Calculate Oxidation Numbers - Online Tutor, Practice Problems & Exam Prep

Understanding oxidation numbers is crucial for grasping oxidation-reduction (redox) reactions. An oxidation number indicates an element's ability to gain, lose, or share electrons in its elemental or compound form. In its standard state, an element's oxidation number is zero. Group 1A elements have a +1 oxidation number, Group 2A has +2, and Group 3A has +3, while Groups 5A, 6A, and 7A have -3, -2, and -1, respectively. Noble gases typically have no charge due to their stability. Diatomic molecules like H₂, N₂, O₂, F₂, Cl₂, Br₂, and I₂, as well as P₄ and S₈, are examples of elements in their natural forms with an oxidation number of zero.

When an element forms an ion, its oxidation number equals its charge. For instance, an aluminum ion (Al³⁺) has an oxidation number of +3. However, oxidation numbers can vary widely in compounds, necessitating a set of rules to calculate them. For example, hydrogen is +1 when bonded to nonmetals and -1 when bonded to metals. Oxygen is usually -2, except in peroxides (-1).

Calculate Oxidation Numbers

Now, the oxidation number is an important idea as we make our way towards oxidation and reduction reactions, otherwise known as redox reactions. Now first we're going to say that the oxidation number itself is an element's ability to gain, lose or share electrons when a loan or in a compound. And we're going to say, when it comes to oxidation numbers, it's important to talk about the natural state of an element.

Now for an atom in its natural state, also called its standard state, its oxidation number or oxidation state is equal to 0. Now remember we have our periodic table here. We know that with our periodic table there are charges that are unique to different groups, and we know that certain elements exist in certain forms in nature. Now when it comes to group 1A, we know that the charge is plus one. Group 2A is plus 2. 3A is +3 and 4A we tend to skip because some elements have varying positive charges. So we're just going to skip that.

We know that group 5A is -3, -, 2, -, 1. And remember, everything is trying to become a noble gas. They're perfect. So they tend not to have charges. If you don't remember this or you haven't seen my videos on it and you want to explore this a little bit further, make sure you take a look at my videos on periodic table charges. Now also remember we've talked about natural states of elements in the past as well.

So when it comes to elements of the periodic table, we have our diatomic molecules. Those are hydrogen, which is H2, Nitrogen, oxygen, fluorine, chlorine, bromine, and of course our iodine. Also remember that phosphorus tends to exist as P4 in its natural state when found in nature, and that sulfur is S8, and since selenium is right below sulfur, it is SE8. These are the natural forms of these different elements when they're found in nature.

And remember all the others? I'm not writing them in because in nature they exist as monoatomic atoms. This you can also remember. Refresh your memory by taking a look at my periodic table videos. We talk about the natural states of elements, but right now just realize that if you find an atom in its elemental state or natural state or standard state, whatever term you're most comfortable with, its oxidation number will be equal to 0.

Calculate Oxidation Numbers Example 1

Which of the following choices represents a natural state with an oxidation number equal to 0? All right, Remember when you're in your natural state as an element, your oxidation number oxidation state will equal 0.

If we take a look here, we have Na3, Cl2, He, and Mn4. Remember that when it comes to chlorine, chlorine exists as a diatomic molecule, so Cl2 is its natural state. So this would not have an oxidation number equal to 0.

Here Na3, three and Minnesota. These are two metals. Remember these metals exist as monoatomic elements in their natural states so this wouldn't work and this wouldn't work.

The answer here is C because helium here is shown in its natural state and as a result of this would have an oxidation number equal to 0. So here option C is the correct answer.

Calculate Oxidation Numbers

Now when an element gains a charge, it's no longer in its standard or natural state, so its oxidation will no longer be equal to 0. Now ions recall an ion is an element or compound with a positive or negative charge. Remember your positive ions are called cations and your negative ions are called your anions.

Now for a monoatomic ion, the oxidation number is equal to its charge. So if they gave us for example, the aluminum ion, aluminum is in Group 3A, so it charges 3+. If we see a charge present for that element, that is also its oxidation number. OK. But remember, it has to be in its ion form for the charge and oxidation number to equal one another.

OK, that's not always going to be the case. We're going to see later on there going to be cases where a particular element, its oxidation number, can range widely based on what it's connected to. But for right now, if you see an ion, you see its charge. That charge is equal to its oxidation number. So keep that in mind as we click on the next video and attempt an example question.

Calculate Oxidation Numbers Example 2

So here it says which of the following elements would have the most positive oxidation number based on its ionic form. Now if you've watched my videos on periodic table charges, you'll remember that the periodic table has some key things we need to remember when it comes to particular elements and their charges.

For example, silver. Silver is a transition metal. Most transition metals have multiple charges. Silver on the other hand does not. Remember, silver is always plus one for its oxidation number, scandium. Scandium is also a transition metal. It's in Group 3B, remember? So group 3B elements. They tend to have a charge of +3.

Next we have sodium. It's in Group 1A, so its charge is +1, and then sulfur is in Group 6A, so its charge is -2. We're looking for the most positive oxidation number, and that would have to mean that option B is the correct choice. It has the biggest positive value of +3.

Remember, if you haven't gotten a chance to look at my periodic table videos on charges, and you just want to go and take a look at the different types of elements and their charges, I highly suggest that you go back and take a look, because when it comes to charges of elements, it's going to form a foundation for a lot of later topics. We're going to cover in Chemistry one in Chemistry 2.

All right, so just remember in this question the answer would have to be option B. It has the most positive charge and therefore the most positive oxidation.

Which of the following elements would have the lowest oxidation number?

• A. Indium, In
• B. Strontium, Sr
• C. Argon, Ar
• D. Manganese, Mn

Calculate Oxidation Numbers

Like we said before, oxidation numbers don't always correspond to real charges and therefore a list of rules will be necessary. Now we're going to say when different elements are in a compound, these specific rules will be used to calculate our oxidation numbers. Now here let's take a look at these specific oxidation number rules.

When it comes to group 1A elements, they will be plus one when connected to any other element. So that's something similar to charges of groups and we're going to see some similarities between oxidation numbers and charges, but then we're going to see a huge deviation from that group, 2A plus two when they're connected to another element. Fluorine, it's -1 when it's connected to any other element.

Now this is where things start to change. Hydrogen can be either plus one or -1. It's plus one when it's connected to non metals. For example, if hydrogen is connected to chlorine or hydrogen is connected to oxygen or it's connected to the non metal of nitrogen. In all these instances, since hydrogen is connected to a nonmetal, its oxidation number will be plus one. Now it's -1 when it's connected to a metal or boron. So for example NaH or CaH₂ or BH₃. In these cases hydrogen will be -1 for its oxidation number.

Now oxygen is even more varied. We're going to say here that oxygen, when it's not a peroxide or superoxide, its oxidation number is -2. When it's in its peroxide form, its oxidation number will be -1. Now what exactly is a peroxide? Well, a peroxide we're going to say is when you have two group 1A elements connected to two oxygens, for example H₂O₂ (2 hydrogens which are in Group 1A, 2 oxygens). This would be hydrogen peroxide, Li₂O₂ (lithium peroxide) or K₂O₂ (potassium peroxide). Now if oxygen is a superoxide, its oxidation number is -1/2.

What is a superoxide? A superoxide is when you have one group 1 element with two oxygens. So here we could have potassium superoxide, cesium superoxide, sodium superoxides. So just remember oxygen can vary. So just look, be on the lookout, do you have a superoxide or peroxide? If not, its oxidation number is -2.

And then finally, when we're talking about Group 7A, we're talking about chlorine, bromine and iodine. They are going to be minus one, except when they're connected to oxygen. In that case, we won't know what their new oxidation number will be, and we'll have to calculate it. So again, we use these specific oxidation number rules when we're talking about different elements connected together. So we're going to have to utilize them to determine the oxidation number of any element given to us within a compound.

Calculate Oxidation Numbers Example 3

Here in this example it says which compound has oxygen with the lowest oxidation state. If we take a look at the first one, we have Na, O₂. We have one group 1A element with two oxygens. Remember that pattern tells us we have a superoxide. And if you have a superoxide, that means the oxidation number of oxygen is -12.

For the next one we have CO₂. Now it's not a superoxide. For it to be a superoxide, we would need one group 1A element with the two oxygens. Carbon is not in Group 1A, it's in Group 4A, so it's not a superoxide. Next, it's not a peroxide either. To be a peroxide, you need two group 1A elements with two oxygens. So it's not a peroxide, it's not a superoxide. Therefore oxidation number is -2.

For the next one we have C₂O₂. That's two Cs which are in Group 1A with two oxygens. That fits the definition of a peroxide. This is a cesium peroxide. So with oxidation number for oxygen would be -1.

And then finally we have O₂. This is the standard or natural state of oxygen. Remember, an element found in its standard or natural state has an oxidation number of 0. So out of all our choices, the one that has the lowest oxidation number, the one with the most negative value would be CO₂.

So just remember, if we have specific rules for that given element, utilize them to find the correct answer.

Calculate Oxidation Numbers

Now when asked to determine the oxidation number of a non listed element within a compound, we're going to follow the next four steps. Now remember we've gone over the specific oxidation number rules for certain elements. This applies to elements that are not found within that list.

So step one, you're going to treat the non listed element as X. You're gonna write, you're gonna use the list to write the known oxidation number of other elements within that compound. For example if you see oxygen is in that compound and it's not a peroxide or superoxide, we know its oxidation number would therefore be -2.

Step three, if an element has a subscript then remember to distribute it. And then step four, add up the oxidation numbers, create an equation and make it equal to the charge of the compound. So these are the rules that we're going to employ when looking for non listed elements within a compound.

And we'll see how we utilize these four steps to get to our oxidation number for that particular element. Now that we've gone over them, click on the next video and let's take a look at an example question.

Calculate Oxidation Numbers Example 4

Here it says give the oxidation number of the carbon atom in the acetate ion. So we're going to say we're looking for carbon. So carbon is going to be our X. Hydrogen when it's with nonmetals is +1. And then oxygen here in it's not a superoxide or a peroxide, so it's -2.

So we have two carbons, each one is X + 3 hydrogens. Each one is plus 1 + 2 oxygens. Each one is -2. This equation equals the charge of the ion which is -1. So all we have to do now is solve for X. So two X + 3 - 4 = -1, two X - 1 = -1. So add 1 to both sides here.

So here you know weird turn of events. You can see that X now equals 0. So here the oxidation number of carbon within the structure is 0. So there are moments where your oxidation number could be 0 and the element is not in its natural state, such as in this example. So again, the oxidation number of carbon when we calculated it comes out to be.

2 X + 3 - 4 = - 1 2 X - 1 = - 1 X = 0

Here’s what students ask on this topic:

How do you calculate oxidation number?

To calculate the oxidation number of an element in a compound, follow these rules:

1. The oxidation number of an atom in its elemental form is always 0. For example, $O_2$ or $H_2$ has an oxidation number of 0 for O and H, respectively.
2. For monoatomic ions, the oxidation number is equal to the charge of the ion. For example, ${\mathrm{Na}}^{+}$ has an oxidation number of +1, and ${\mathrm{Cl}}^{-}$ has an oxidation number of -1.
3. Oxygen usually has an oxidation number of -2 in compounds, except in peroxides like $H_2{O}_2$ where it is -1.
4. Hydrogen typically has an oxidation number of +1 when bonded to nonmetals and -1 when bonded to metals.
5. The sum of oxidation numbers in a neutral compound is 0. In polyatomic ions, the sum of oxidation numbers equals the ion's charge.
6. Fluorine always has an oxidation number of -1 in its compounds.

By applying these rules, you can assign oxidation numbers to each element in a compound and balance redox reactions accordingly. Remember that some elements can have multiple oxidation states, so the context of the compound is important for determining the correct oxidation number.

How do you calculate the oxidation number of an element in a compound?

To calculate the oxidation number of an element in a compound, follow these guidelines:

1. The oxidation number of a free element (not combined with other elements) is zero. For example, $O_2$, $H_2$, or $S_8$ all have oxidation numbers of zero.
2. For ions composed of only one atom, the oxidation number is equal to the charge on the ion. For example, ${\mathrm{Na}}^{+}$ has an oxidation number of +1, and ${\mathrm{Cl}}^{-}$ has an oxidation number of -1.
3. Oxygen almost always has an oxidation number of -2, except in peroxides (like $H_2{O}_2$) where it is -1, and in compounds with fluorine where it can be positive.
4. Hydrogen is usually +1 when bonded to non-metals and -1 when bonded to metals.
5. The algebraic sum of the oxidation numbers of all atoms in a neutral compound is zero. In polyatomic ions, the sum of the oxidation numbers must equal the charge of the ion.
6. Fluorine always has an oxidation number of -1.
7. For other halogens (Cl, Br, I), the oxidation number is usually -1, unless they are bonded to oxygen or fluorine.

By applying these rules, you can assign oxidation numbers to elements in compounds.

What is an oxidation number?

An oxidation number, also known as an oxidation state, is a number assigned to an element in a chemical compound that represents the number of electrons lost or gained by an atom of that element in the compound. It is a theoretical construct that helps in understanding the electron distribution in molecules and ions.

Oxidation numbers can be positive, negative, or zero. A positive oxidation number indicates that the element has lost electrons, becoming more positive, while a negative oxidation number means the element has gained electrons, becoming more negative. Elements in their elemental form have an oxidation number of zero.

The rules for assigning oxidation numbers include:

1. The oxidation number of a free element is always 0.
2. The oxidation number of a monatomic ion is equal to its charge.
3. Oxygen usually has an oxidation number of -2, except in peroxides where it is -1, and in compounds with fluorine, where it can be positive.
4. Hydrogen has an oxidation number of +1 when bonded to non-metals and -1 when bonded to metals.
5. The sum of oxidation numbers in a neutral compound is 0, while in a polyatomic ion, it is equal to the ion's charge.

Understanding oxidation numbers is crucial for balancing redox reactions and determining the electron transfer processes in chemical reactions.