- The periodic table, electron shells, and orbitals
- Shells, subshells, and orbitals
- Introduction to electron configurations
- The Aufbau principle
- Valence electrons
- Electron configurations of ions
- Electron configurations of the 3d transition metals
- Atomic structure and electron configuration
To find the electron configuration for an ion, first identify the configuration for the neutral atom. Then, add or remove electrons depending on the ion's charge. For example, to find the configuration for the lithium ion (Li⁺), start with neutral lithium (1s²2s¹). Then, since the lithium ion has one less electron, remove an electron from the 2s subshell to get 1s². Created by Sal Khan.
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- I don't really understand what he means by shells or the numbers and letters like F:1S2(7 votes)
- So electrons exists in clouds in a way around the nuclei of atoms. These regions where we find electrons are represented by the quantum numbers, of which there are four. These include the principal quantum number, the angular quantum number, the magnetic quantum number, and spin quantum number. Due to the Pauli exclusion principal no two electrons can have exactly the same quantum numbers.
The principal quantum number, often represented by n, is another term for the electron shell and are positive integer values (n=1,2,3,etc.) These electron shell numbers broadly correspond to the period, or row, an element is in. So elements in the first row of the periodic table have the first electron shell, or n=1, as their valence shell or outermost electron shell.
The angular quantum number tells you the subshell an electron is in within an electron shell and is represented by the letter 'l' often. The allowed values of l in an electron shell are determined by: l = n-1. So electrons in the first electron shell can only have a value of l=0 since n=1. Electrons in the second electron shell can have a value of l=0 AND 1 since n=2, and so on. The subshell tells us broadly the shape the orbitals which hold the electrons take on and these are represented by the letters s, p, d, and f. The values of l correspond to the subshell letters, so: l=0 is s, l=1 is p, l=2 is d, and l=3 is f.
The magnetic quantum number, represented by ml, tells you the specific orbital within the subshell an electron is in. The values of ml in a subshell range from +l to -l. So if l=1, then ml could be -1,0, or +1. This tells us why the s subshells only have 1 orbital while p subshells have 3, and the other subshells have the amount of orbitals that they do (and by extension how many electrons they can hold).
Finally there is spin quantum number, represented by ms, of which an electron in an orbital can only take on two values: +1/2 or -1/2, broadly referring to the electron spinning clockwise or counterclockwise if you want to think of it like that. This is why orbitals can only hold a maximum of 2 electrons because they can only take on one of two ms values.
Putting all these together you can specify exactly which electron you are referring to in an atom.
For the purposes of electron configuration though, we only care about the principal quantum number (the shell number), the angular quantum number (the subshell or the letter), and the number of electrons in those subshells which are basically a result of the magnetic quantum and spin quantum numbers. Shell numbers are written first, then the subshell as a letter, then the number of electrons in that subshell are written as a superscript to the letter.
For for example neutral hydrogen has the electron configuration of: 1s^(1). The first 1 is the shell number, the s is the subshell the electron exists in, the 1 superscript is the number of electrons in that subshell. The order which you fill subshells for more massive elements follows the aufbau principal.
Sorry for the long response, but that's the stuff you need to know to do electron configurations well. Hope that helps.(14 votes)
- Where did we learn about electron configuration before??(5 votes)
- since the postive charged Ca ion has the same configuration as Ar does that mean that the Ca ion is now considered Ar??(2 votes)
- piggybacking on what zelmen said, the thing that determines which element it is, is the number of protons not electrons. It can happen where an element loses a proton because it is unstable, that is what we call radioactive decay.(9 votes)
- does this mean you can't differentiate between atoms just by their electron configuration?(3 votes)
- If we were dealing with neutral atoms which never had electron transfers, then yes we would have unique electron configurations and you could identify an element solely by them. However because electrons are transferred frequently we can have two particles with the same electron configuration, but from different elements. This is a phenomenon called isoelectronicity. So this means we can have a neon atom, a fluoride anion, and a sodium cation and we wouldn't be able to discern any difference if we only inspected their electron configurations since they are all the same.
Hope that helps.(5 votes)
- I think adding a video on an intro to electron configuration and how the orbitals are filled before this video would help a lot.
I had to search up a separate video on these topics to understand what was going on.(4 votes)
- I agree, students need more material before they start doing electron configurations. It's odd because here in the high school chemistry section it's excluded (the extra material), but included in the AP chemistry section.(2 votes)
- Can fluoride exist by itself? I understand that it is an ion and that it must take an electron to be negatively charged, but can't it just steal that electron from like potassium and those elemental ions exist and be separated as ions?(1 vote)
- Yes. Fluoride as part of an ionic compound like sodium fluoride (NaF) is a solid at room temperature and is bound to the cation. But many ionic compounds are soluble in water and dissociate into their constituent ions. So when sodium fluoride is in water, the water molecules essentially separate the sodium cations and fluoride anions from each other and they exist as individual ions.
Hope that helps.(3 votes)
- Are there 18 electrons in the third shell or 8 and if there are 18 electrons in the third shell why do we only fill it up to 8(1 vote)
- So in the third shell there's s, p ,and d orbitals which, if fully filled, do add up to 18 electrons. But if we have an element in the third row (or period) like sulfur or phosphorus, we're only going to fill up the 3s and 3p orbitals in their electron configurations because that's where their valance electrons are. Valence electrons being the electrons that are important to chemical bonding. Once those 3s and 3p orbitals are filled, we'll have used 8 electrons. Only when we start doing the electron configurations of 4 row transition metals like iron and cobalt do we start back filling the 3d orbital. So third row elements like to have an octet to have an electron configuration similar to a noble gas like argon because it makes them more stable. So third row elements are similar to second row elements like carbon and oxygen who also want to have an octet, however third row elements aren't limited to only have 8 electrons because of that 3d orbital and can have greater than 8. It's just that they don't use that 3d orbital as frequently because once they've reached that noble gas electron configuration, they're stable and don't have an overpowering desire to exceed that octet. Hope this helps.(4 votes)
- I tried following his advice on trying some electron configurations for myself, but ended up randomly picking the exception to the Aufbau principle: Rh.
Why is the electron configuration for Rh: [Kr] 4d^8 5s^1 instead of [Kr] 4d^7 5s^2?(2 votes)
- Yep that element, rhodium, is a good example how general rules, like the order of filling electrons, are not always correct and why we have to keep in mind exceptions. Sometimes there is more nuance in chemistry than a simple rule can handle.
Most of the exceptions to normal electron configuration rules for the transition metals occur in rhodium’s period, the fifth period. Elements like molybdenum and ruthenium also have different electron configurations than what we would predict using the usual rules. The difference is usually that one of the electrons in the s subshell is demoted to the lower d subshell. Except for palladium which demotes both of its s electrons into the d subshell. Whatever the change is though between the rule and what is observed, the goal is still the same, to achieve a lower energy state. Essentially the rhodium atom gains more stability by having that s electron in the d subshell than allowing it to remain in the subshell.
Hope that helps.(2 votes)
- I have a question.
Let's say you had to find the electron configuration of Neon
Well, that's pretty easy: [He]2s^2 2p6
However, what if we had to find the electron configuration of Neon with an extra electron? (Negatively charged ion).
Where would that negatively charged ion go? it can't fit into the p sub shell since the p sub shell can only hold 6 electrons.
Would it become [He] 2s^2 2p^6 3^s1 ?(1 vote)
- Yup, you are correct. To approach it from another angle, you could think of adding an electron as changing the electron configuration to the next element on the periodic table. So adding an electron to Neon would net the same electron configuration as Sodium (Na).(3 votes)
- Should you always use noble gas configuration? He switches back and forth, not always using it. It seems like it would make a difference in figuring out which are the valence electrons.(2 votes)
- Usually electron configurations are expressed using noble gas notation simply because it’s the same information, but condensed to its most essential information. But you should be able to convert between the long form and the noble gas form with ease still. Only then do you really know your electron configurations.
Hope that helps.(1 vote)
- [Instructor] In many videos we have already talked about electron configuration and now in this video we're going to extend that understanding by thinking about the electron configuration of ions. These are going to be charged atoms. Let's just start looking at some examples. Let's say we are dealing with fluorine. Now, we know what a neutral fluorine atom's electron configuration would be. In fact, if you want a little bit of practice, try to pause this video and think about what is the electron configuration of a neutral fluorine atom? All right, now let's work through this together. A neutral fluorine atom has nine electrons, and we could just use our Periodic Table of Elements. So first, we're going to have two electrons in 1s. So we'll have 1s two. And then we're going to go to the second shell. So then we go to 2s two. So far we have filled in four electrons. And next we got to the 2p sub-shell. And we are going to have, we're talking about a neutral fluorine, we are going to have one two three four five electrons in that 2p sub-shell. So it's 2p five. So if that's the electron configuration for fluorine, what do you think the electron configuration for fluoride would be? This is just the anion that has one extra electron. It is a negatively charged ion. Pause this video and try to figure it out. Well, here you're now going to have one extra electron. The fluorine has nabbed an electron from someplace and so where will that extra electron go? Well our 2p sub-shell has space for one more electron. So that's where it will go. So the fluoride anion is going to have an electron configuration of 1s two, 2s two, 2p, now it's going to have an extra electron here, 2p six. 2p six. Now let's do another example. Let's say we wanted to figure out the electron configuration of a part positively charged calcium ion. So calcium, let's make it two plus. It has a positive charge of two. You could do this as a neutral calcium that has lost two electrons. What would be its electron configuration? Pause this video and try to figure that out. All right, well one way to figure this out is first we could figure out the electron configuration of a neutral calcium atom and then from that, we can take two of the highest energy electrons away. And so neutral calcium, you could view it, actually let's do it in noble gas configuration. Neutral calcium, the noble gas that comes right before calcium is argon. So it's going to have the electron configuration of argon and then we are going to have two electrons for that fourth shell. It's going to fill in the 4s sub-shell. And so we're going to have argon and then we're going to have, let me do this in a new color, let's call this 4s two. Now what do you think is going to happen if we were to lose two electrons? Well those two electrons in that 4s sub-shell, in the fourth shell, are gonna go away. And so the electron configuration here for calcium with a positive two charge, this calcium cation, is going to be the electron configuration of argon and no 4s two. So it's actually going to have the exact same electron configuration as argon. So I will leave you there, just a couple of examples. And I encourage you, if you're in the mood, just pick any of these atoms, any of these elements, and think about what would happen if they gained or lost an electron and what their electron configurations might be.