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Course: Chemistry library > Unit 7
Lesson 4: Dot structures and molecular geometry- Drawing dot structures
- Drawing Lewis diagrams
- Worked example: Lewis diagram of formaldehyde (CH₂O)
- Worked example: Lewis diagram of the cyanide ion (CN⁻)
- Worked example: Lewis diagram of xenon difluoride (XeF₂)
- Exceptions to the octet rule
- Counting valence electrons
- Lewis diagrams
- Resonance
- Resonance and dot structures
- Formal charge
- Formal charge and dot structures
- Worked example: Using formal charges to evaluate nonequivalent resonance structures
- Resonance and formal charge
- VSEPR for 2 electron clouds
- VSEPR for 3 electron clouds
- More on the dot structure for sulfur dioxide
- VSEPR for 4 electron clouds
- VSEPR for 5 electron clouds (part 1)
- VSEPR for 5 electron clouds (part 2)
- VSEPR for 6 electron clouds
- Molecular polarity
- VSEPR
- 2015 AP Chemistry free response 2d and e
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Drawing Lewis diagrams
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.
Want to join the conversation?
Why does every line in a Lewis diagram represent two electrons? Does that mean covalent bonds always share even numbers of electrons?(3 votes)
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.(14 votes)
What is the chemistry behind the least electronegative atom being central?(4 votes)- 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)(11 votes)
- 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)
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)
- Does the Lewis structure apply to both covalent and ionic bonds?(1 vote)
- 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.(4 votes)
Is there a trick to remember the valence electrons of various elements without taking help of the periodic table?(1 vote)- 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)
Is every element trying to reach 8 in its outer shell? does that number ever change(1 vote)- 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.(3 votes)
Is the reason that the central atom is the least electronegative except Hydrogen because Hydrogen wants only 2 valence electrons so it won't form that many bonds? Or is it because Hydrogen is electronegative?(1 vote)- The central atom of a molecules is going to be one sharing its electrons with all the other atoms. An atom is more likely to be able to share electrons if it’s not holding onto its electrons as tightly. Since electronegativity is a measure of how strongly an atom attracts electrons to itself, low electronegativity values signify an atom more willing to share its electrons. So having the atom with the lowest electronegativity makes it more likely for multiple covalent bonds to form.
Hydrogen wants to gain noble gas electron configuration like all other elements since noble gases are more stable. The nearest noble gas is helium with two valence electrons, so hydrogen wants to also have two valence electrons. A single hydrogen atom has one valence electron, so it only needs to react with one other atom covalently with a single bond to get a second electron.
So hydrogen is never a true central atom of a molecule simply because its limited in its bonding abilities. You would never create a molecule any more complicated than a diatomic molecule with hydrogen as the central atom. You could argue that hydrogen is the central atom in diatomic molecules like H2 or HF. But again, we would just be limited to just these diatomic examples, and we know that there are molecules which are much larger than these which means hydrogen isn’t acting as the central atom in these larger molecules.
Hope that helps.(2 votes)
Why do some molecules need double bonds to be stable? is it because they are different kinds of atoms from the 2p subshell type, but 2 oxygens for example (which are 2p subshell) don't need double bonds because they re the same type of element.(1 vote)
the need for double bonds arises when the atoms involved require additional electrons to achieve stability and satisfy the octet rule. It is not limited to different types of atoms or specific subshells but rather depends on the specific electron configuration and bonding requirements of the atoms in the molecule(2 votes)
- why lone pairs of electron should be write as a pair? does it related to its real appearance??(1 vote)
- I've explained why electrons spin-pair in a previous question of yours. An electron's spin is related so its spin quantum number.(1 vote)
- In what order should I draw the dots that surround the letter in the Lewis structure? Is it clockwise or counterclockwise? Where do they start (top, right, or left)?(1 vote)
- I mean, it really is irrelevant how you place the dots in a Lewis structure since what really matters whether atoms have complete octets or not. If you place the same number of electrons around an atom clockwise or counterclockwise, you get the same structure.(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.