what it's saying. So here, so it's saying that iodine is reducing chlorine gas. So that means that iodine would have to be oxidized. It's losing an electron and the electron that it's losing, it's giving over to chlorine. For it to be oxidized, it would have to have the smaller cell potential. So here, the equation with iodine does have a smaller cell potential than the equation with chlorine and if it's being oxidized, it would have to be on the side away from the electrons. Here, the iodide ion, yes. It is on the side that is not next to the electrons. So my iodide ion would reduce my chlorine gas and as a result become oxidized in itself.

So here we'd say, yes. So remember, for questions like this, you really have to go back and take a look and remember the principles that we talked about in the past. Leo the lion goes Grr. If you're losing electrons, you're being oxidized so you represent the reducing agent. Things that are oxidized are part of the anode which have the smaller cell potential.

Grr gain electrons reduction so you're the oxidizing agent. If you've been So you're the oxidizing agent. If you've been reduced, you represent the cathode. The cathode typically has the larger cell potential. Remember, this is true when we're dealing with a spontaneous electrochemical cell, when we're dealing with a galvanic or a voltaic electrochemical cell.

Those are spontaneous, so the cathode should have the higher cell potential and the anode should have the smaller cell potential. When we're talking about non-spontaneous electrochemical cells, like the electrolytic cell, everything is reversed. The cathode would actually be the smaller cell potential value and the anode would be the larger cell potential value. You'd be you'd be indicated which type of cell you're dealing with. Here, we're assuming that we're dealing with a spontaneous electrochemical cell, so we're dealing with a galvanic or a voltaic cell.

Here’s what students ask on this topic:

What is the role of auxiliary reducing agents in redox titrations?

Auxiliary reducing agents are used in redox titrations to convert the analyte to a single oxidation state. This pre-reduction step is crucial for accurate quantitative analysis. Common auxiliary reducing agents include zinc, silver, tin, and cadmium. These agents donate electrons to the analyte, reducing it to a consistent oxidation state. For example, zinc can reduce Fe³⁺ to Fe²⁺. Once the analyte is in a single oxidation state, the reducing agent is removed, and the redox titration can proceed effectively.

How do auxiliary oxidizing agents function in redox titrations?

Auxiliary oxidizing agents are used to oxidize the analyte to a single oxidation state in redox titrations. This pre-oxidation step ensures accurate analysis. Common auxiliary oxidizing agents include peroxy disulfate and hydrogen peroxide. These agents accept electrons from the analyte, oxidizing it to a consistent state. For instance, peroxy disulfate can oxidize Fe²⁺ to Fe³⁺. Once the analyte is in a single oxidation state, the oxidizing agent is removed, allowing the redox titration to proceed accurately.

Why is it important to convert the analyte to a single oxidation state in redox titrations?

Converting the analyte to a single oxidation state in redox titrations is essential for accurate quantitative analysis. Multiple oxidation states can lead to inconsistent and unreliable results. By using auxiliary reducing or oxidizing agents, the analyte is brought to a consistent state, ensuring that the titration measures the exact amount of the analyte. This consistency is crucial for precise and reproducible results in redox titrations.

What are some common auxiliary reducing agents used in redox titrations?

Common auxiliary reducing agents used in redox titrations include zinc, silver, tin, and cadmium. These agents are chosen for their ability to donate electrons and reduce the analyte to a single oxidation state. For example, zinc can reduce Fe³⁺ to Fe²⁺. These agents are often used in solid, powder, or column form, such as the Jones Reductor, which uses a zinc amalgam to achieve the reduction.

How does hydrogen peroxide function as both an oxidizing and reducing agent in redox titrations?

Hydrogen peroxide (H₂O₂) is versatile and can function as both an oxidizing and reducing agent in redox titrations, depending on the solution's pH. In acidic solutions, it acts as an oxidizing agent, converting Fe²⁺ to Fe³⁺. In basic solutions, it can also oxidize other species like Co²⁺ to Co³⁺. Conversely, in acidic solutions, it can act as a reducing agent, converting Cr₂O₇^2- to Cr³⁺. This dual functionality makes hydrogen peroxide a valuable reagent in redox titrations.