- Oxidation and reduction
- Oxidation state trends in periodic table
- Practice determining oxidation states
- Unusual oxygen oxidation states
- Balancing redox equations
- Oxidizing and reducing agents
- Worked example: Balancing a simple redox equation
- Worked example: Balancing a redox equation in acidic solution
- Worked example: Balancing a redox equation in basic solution
- Redox titrations
- Oxidation–reduction (redox) reactions
Unusual oxygen oxidation states
Determining oxidation numbers in hydrogen peroxide, H₂O₂, and oxygen difluoride, OF₂. Created by Sal Khan.
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- So what about LiO2? What oxidation state does oxygen have, -.5?(15 votes)
- Yes. In LiO₂, Li has a +1 oxidation state. The two oxygens each have a -½.
You should be aware that all superoxides are radicals that are extremely reactive. Lithium superoxide (that is the name of LiO₂) is exceptionally extremely reactive. To the best of my knowledge, it has never been confirmed to exist at temperatures above 15 K (I did see a paper once about LiO₂ possibly existing momentarily at a temperature as high as 40 K, but I don't know whether that was ever confirmed.)(26 votes)
- In hydrogen peroxide the oxygen has one oxidation state right but the oxygen is bonded with another oxygen do that doesnt count as another oxidation state ?(8 votes)
- When an element is bonded to another atom of itself, that contribute 0 to its oxidation state. So, in other words, you just ignore the bond an atom makes to another atom of itself. So, in hydrogen peroxide we have:
H-O-O-H. The O-O bond doesn't count. So each O is bonded to one H. Since H has an oxidation state of +1, O must cancel that out, so its oxidation number is −1.(13 votes)
- How many oxidation states are there besides -2 and -1 ?(8 votes)
- There are a whole range of oxidation states that can range from about -7 to +7. It is rare that you go above 7. Chromium, a transition metal, with my knowledge, has the most oxidation states, which range I believe from +7 to -2.(5 votes)
- Sal says that with oxidation states we write the sign after the number (2+ or 2-), but according to my textbook oxidation states should be written as +2 and -2.
Which way is correct?(7 votes)
- I was always taught to use the +2 and -2 notation (i.e. number after the sign) in order to distinguish oxidation states from actual ionic charges. So when someone looks at your chemical equation they know it's an oxidation state you have written and not an ionic charge on the molecule.(6 votes)
- is it same for every compound that one element has to become electro+ while the other element has to become electro-(5 votes)
- No, that is not always the case. It is very common, but is not necessarily the case all of the time.
Remember, the oxidation states are ASSIGNED numbers, they are not necessarily actual charges that are physically present.(6 votes)
- what is the difference between a superoxide and peroxide and how to determine whether an oxide is superoxide or peroxide?(5 votes)
- At a very basic level, in a peroxide, the oxygen has an oxidation number of -1, where in a superoxide, the oxidation number of oxygen is -1/2. To determine whether the anion is a superoxide or peroxide, use the total oxidation number of a compound and the known oxidation numbers of other elements. For example, in KO2, we know that the oxidation number is 0, and K is +1; So we can do 1 + 2O = 0, simple algebra tells us that the oxidation number of the Oxygens is -1/2, which means it is the superoxide anion.(6 votes)
- what does covalent bond mean?(3 votes)
- A covalent bond is a bond formed between two non metal atoms where they share a bonding pair of electrons.
However you can have a dative covalent bond where one atom shares both of the bonding electrons(3 votes)
- at1:51why is hydrogen more electronegative(1 vote)
- You misheard Sal, he said hydrogen is less electronegative. Hydrogen is less electronegative than oxygen because oxygen has six electrons in its outer shell and so really really wants more electrons to become stable...wheras hydrogen doesn't want to gain electrons as much, so in the covalent bond the oxygen 'hogs the electrons' they spend more time near the oxygen and less time near the hydrogen.
Sal explain this really well in this video.https://www.khanacademy.org/science/chemistry/periodic-table-trends-bonding/v/other-periodic-table-trends(6 votes)
- At4:12, if Florine is more electronegative than oxygen, why not we consider it as standard? That is, why don't we say 'Floridised' instead of 'Oxidised' ?(3 votes)
- Our atmosphere is 20% oxygen and almost all metals are found naturally in compounds with oxygen (ie oxidised) on Earth.
The term was coined before we knew that oxygen was not the only element that could do this.(2 votes)
- so the oxidation number for an element is not absolute? i.e. O can be 2- or 1- or whatever the heck it wants?
if so, then how are we supposed to work out what its oxidation number is in any given question?(1 vote)
- Oxidation numbers can vary quite a lot for some atoms, but are very consistent for others.
Oxygen is actually among the most reliable atoms, as far as oxidation number. It nearly always counts for 2-, except when it's part of a peroxide, when it is instead 1-.
Fluorine has a 1-, with no exceptions that I know of. Other halogens also tend to have 1-, unless there is oxygen involved, as in polyatomic ions like chlorate and chlorite.
This makes sense if you think about electronegativity. Halogens are pretty electronegative, and tend to pull an electron toward them. But if you compare the actual EN values, oxygen is actually more electronegative than anything but fluorine, and tends to pull electrons away from even chlorine and the rest.
Fluorine is always 1-, because it pulls hardest, and only wants one electron, while oxygen pulls almost as hard, and wants two electrons, so it is almost always 2-.
It's also important to remember that in a pure element, the oxidation number is 0, by definition. So, in O2, both oxygen atoms have oxidation number 0, just as both fluorines in F2 have 0, and so on.(4 votes)
So we have two different molecules here. This is hydrogen peroxide. We call it peroxide, because it has this oxygen-oxygen bond. And here we have oxygen difluoride, where oxygen is bonded to two different fluorines. And what I want you to do is pause this video, use this periodic table of elements I have here, and this is more than just a typical information of periodic table of elements. It also gives you the electronegativities of these different elements. And these electronegativities are based on the Pauling scale named after famous biologist and chemist Linus Pauling. And so using the information here and what you know already about oxidation states, think about the oxidation states or the oxidation numbers for each of the constituent elements in these molecules. So pause the video now. So I'm assuming you have given a shot at it. And you might have immediately realized that something very interesting is going on. We've said in the past that because it's two valence electrons away from a full valence shell, because it is so electronegative, oxygen typically takes electrons from other things, typically two electrons, which typically gives it an oxidation state of negative-- an oxidation number or oxidation state of negative 2. This is so electronegative, and it so typically oxidizes other things that we've called the whole phenomenon "oxidation." But what's interesting here is that oxygen isn't purely bonded to things less electronegative than itself. And the hydrogen peroxide, yes, it is bonded to the hydrogen. But it's also bonded to another oxygen. And obviously, these two are going to be equally electronegative. So what would be the oxidation states or the oxidation numbers here? Well, hydrogen, once again, we portend-- hydrogen, because it's less electronegative, it would have a partially positive charge, because the electrons would spend more time around this oxygen. But when we're talking about oxidation states, we don't like this partial charge business. We want to pretend like these covalent bonds are ionic bonds, hypothetical ionic bonds. And if they were hypothetically ionic bonds, what would happen? Well, if you had to give these electrons to somebody, you would give them to the oxygen, the electrons in this period, give them to the oxygen, giving it an oxidation state of negative 1. With the hydrogen having these electrons taken away, it's going to have an oxidation state of positive 1. And the same thing's going to be true for that oxygen and that hydrogen right over there. So this is fascinating, because this is an example where oxygen has an oxidation state not of negative 2, but an oxidation state of negative 1. So this is already kind of interesting. Now it gets even more interesting when we go to oxygen difluoride. Why is this more interesting? Because fluorine is the one thing on this entire table that is more electronegative than oxygen. This is a covalent bond, but in our hypothetical ionic bond, if we had to give these electrons to one of these atoms, you would give it to the fluorine. So the fluorine, each of them would have an oxidation state of negative 1. And the oxygen here-- now, you could imagine, this is nuts for oxygen. The oxidation state for oxygen, it's giving up these electrons. It would be a positive 2. And we talk about oxidation states when we write this little superscript here. We write the sign after the number. And that's just the convention. But it has an oxidation state of positive 2. Oxygen, the thing that likes to oxidize other things, it itself has been oxidized by fluorine. So this is a pretty dramatic example of how something might stray from what's typical oxidation state or it's typical oxidation number. And in general, oxygen will have an oxidation state or oxidation number in most molecules of negative 2. But unless it's bonded with another oxygen or it's bonded to fluorine, which is a much more electronegative-- or actually, not much more, but it's the only atom that is more electronegative than-- or the only element is more electronegative than oxygen.