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Drawing Lewis diagrams

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A Lewis diagram shows how the valence electrons are distributed around the atoms in a molecule. Shared pairs of electrons are drawn as lines between atoms, while lone pairs of electrons are drawn as dots next to atoms. When constructing a Lewis diagram, keep in mind the octet rule, which refers to the tendency of atoms to gain, lose, or share electrons until they are surrounded by eight valence electrons (an octet). Created by Sal Khan.

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  • blobby green style avatar for user Peter Patterson
    Why does every line in a Lewis diagram represent two electrons? Does that mean covalent bonds always share even numbers of electrons?
    (3 votes)
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    • leaf red style avatar for user Richard
      Yes, covalent bonds come in pairs which are represented by lines in Lewis structures. One line is a single bond with 2 bonding electrons, two lines is a double bond with 4 bonding electrons, and three lines is a triple bond with 6 bonding electrons. Covalent bonds form when two atoms react such that they share electrons in a bond between them and each atom donates half of the electrons which forms the bond from their original valence electrons. So yes each covalent bond will be a pair of electrons because each atom contributes 1 electron to a bond (And 1+1=2).

      Hope that helps.
      (11 votes)
  • blobby green style avatar for user Mitchell Estes
    I tried the Lewis structure on BeF2, and I came out with each F being double bonded to the Be,with extra electrons around each F, fulfilling the octet rule. But their answer had each F single bonded to the Be, with extra electrons around the F. Could you please explain this? Thanks.
    (2 votes)
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    • hopper cool style avatar for user Iron Programming
      When 2 atoms share electrons, what is actually happening is both atoms take one electron to share, making a pair of electrons shared. It isn't just that one atom shares one electron with another (that is not how it works).

      So we have Beryllium (Be) which has 2 valence electrons and then we take 2 Fluorine (F) which both have 7 valence electrons.

      If we make a single bond of F with anything else, then the Fluorine atom now has 8 total valence electrons, so a neutral fluorine atom will never make more than one covalent bond because otherwise it would have more than 8 valence electrons & overflow into the next shell.

      Thus once the Fluorine atoms each make a single bond with Beryllium then they no longer have a reason to share anymore so Beryllium must make due with what it has. If Beryllium isn't completely full then it may try to get more electrons making making additional bonds to make the overall molecule even larger.

      Hope this helps,
      - Convenient Colleague
      (4 votes)
  • blobby blue style avatar for user P
    Is there a trick to remember the valence electrons of various elements without taking help of the periodic table?
    (1 vote)
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    • leaf red style avatar for user Richard
      You could memorize the number of valence electrons for every element, but that's unproductive and unnecessary unless you're working with that element a lot. The periodic table is the best way to remember the valence electrons for the elements in my opinion since it partially organizes elements by their valence electrons. It's a good tool which doesn't make us burden our memories with minor facts.
      (3 votes)
  • blobby green style avatar for user algifarihaikal123
    why lone pairs of electron should be write as a pair? does it related to its real appearance??
    (1 vote)
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  • scuttlebug blue style avatar for user Hannah Davidson
    Is there a significant difference between a dot structure and Lewis diagram?
    (1 vote)
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    • leaf red style avatar for user Richard
      Not really, they're basically interchangeable. A dot structure is any representation of atoms/molecules using dots for electrons. And a Lewis diagram (or Lewis structure or Lewis dot structure) is a type of dot structure created by the chemist Gilbert N. Lewis which is most commonly used in chemistry nowadays. There's a slight difference, but they effectively mean the same thing.

      Hope that helps.
      (1 vote)
  • starky tree style avatar for user Tzzy7
    What is the chemistry behind the least electronegative atom being central?
    (1 vote)
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    • starky tree style avatar for user Tzzy7
      Atoms which have lower electronegativities hold onto their electrons less tightly and therefore are more prone to share their electrons. The central atom of a molecule needs to be sharing its electrons with multiple atoms which is easier to do so with a less electronegative atom which isn't as reluctant to share its electrons.

      (pasted from this same question elsewhere)
      (1 vote)
  • female robot grace style avatar for user inquisitivechild
    Is every element trying to reach 8 in its outer shell? does that number ever change
    (1 vote)
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    • leaf red style avatar for user Richard
      Not every element follows the octet rule and strives to reach eight valence electrons. A simplest exception to this rule is hydrogen and helium and the first period. They strive to reach two valence electrons and hence follow the duet rule. Another exception are the transition metals which follow an 18-elecron rule. Primarily the octet rule is followed by main block elements (groups 1-2 & 13-18) and even then there are plenty of exceptions. For example phosphorus usually would want to follow the octet rule, but in a chemical like phosphorus pentachloride it has ten valence electrons. Elements in period 2 is where the octet rule best applies. A lot of chemistry is learning simple rules and finding out about all the exceptions.

      Hope that helps.
      (1 vote)
  • blobby green style avatar for user mn103050
    Does the Lewis structure apply to both covalent and ionic bonds?
    (1 vote)
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    • leaf red style avatar for user Richard
      Lewis structures are mostly applied to covalent molecules, and while it is exceedingly uncommon you should do the same for ionic compounds.

      It's just for ionic compounds electrons aren't shared so you won't have things like single bonds between atoms. Instead ionic compounds stick together through electrostatic forces (different electrically charged ions) which we usually represent with brackets and the charge in the upper right corner. Additionally ionic compounds don't exist as individual molecules (as their formula unit suggests) but as a repeating pattern of these formula units in a lattice. So you can use Lewis structures for ionic compounds too, but the bonding is different enough from covalent compounds that it's simply not used to the same amount.

      Hope that helps.
      (1 vote)

Video transcript

- [Sal] In this video we're going to think about constructing Lewis diagrams, which you've probably seen before. They're nice ways of visualizing how the atoms in a molecule are bonded to each other and what other lone pairs of valence electrons various atoms might have. And so let's just start with an example, then we'll come up with some rules for trying to draw these Lewis diagrams. So the first example that we will look at is silicon tetrafluoride, and tetrafluoride is just a fancy way of saying four fluorines, so tetrafluoride. Now the first step is to say, "Well, what are the electrons that are of interest to us?" And if we're talking about the electrons that are likely to react, we're talking about the valence electrons, so V.E. for short, valence electrons. So first let's think about how many total valence electrons are involved in silicon tetrafluoride. Well, to think about that, we could think about how many valence electrons does silicon have, and then how many valence electrons does each of the fluorines have if they were just free atoms and neutral, and then multiply that times four, 'cause you have four fluorines. So let's get out our periodic table of elements, and then you can see here that silicon, its outer shell is the third shell, and in that third shell it has one, two, three, four valence electrons. So silicon here has four valence electrons, and then to that, we're going to add the valence electrons from the four fluorines. A free, neutral fluorine atom, its outer shell is the second shell, and in that outer shell, it has one, two, three, four, five, six, seven electrons. So each of these fluorines has seven valence electrons, but there are four of them. So one silicon tetrafluoride molecule is gonna have four plus 28 valence electrons. So this is going to be a total of 32. Now the next step is to think about how might these be configured? And as a general rule of thumb, we'd wanna put the least electronegative atom that is not hydrogen at the center. And we've talked about this before, but you can even see from the periodic table of elements, fluorine is actually the most electronegative element, and so we would at least try to put silicon at the center and make fluorine a terminal atom, something on the outside. So let's try to do that. So let's put silicon in the center, and then we have to put the four fluorines some place. Let's just put one fluorine there, one fluorine there, one fluorine there, and one fluorine there. Now the next step is, well let's just say for simplicity that we just have single bonds between the silicon and each of the fluorines. So let's do that. So one bond, a bond, a bond, a bond. Now each of these covalent bonds, each of these lines in our Lewis diagram, they represent two electrons. So for example, this one right over here that I'm doing in yellow, that represents two electrons that are shared by this fluorine and this silicon. This represents another two electrons that is shared between this fluorine and the silicon. This is another two electrons that's shared between this fluorine and this silicon. And this is another two electrons shared between that fluorine and the silicon. So, so far, how many electrons have we accounted for? Well, each of these represent two electrons, so two, four, six, eight electrons. So if we subtract eight from this, we are left with 24 electrons to account for, 24 valence electrons. So now, our general rule of thumb would be, try to put those on those terminal atoms with the goal of getting those terminal atoms to having eight valence electrons. In general we try to get the octet rule for any atom except for hydrogen. Hydrogen, you just need to get to two in that outer shell. But fluorine, you want to get it to eight. It already has two that it can share, so it needs six more, so let's add that. Two, four, six. Let's do that again for this fluorine. Two, four, six. Do it again for this fluorine. Two, four, six. And then last but not least, for this fluorine. Two, four, and six. Now how many more electrons are now accounted for? Well, six in this fluorine, six in this fluorine, six in this fluorine, six in this fluorine, so six times four, we've now accounted for 24 more electrons. We've now used up all of the valence electrons. Now that's good, because we wanted to account for all of the valence electrons. We wanna represent them somehow in this Lewis diagram. The next thing to check for is how satisfied the various atoms are relative to to the octet rule. We've already seen that the fluorines are feeling pretty good. They each have six electrons that are not in a bond, and then they're able to share two electrons that are in a bond, so each of them can kind of feel like they have eight outer electrons, eight valence electrons hanging out with them. And then the silicon is able to share in four bonds. Each of those bonds have two electrons, so the silicon is also feeling good about the octet rule. So I would feel very confident in this being the Lewis diagram, sometimes called the Lewis structure, for silicon tetrafluoride. So just to hit the point home on what we just did, I will give you these steps, but hopefully you find them pretty intuitive. That's why I didn't wanna show you from the beginning. But as you see, step one was, find the total number of valence electrons. We did that. That's the four from silicon and then the 28 from the fluorines. It says add an electron for every negative charge. Subtract an electron for every positive charge. We didn't have to do that in this example because it's a neutral molecule. Then it says decide the central atom, which should be the electronegative except for hydrogen. That's why we picked silicon, because fluorine is the most electronegative atom. And then we drew the bonds. We saw that the bonds accounted for eight electrons, and we subtracted those electrons from the total in step one, and that's just to keep track of the number of valence electrons that we are accounting for. And then we had 24 left over. And then the next step, it says assign the valence electrons to the terminal atoms. That's where we assigned these extra lone pair electrons to the various fluorines, giving them an extra six each so that they were all able to fulfill the octet rule. And then we subtracted that from the total, really just to account, to make sure that we're using all of our electrons. It says it right here: subtract the electrons from the total in step two. And then we saw that all of our electrons were accounted for. But then in step four, it says if necessary, assign any leftover electrons to the central atom. We didn't have to do that in this example. If the central atom has an octet or exceeds an octet, you are usually done. In this case, it had an octet, so we felt done. And it finally says, if a central atom does not have an octet, create multiple bonds. Once again, in this example we were able to stay pretty simple with just single bonds. But in future examples, we're going to see where we might have to do some of these more nuanced steps. So I will leave you there, and I'll see you in the next example.