- Shells, subshells, and orbitals
- Introduction to electron configurations
- Noble gas configuration
- Electron configurations for the first period
- Electron configurations for the second period
- Electron configurations for the third and fourth periods
- Electron configurations of the 3d transition metals
- Electron configurations
- Paramagnetism and diamagnetism
- The Aufbau principle
- Valence electrons
- Valence electrons and ionic compounds
- Valence electrons and ionic compounds
- Atomic structure and electron configuration
- Introduction to photoelectron spectroscopy
- Photoelectron spectroscopy
- Photoelectron spectroscopy
Valence electrons are the electrons in the outermost shell, or energy level, of an atom. For example, oxygen has six valence electrons, two in the 2s subshell and four in the 2p subshell. We can write the configuration of oxygen's valence electrons as 2s²2p⁴. Created by Sal Khan.
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- At4:26, Sal says "In most cases, your valence electrons are going to be your outermost electrons." What are some cases where valence electrons won't be the outermost electrons?(6 votes)
- Transition metals, lanthanides, and actinides have valence electrons which do not belong to the same electron shell.(9 votes)
- Mr. Khan says a full outershell is where the s and p orbitals are filled. What about the d, f etc orbitals for atoms with more shells?(4 votes)
- That isn’t strictly true for all elements. It is true for less massive elements in the first and second periods. This pattern begins to break down for elements in the third period (like sulfur and chlorine) who can still have an octet and achieve stability, but still have an unfilled d subshell.
The key is the filling of certain subshells, like the s and p, and not necessarily the entire shell that makes certain electron configurations stable.
Hope that helps.(7 votes)
- Why is it that Oxygen has 6 Valence electrons and not 8? Like why don't we count the first subshell s2 to add up for valence electrons?(4 votes)
- Sal mentioned it at1:33, but the 1s electrons are core electrons. Core electrons are different from valence electrons.(6 votes)
- I’m still a little confused, how do you know how many valence electrons the element has?(4 votes)
- Well, in a ground-state atom, it is usually equal to the column number on the Periodic Table (only for those that are not transition metals or lanthanides or actinides). You can also calculate which electrons are in which subshell (the estimation is not hard, KA has videos on it), and those in the shell furthest out are valence electrons. In a bonded atom, it depends on how it is bonded, and how you are counting them. Don't worry too much, as long as you get the idea, you'll learn about them as you progress through chemistry.(3 votes)
- Why did Sal skip the transition metals when calculating for valence electrons?
- To save you a headache. You should ignore transition metals for now, they don’t behave like the other elements(9 votes)
- What is the definition of valence electron for transition metal? I find many kinds of answers online:
1. just outermost shell (highest n, like main group)
2. (Wikipedia) electrons outside noble gas core (n s, n-1 d, n-2 f)
3. "accessible" electrons (if d/f orbitals are full don't include them)
For example, what is the number of valence electrons for Cu?
1&3 say one, but then they have to add "core electrons sometimes participate" (Cu3+); 2 says 11.(1 vote)
- The best definition of valance electrons in general are the electrons that participate in chemical reactions. For transition metals that means d orbitals and a higher s orbital. So for a transition metal in the fourth period like copper, Cu, this would mean a 4s and 3d orbital. You could count how many groups to the right copper is to find how many valence electrons it has. So 11. And then for copper ions you just subtract from that 11 number. So Cu3+ would have 8 valence electrons now.
Hope that helps.(9 votes)
- So, what does being stable mean here exactly?
If an atom is already neutral and has no net charge, then why it even needs to gain or lose electrons?
Why do they need to be stable?
Are they preventing reactions?(3 votes)
- So being stable when talking about valence electrons means that the valence shell has been filled completely (or half filled). The valence shell meaning the outermost electron shell. Stability meaning that something is unreactive, that it won't engage in some kind of chemical reaction to reach a new state. And vise versa, something which is unstable is reactive and will engage in chemical reactions to reach a new state.
So even if an atom is neutral, that is not necessarily its most stable state. For instance, if we take an atom of fluorine, F, in it's neutral state it has no net charge however it is VERY reactive. And this is explained if we look at fluorine's valence shell (the second shell) electron configuration: 2s^(2)2p^(5). The second electron shell holds a maximum of 8 electrons and 8 electron would be considered a filled valence shell and would therefore be stable and unreactive. But fluorine has 7 valence electrons, 1 away from being filled. And since it's so close to being filled, it will desperately want an extra electron which it will take from any nearby atom to do so. In other words, it will react in such a way to gain a new electron and fulfill its valence shell.
Now if we take a neutral atom of neon, Ne, which is just to the right of fluorine, neon is practically the opposite of fluorine in terms of reactivity despite both being neutral atoms. And again this can be explained by neon's valence electron configuration which is: 2s^(2)2p^(6). Neon has 8 valence electrons which is a filled second electron shell and is therefore stable and will be unreactive since it does not need to take electrons from other atoms.
This need to gain a filled valance electron shell by having 8 valence electrons is known as the octet rule and explains why certain elements are stable or unstable despite being electrically neutral. This octet rule holds for elements in the second and third periods (or rows) of the periodic table. Once you reach the fourth period and the transition metals they follow an 18 electron rule of stability, but it's the same idea as before in that they are attempting to fill their valence electron shells in order to become stable.
Hope that helps.(5 votes)
- How do I tell which 2 electrons would form an ion?(3 votes)
- I mean, do you mean from the electron configuration?
Even if we're counting electrons of an atom from the electron configuration, the it's the same idea for making an ion. If the protons and electrons of an atom are the same, then it is no an ion; it is neutral. If proton number differs from the electron number, then it is an ion.(3 votes)
- In the video, Sal emphasizes the idea of "core electrons", the electrons of the inner shells. Is knowing about the core electrons or how many there are important?(3 votes)
- Why does my textbook have, for instance, have the elctron config of phosphorus as 1s2 2s2 2p6 3s2 3px 1 3py1 3pz1
What are the x's y's and z's for?(1 vote)
- The p orbital have 3 sub-orbitals which are oriented in different directions according to their magnetic quantum number. If you imagine a 3D coordinate system with the nucleus at the origin, the p sub-orbitals would be shaped like two lobes (almost like a peanut) extending from the origin along one of the three axes. So you have a px orbital which lies on the x-axis, a py orbital on the y-axis, and a pz orbital on the z-axis. Phosphorus has 3 valance electrons in the 3p orbital and according to Hund's rule they must be placed into each sub-orbital singly before they are to be paired. So each p sub-orbital gets 1 electron in phosphorus therefore. So writing the electron configuration with 3p3 is the same as 3px1 3py1 3pz1, except the second notation is more detailed as to what's happening. Hope that helps.(5 votes)
- [Instructor] We are now going to talk about valence electrons, and non-valence electrons, which are known as core electrons and so one question that you might have been asking yourself this whole time that we've been looking at electron configurations is, what is the point? And the point of electron configurations is, is they can give us insights as to how a given atom or how a given element is likely to react with other atoms. And so just to make that point, or make it a little bit clearer, let's look at the electron configuration of an element that we'll see a lot of in chemistry, of oxygen. So oxygen's electron configuration is what? Pause this video and see if you can work through that. Well, in a neutral oxygen atom, you have eight protons and eight electrons, so first you're gonna fill the one shell, then you are going to start filling the second shell, so you're gonna go 2s2, so I have four right now, I have to have four more, so then you're going to have 2p4. And then notice, if I add up all the electrons here, I have exactly eight electrons. Now if I'm thinking about how might oxygen react, it's interesting to look at the outer oxygen electrons. The electrons that are in the outermost shell. So the outermost shell is being described right over here, this second shell. So how many electrons are in the outermost shell? You have six electrons here. So oxygen has six valence, valence electrons. And how many core electrons does it have? And the core electrons generally aren't reactive, or aren't involved as much in reactions? It has two core, two core electrons. Now, why is six valence electrons interesting? Well, atoms tend to be more stable when they have a filled outer shell, or in most examples, at least a filled SNP subshells in their outer shell. And so in this situation, you say, okay, oxygen has six valence electrons, and oftentimes that's drawn with a Lewis structure, and it might look something like this, where oxygen has one, two, three, four, five, six valence electrons, and you might say, hey, it would be nice if oxygen somehow were able to share, or get ahold of, two more electrons, because then that outermost shell will have a full number of eight electrons. The 2s and the 2p would be filled then, we would have 2p6. And so you'd say, alright, well maybe they can grab those electrons from something else and that's actually what oxygen does a lot of, it grabs electrons from other things. You can look at something like calcium. Pause this video, think about what the electron configuration of calcium is, and then think about how calcium is likely to react given that atoms tend to be more stable when they have a full outer shell, where both their S and P subshells are completely filled. Well, calcium's electron configuration, I could do it in noble gas notation or configuration, it'd have the electron configuration of argon, and one of the reasons why the noble gases are so stable is that they have a completely full shell. Argon for example has a completely full first shell, second shell, and third shell, and then to build calcium, will then have two electrons in that fourth shell, so it is argon and then 4s2. So how many valence electrons does calcium have? Well, you could see it right over there, it has two valence electrons. What about its core electrons? Well, a neutral calcium atom would have 20 electrons, 'cause it has 20 protons, so it would have 18 core electrons. Electrons that are less likely to react. And so you can say, what's the easiest way for calcium to get to a full outer shell? Well, instead of trying to gain six electrons, it might be a lot easier to just lose these two electrons. It is actually the case that many times, calcium will lose electrons, and become ionized, will get a positive charge. So the big picture here is, one of the values of electron configuration is to think about which of your electrons are most likely to react. Those are your valence electrons. In most cases, your valence electrons are going to be your outermost electrons. They're going to be the electrons in that outermost shell. Generally speaking, if you're talking about elements that are in the S block or the P block, you can think about how many valence electrons they have just based on what column they're in. This column right over here has one valence electron. This column over here has two valence electrons. This column out here has three valence electrons, four valence electrons, five valence electrons, six valence electrons, and seven valence electrons. The noble gases here, they are very unreactive, so one way to think about it is they are very very very stable, they have filled their outer shell.