- 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
Oxidation state trends in periodic table
Trends in common oxidation states for main group elements. Created by Sal Khan.
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- Is there any difference between oxidation state and valency?? if yes what??(47 votes)
- valency always comes in whole numbers but oxidation number can come in fractions(2 votes)
- Is the oxidation state for groups 15, 16, and 17 are not always -3, -2, and -1 right ? and are groups 1 and 2 always +1 +2 ?(9 votes)
- yeah group 1 and 2 have stable oxidation state. And group 15,16,17 do not necessarily have same oxidation states of -3,-2,-1.
E.g. in H2O2 oxidation state of oxygen is -1 and in OF2 oxidation state of oxygen is +2(6 votes)
- what is the oxidation state for group 4 elements?
C, Si, Ge, Sn, Pb.?(5 votes)
- As tetravalent transition metals, all three elements (Ti, Zr, Hf) form various inorganic compounds, generally in the oxidation state of +4.(5 votes)
- At3:53, I always thought Hydrogen was part of the Alkali metals, so why does he say it can be a Halogen?(3 votes)
- Hydrogen is part of the alkali metals, but it can act like a halogen. This is because it is essentially just a proton. It can be "happy" either losing its sole electron and having empty electron shells, or it can be just as "happy" gaining an electron and having a full 1s shell. By "happy" I mostly mean stable.(10 votes)
- what about the transition elements?(3 votes)
- Transition elements are "weird" and you typically will not discuss them in much details until further studies. Nevertheless, you did ask so let me try to explain:
The elements in the first 3 rows (until 18Ar) have s and p orbitals, where higher shell number has higher energy and p has higher energy than s (for example, 2s > 1s, and 2p > 2s). When transition metals begin, they have an additional d-orbital but unlike in the first few rows, where it is clear where the electrons shells are in relation to others, the energy is quite close to that of nearby orbitals. Therefore it is not always clear, which of the electrons are removed to have a completely filled outer shell. For this reason, you will find that transition metal do not have one predefined oxidation state. Depending on the conditions, you have iron +2 or iron +3 (also noted as iron II or iron III). The same basic principles in the first few rows in the periodic table don't apply as easily as in the transition metals.(9 votes)
- 3:12Sal said "Group 5 elements" but meant "Group 15 elements."
Isn't group 5 elements correct?? Aren't there only 8 groups max?(2 votes)
- There are 18 groups (not counting the f-block). In an older system, the groups were labeled with a number and a suffix of either A or B. But there were inconsistencies in how these groups were labeled so a more straightforward system has been developed.
In the modern system, the groups are numbered 1 through 18 with no A or B suffixes. So, the modern Group 15 was previously called Group 5A or VA (or Group VB in Europe).(6 votes)
- At7:10, Sal says that O in -OH has an oxidation state of 2- and hydrogen has 1+. I'm having a hard time understanding why oxygen is 2- here. Could anyone clarify?(4 votes)
- Remember that oxygen is a diatomic molecule, meaning it takes two O atoms to make one molecule of oxygen. Hydroxide is an anion (negatively charged ion) made up of one oxygen atom to one hydrogen atom. Since oxygen is highly electronegative it is a common oxidizing agent. The oxidation number for O is 2- in most compounds. In this video, 2- is written above O. Another way to think about it is 1- being written above O2. Also check out the video about unusual oxygen oxidation states. https://www.khanacademy.org/science/chemistry/oxidation-reduction/redox-oxidation-reduction/v/unusual-oxygen-oxidation-states(1 vote)
- Whoa whoa whoa, Sal said in the previous video that the oxidation state of hydrogen in H2O is (-1). Is that supposed to mean that the charge of water molecule is equal to (-1) just like hydroxide?(2 votes)
- The oxidation state of hydrogen in water is +1 and oxygen is -2.
Hydrogen almost always has an oxidation state of +1. The main exception is that in hydrides of metals (such as NaH) hydrogen has a -1 oxidation state. And, of course, by definition, H₂ has hydrogen in the 0 oxidation state.
If Sal said that hydrogen had a -1 state in some compound other than a metal hydride, then he was mistaken. With that exception and the exception of elemental hydrogen, all hydrogen compounds have hydrogen in the +1 oxidation state.(5 votes)
- So is a hydroxide ion ionic or covalent ?
Can one find ions in nature ?
Can one create ions ?
How does oxygen just 'nab' that electron and 'ditch' the proton/hydrogen ion ?
Thanks in advance(2 votes)
- 1. Hydroxide is an ion. The H and O are covalently bonded to each other. Hydroxide ions can also form ionic bonds to cations, such as Na+.
2. Absolutely! Your body is full of ions, such as hydroxide, hydronium, and Na+!
3. Yes, you can create ions by mixing the correct chemicals together. For example, if you mixed HCl with H2O, then you would make the new ion H3O+.
4. If suppose you have two water (H2O) molecules. The proton/hydrogen ion from one water molecule can transfer to the other one. This gives you HO- and H3O+. It is thermodynamically favourable for a small percentage of water to react like this with itself, even in neutral conditions.(3 votes)
- 1:27Do metals ever form covalent bonds?(3 votes)
- What helps me is I have draw a border between the non-metals, metaloids, and metals and then write in big letters the types of bonds each type can have. Non-metal + Non-metal bonds are always covalent aka they always lend electrons to fill each other's needs to be stable. Since they are lending, it doesn't take much to seperate them again unlike Ionic bonds (strongest type of bonds)(1 vote)
Let's see if we can come up with some general rules of thumb or some general trends for oxidation states by looking at the periodic table. So first, let's just focus on the alkali metals. And I'll box them off. We'll think about hydrogen in a second. Well, I'm going to box-- I'm going to separate hydrogen because it's kind of a special case. But if we look at the alkali metals, the Group 1 elements right over here, we've already talked about the fact they're not too electronegative. They have that one valence electron. They wouldn't mind giving away that electron. And so for them, that oxidation state might not even be a hypothetical charge. These are very good candidates for actually forming ionic bonds. And so it's very typical that when these are in a molecule, when these form bonds, that these are the things that are being oxidized. They give away an electron. So they get to-- a typical oxidation state for them would be positive 1. If we go one group over right over here to the alkaline earth metals, two valence electrons, still not too electronegative. So they're likely to fully give or partially give away two electrons. So if you're forced to assign an ionic-- if you were to say, well, none of this partial business, just give it all away or take it, you would say, well, these would typically have an oxidation state of positive 2. In a hypothetical ionic bonding situation, they would be more likely to give the two electrons because they are not too electronegative, and it would take them a lot to complete their valence shell to get all the way to 8. Now, let's go to the other side of the periodic table to Group 7, the halogens. The halogens right over here, they're quite electronegative, sitting on the right-hand side of the periodic table. They're one electron away from being satisfied from a valence electron point of view. So these are typically reduced. They typically have an oxidation state of negative 1. And I keep saying typically, because these are not going to always be the case. There are other things that could happen. But this is a typical rule of thumb that they're likely to want to gain an electron. If we move over one group to the left, Group 6-- and that's where the famous oxygen sits-- we already said that oxidizing something is doing to something what oxygen would have done, that oxidation is taking electrons away from it. So these groups are typically oxidized. And oxygen is a very good oxidizing agent. Or another way of thinking about it is oxygen normally takes away electrons. These like to take away electrons, typically two electrons. And so their oxidation state is typically negative 2-- once again, just a rule of thumb-- or that their charge is reduced by two electrons. So these are typically reduced. These are typically oxidized. Now, we could keep going. If we were to go right over here to the Group 5 elements, typical oxidation state is negative 3. And so you see a general trend here. And that general trend-- and once again, it's not even a hard and fast rule of thumb, even for the extremes, but as you get closer and closer to the middle of the periodic table, you have more variation in what these typical oxidation states could be. Now, I mentioned that I put hydrogen aside. Because if you really think about it, hydrogen, yes, hydrogen only has one electron. And so you could say, well, maybe it wants to give away that electron to get to zero electrons. That could be a reasonable configuration for hydrogen. But you can also view hydrogen kind of like a halogen. So you could kind of view it kind of like an alkali metal. But in theory, it could have been put here on the periodic table as well. You could have put hydrogen here, because hydrogen, in order to complete its first shell, it just needs one electron. So in theory, hydrogen could have been put there. So hydrogen actually could typically could have a positive or a negative 1 oxidation state. And just to see an example of that, let's think about a situation where hydrogen is the oxidizing agent. And an example of that would be lithium hydride right over here. Now, in lithium hydride, you have a situation where hydrogen is more electronegative. A lithium is not too electronegative. It would happily give away an electron. And so in this situation, hydrogen is the one that's oxidizing the lithium. Lithium is reducing the hydrogen. Hydrogen is the one that is hogging the electron. So the oxidation state on the lithium here is a positive 1. And the oxidation state on the hydrogen here is a negative. So just, once again, I really want to make sure we get the notation. Lithium has been oxidized by the hydrogen. Hydrogen has been reduced by the lithium. Now, let's give an example where hydrogen plays the other role. Let's imagine hydroxide. So the hydroxide anion-- so you have a hydrogen and an oxygen. And so essentially, you could think of a water molecule that loses a hydrogen proton but keeps that hydrogen's electron. And this has a negative charge. This has a negative 1 charge. But what's going on right over here? And actually, let me just draw that, because it's fun to think about it. So this is a situation where oxygen typically has-- 1, 2, 3, 4, 5, 6 electrons. And when it's water, you have 2 hydrogens like that. And then you share. And then you have covalent bond right over there sharing that pair, covalent bond sharing that right over there. To get to hydroxide, the oxygen essentially nabs both of these electrons to become-- so you get-- that pair, that pair. Now you have-- let me do this in a new color. Now, you have this pair as well. And then you have that other covalent bond to the other hydrogen. And now this hydrogen is now just a hydrogen proton. This one now has a negative charge. So this is hydroxide. And so the whole thing has a negative charge. And oxygen, as we have already talked about, is more electronegative than the hydrogen. So it's hogging the electrons. So when you look at it right over here, you would say, well, look, hydrogen, if we had to, if we were forced to-- remember, oxidation states is just an intellectual tool which we'll find useful. If you had to pretend this wasn't a covalent bond, but an ionic bond, you'd say, OK, then maybe this hydrogen would fully lose an electron, so it would get an oxidation state of plus 1. It would be oxidized by the oxygen. And that the oxygen actually has fully gained one electron. And you could say, well, if we're forced to, we could say-- if we're forced to think about this is an ionic bond, we'll say it fully gains two electrons. So we'll have an oxidation state of negative 2. And once again, the notation, when you do the superscript notation for oxidation states and ionic charge, you write the sign after the number. And this is just the convention. And now, with these two examples, the whole point of it is to show that hydrogen could have a negative 1 or a positive 1 oxidation state. But there's also something interesting going on here. Notice, the oxidation states of the molecules here, they add up to the whole-- or the oxidation state of each of the atoms in a molecule, they add up to the entire charge of the molecule. So if you add a positive 1 plus negative 1, you get 0. And that makes sense because the entire molecule lithium hydride is neutral. It has no charge. Similarly, hydrogen, plus 1 oxidation state; oxygen, negative 2 oxidation number or oxidation state-- you add those two together, you have a negative 1 total charge for the hydroxide anion, which is exactly the charge that we have right over there.