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Course: Chemistry library > Unit 9
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 dot structures
Guidelines for drawing Lewis dot structures. Created by Jay.
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- Why does the least electronegative atom belong in the the center? Doesn't the central atom receive electrons from other atoms, so wouldn't it make sense for the most electronegative one to go in the center?(20 votes)
- no because in order to satisfy the octet rule which says the second and the third period elements under the neither electropositive nor electronegative must not become the non-central atom in the compound. oxygen satisfies the octet rule as the most electronegative and carbon the least electronegative between the two and hence we see the electron dot structure as this.(4 votes)
- Does the central atom always form a double bond with some other atom to complete its octet, or can it also have a lone pair instead?(7 votes)
- There are few rules in chemistry that are "always" followed by all substances.
Sometimes central atoms have double bonds to complete the octet, sometimes lone pairs, and sometimes they just plain don't follow the octet rule (the octet rule is not a firm law of chemistry: though it is very often followed there are numerous exceptions).(17 votes)
- This is killing me right now, one of the MANY things i'm confused about..... wouldn't an oxygen bond be left with 4 electrons, not 6 because it's already sharing two? And how do you know when the structure is bent or not?(8 votes)
- You can tell a structure is bent if there are lone pairs left behind. Oxygen and Ammonia are bent because they have lone pairs.
An oxygen normally has 6 electrons in it's outer shell (aka valance electrons because they're on the outermost shell). It wants 8 electrons (octet rule). It gets 2 more by sharing it with other atoms, like 2 Hydrogens.(5 votes)
- At, the description of the molecule/cation shows that Xe has five co-valent bonds plus two additionally allocated electrons... my question is where do these 12 electrons "reside" 5s and 5p together give six quantum combinations. The 4d is already populated and was not counted when totaling up the number of valence electrons. Do the bonds invoke or consume some of the 6s and 5d orbitals? 13:40(4 votes)
- You may not yet have gotten to these topics. One video is "VSEPR for 6 Electron Clouds". Then you look at the videos on "Hybridization and Atomic Orbitals".
The Xe atom says to itself, "I need six orbitals to hold 12 electrons. I can use the 5s and the three 5p orbitals, but I will have to use two of my empty 5d orbitals. I will temporarily take out the electrons from the 5s and 5p orbitals. Then I will hybridize all six empty orbitals to form six new equivalent sp³d² orbitals."
"Now I have the 8 electrons from the 5s and 5p orbitals, plus 1 from each of the 5 F atoms, minus 1 that I lost to form a cation. That makes 12 electrons, so I have just enough sp³d² orbitals to hold all the electrons."(13 votes)
- Could you please explain how you decide--which rule or criteria you use--which of the two atoms is "least electronegative"?(2 votes)
- It is one of the trends of the periodic table. Electronegativity increases upwards and from left to right. So, fluorine would be the most electronegative element.(15 votes)
- How come the central atom can exceed an octet?(3 votes)
- For elements from phosphorus up, the d orbitals are close in energy to the s and p orbitals, so they can use one or two d orbitals as well to form bonds.(3 votes)
- At, how do you know that it can reach octet? Is there a proof or a way to determine? Does it have 8 valence electrons just because it is in the second period? Then what about other periods? 5:28(3 votes)
- In NO2+ how come N has 4 bonds? I thought N could only have a maximum no. Of 3bonds??(2 votes)
- Nitrogen will always obey the octet rule. Nitrogen will only have three bonds if there is a lone pair attached to the nitrogen as in NH3 (ammonia).
NH4+ (ammonium) is another very common example of 4 bonds to Nitrogen. If there are four bonds to the nitrogen, the nitrogen will always carry a + charge.(4 votes)
- At, how do we know to draw the hydrogens bent? 9:45(3 votes)
- So what would be the Lewis Dot Structure for the compound of N2O2? I understand the valence electron part of it, but I don't understand the actual drawing of the Lewis Structure.(2 votes)
- http://www.chemspider.com/Chemical-Structure.3874228.html This shows the bonds of the H2O2.
https://images.search.yahoo.com/images/view;_ylt=AwrB8pg3sB9XM0UA2CgunIlQ;_ylu=X3oDMTIycnRndjAyBHNlYwNzcgRzbGsDaW1nBG9pZAMyZGI3ZWI3MWY0ZDAzMmM1MmExN2ZlODQzYmI0MDc5YQRncG9zAzIEaXQDYmluZw--?.origin=&back=https%3A%2F%2Fimages.search.yahoo.com%2Fyhs%2Fsearch%3Fp%3DN2O2%26fr%3Dyhs-mozilla-002%26ri%3D7%26hsimp%3Dyhs-002%26hspart%3Dmozilla%26tab%3Dorganic%26ri%3D2&w=486&h=1020&imgurl=2.bp.blogspot.com%2F-pMqQAsq9P8M%2FT0Fgt6rIgaI%2FAAAAAAAAAKM%2FVGP9wQS0aA4%2Fs1600%2Fn2o2_final.png&rurl=http%3A%2F%2Fimgarcade.com%2F1%2Fn2o2-lewis-structure%2F&size=51.8KB&name=%3Cb%3EN2o2%3C%2Fb%3E+Lewis+Structure+Figure+1%3A+lewis+structures+for&p=N2O2&oid=2db7eb71f4d032c52a17fe843bb4079a&fr2=&fr=yhs-mozilla-002&tt=%3Cb%3EN2o2%3C%2Fb%3E+Lewis+Structure+Figure+1%3A+lewis+structures+for&b=0&ni=160&no=2&ts=&tab=organic&sigr=11ccfb4ap&sigb=13n6qctib&sigi=12novehts&sigt=11q5qb3do&sign=11q5qb3do&.crumb=MgZEsw/1Qpl&fr=yhs-mozilla-002&hsimp=yhs-002&hspart=mozilla This shows the Lewis Structure.
I noticed that the first link has double bonds between the N and the second link has triple bonds between the N.(2 votes)
Video transcript
Here's some of the guidelines
for drawing dot structures. So let's say we
wanted to draw the dot structure for this molecule,
so silicon tetrafluoride. The first thing we
would need to do is to find the total number
of valence electrons. And we would account for these
valence electrons in our dot structure. So to find the
valence electrons, we need to look at
it periodic table. So here I have a modified
version of the periodic table. And you see what
I've done is, I've kind of cut out
the d-block here, so we can focus
in on the elements that we're going to be
drawing in our dot structures. It also makes it easier
to see how the group numbers correspond to the
number of valence electrons. For example, if we look at
elements in the first group, like hydrogen, lithium,
or sodium, the first group all have one valence electron. So the group number corresponds
to how many valence electrons something has. So hydrogen has one
valence electron. If we think about
the periods, hydrogen is in the first period,
or the first energy level. So periods are
going horizontally across your periodic table. Hydrogen is in group one,
it has one valence electron. And if I go over
here to helium, this would be two valence
electrons for helium. And so in the
first energy level, you can fit a maximum
of two electrons. And so this is
going to important when we're drawing
our dot structures. Because when we're
drawing hydrogen, we're always going
to surround it by two electrons, or a
single covalent bond. When you take the second
period on the periodic table, so here we are on
the second period, lithium has one valence
electron, beryllium has two, boron has three, carbon has
four, nitrogen has five, oxygen has six, fluorine
seven, and neon eight. And so you have more orbitals
in a second energy level. And so because of that,
you can fit more electrons. So maximum of eight electrons
in the second energy level. And this is where the idea
of the octet rule comes in. So for elements like carbon,
nitrogen, oxygen, fluorine, understanding the
octet rule is going to help you when you're
drawing dot structures. Now it is possible for
some of the elements of the second period to
not have eight electrons. It's possible for them to have
less than eight electrons. So things like boron
will sometimes do that. But it is not
possible for elements to have more than
eight electrons. Always check your
dot structures, and make sure that if
you have an element in the second period, you do
not exceed eight electrons. Once you get to
the third period, you have even more orbitals
available to you now. So in the first energy level,
you have only one s orbital. In the second energy level
you have s and p orbitals, and in the third energy level
you have s, p, and d orbitals. So you can fit more
than eight electrons. And so therefore it's possible
to exceed the octet rule for elements in the
third period and beyond. And we will see a few examples
of that in this video, and some of the
ones to come here. So getting back to our
molecule, silicon tetrafluoride, if I wanted to find out how
many total valence electrons are in this molecule, I need
to find these elements on my periodic table. So I go over here
and I find silicon, and I see it's in group four. So therefore one atom of silicon
has four valence electrons. Fluorine is over here
in a group of seven, and so therefore
each atom of fluorine will have seven
valence electrons. And I have four of them. So 7 times 4 gives me 28 valence
electrons for my fluorine. The total number valence
electrons for my molecule will be 28 plus 4. So I have to account
for 32 valence electrons when I draw this dot structure. So let's go ahead and
move on to the next step. Let's go back up here and
look at our guidelines. So we figured out how
many valence electrons we need to account for
for our dot structure. We don't have any
kind of charges, so we don't need to worry about
the rest of step one here. We move on to step
two, where we decide on the central atom
of our dot structure. And the way to do this is to
pick the least electronegative element that we have here,
and then draw the bonds. And so for our
example, we're working with silicon and fluorine. And so we can go ahead
and find those again on our periodic table. Here's fluorine. Fluorine is the most
electronegative element, and so therefore, for
silicon tetrafluoride, we're going to put the silicon
atom at the center of our dot structure, since it is the least
electronegative of those two. So I'm going to start
with silicon here. And I know that silicon has
four bonds to fluorine atoms. I'm going to go ahead and put
in some fluorines right here. So here's some
fluorines like that. So I just drew four
covalent bonds, and we know that each
covalent bond represents two valence electrons, right? So here's two valence
electrons, here's two, so that's a total of
four, six, and eight. So we've represented eight
valence electrons so far in our dot structure. So we originally
had to represent 32. So I'm going to go ahead
and subtract 8 from 32. So 32 minus 8 gives me 24. So now I only have to account
for 24 valence electrons. Let's go back up and
look at our steps again. So let's find out where we are. So we've decided
the central atom, and we've drawn the
bonds, and we just subtracted the
electrons that we used to draw those bonds from the
total that we got in step one. So we're on to step three,
where we assign the leftover electrons to the terminal atoms. So in this case, the terminal
atoms would be the fluorines. So let's go back down here
and look at our dot structure. So fluorine would be
the terminal atoms, and we're going to assign
electrons to those fluorines. But how many do
we need to assign? Well, going back to our
periodic table over here, so fluorine is in
the second period. So pretty good bet it's going
to follow the octet rule here. So we need to surround
each fluorine atom with eight electrons. Each fluorine already has
two electrons around it, so I'm going to go ahead and put
six more around each fluorine, like that. So each fluorine get six
more valence electrons. And since I'm assigning
six valence electrons to four fluorines, 6
times 4 gives me 24. And so therefore
we've now accounted for all of the
valence electrons. And so this should be the
should be the final dot structure here. And so we don't
even need to go on to step four for this molecule. This is a very simple
molecule to draw. Let's go ahead and look
at another example. So now we have CH2O, which
is the dot structure for, which is the molecular formula,
I should say, for formaldehyde. And so if we're following
our guidelines here, the first thing we need to do is
find out the valence electrons. So we need to look and find
carbon, hydrogen, oxygen on our periodic table. So let's go back up here, and
let's find those elements here. So here's carbon. Carbon's in group four,
so therefore carbon will have four
valence electrons. Here's hydrogen, in group
one, so one valence electron. And oxygen is in group six,
so six valence electrons for oxygen. So let's go back down and let's
calculate the total number of valence electrons
that we need to represent in
our dot structure. We have one atom of carbon, so
that's four valence electrons. Each hydrogen is one
valence electron, but we have two of
them, so 1 times 2. And each oxygen is six, and
we have one of them, so six. So 6 plus 4 plus 2 is
12 valence electrons. So we need to represent 12
valence electrons in our dot structure. All right. Let's go back up
to our guidelines and see where we are now. So we've already figured
out the total number of valence electrons. Once again, we don't
have any charges, so we don't need to worry
about the rest of step one. Step two is to assign
the central atom, which is the least electronegative. And remember you're
going to ignore hydrogen for this example. So if we ignore hydrogen,
the central atom is either going to
be carbon or oxygen. And let's look at
our periodic table to figure out the
relative values for those. Oxygen is more
electronegative than carbon. Here we go. Here's oxygen, right
next to fluorine. And here's carbon, over here. So remember your trends
for electronegativity. Oxygen is more
electronegative, therefore carbon is going to be at the
center of our dot structure. So we're going to go ahead
and put carbon in the center. So go ahead and draw carbon. And we know that
the carbon's going to be bonded to two hydrogens. And we know that the carbon is
going to be bonded to an oxygen here. So let's see how many
valence electrons we've accounted for so far. Two, four, and six. So we needed 12. We just used up 6. So right now we have six
valence electrons leftover. OK, so let's go back up and
look at the next step here. So we are to step three. Assign the left over electrons
to the terminal atoms here. All right, so the
terminal atoms. Well, in this case,
our terminal atoms would be hydrogen and oxygen. But we're not going to assign
any electrons to hydrogen, because each hydrogen is now
surrounded by two electrons. And so therefore we're going to
assign those leftover electrons to the oxygen here. And oxygen's going to
follow the octet rule. So it already has two around
it, so it needs six more. So 2, 4, and then 6. And so that actually
takes care of all of our valence electrons, right? So now we have accounted
for all 12 of them. However we're not done with
our dot structure here. So carbon actually follows
the octet rule almost all of the time. And let's go back up here and
let's look at step four here. So for step four, if the central
atom doesn't have an octet, and it usually
does have an octet, you can give it an octet
by creating multiple bonds. So let's look at what we have
so far for our dot structure, and let's see if we can
create multiple bonds somehow. So if I took one of these
electron pairs here, and I moved them
into here, let's see what that would give
us for a dot structure. Now I'd have carbon with
a double bond to oxygen. And this oxygen would have
two lone pairs of electrons. And then these hydrogens
right down here. And so this would be the correct
dot structure for formaldehyde. You have an octet of
electrons around carbon. You have an octet of
electrons around oxygen. And hydrogen has two
electrons around it, which it is happy with. So this is the
correct dot structure. And let me just talk about some
terminology really fast here. So these electrons
in here, so this would be a double bond between
the carbon and the oxygen, these electrons you would
call bonding electrons. Because, obviously, they're
involved in bonding. And then these
electrons out here would be non-bonding electrons,
or lone pairs of electrons. So that's just
terminology that you'll hear people use when you're
talking about dot structures. All right. And then let me just go
ahead and highlight the fact that carbon has an
octet of electrons here. So this would be two, four, six,
and eight electrons surrounding our carbon. All right. Let's do one more example here. Let's look at, this time an ion. So this would be the xenon
pentafluoride cation here. So we need to look and
find xenon and fluorine on our periodic table. So we can figure
out how many valence electrons we're
dealing with here. So let's find xenon first. So here is xenon. It is in group eight, so eight
valence electrons for xenon. And then we've seen fluorine
of course in group seven. So let's go back
down and figure out how many valence electrons
we need to account for here. So we have one xenon, and
that's eight valence electrons. Each fluorine is seven,
and we have five of them. So 7 times 5 is 35. So we have 35 plus
8 gives us 43. Now this is actually not the
total number of electrons that we're going to account
for, because this is an ion. This is a plus 1,
a positive charge. Which means that we
lost an electron. Remember, and electron
is negatively charged, so if you lose a
negative charge, you have a positive
charge left over. And so we're going to
take away an electron. So instead of 43,
we are now going to account for 42 in
our dot structure. Let's go back up
to our guidelines, and I'll show you
where I have that written down right up here. So find the total number of
valence electrons, and then, in this case we had a
positive charge in or ion. So we subtracted an electron. And then, again, step two,
find the central atom, least electronegative. OK. So we know that fluorine is
the most electronegative, so that means we're going
to put xenon in the center. So we'll go ahead, and let's
get some room right down here, and we'll go ahead and put
xenon in the center, like that. On a two five fluorine,
so we can go ahead and put in these bonds around
here, like that. And how many valence electrons
have we accounted for so far? Well, let's see. That would be two, four,
six, eight, and ten. So we had to represent 42,
and we just represented ten. So now we have 32 left over
that we need to represent. And notice, xenon is already
violating the octet rule. It's exceeding the octet rule. And it's OK. It can expand its outer
shell, because xenon is past the third period
on the periodic table, obviously, like what we
talked about earlier. So we can go ahead and
go to the next step now. We're going to assign
the leftover electrons to the terminal atoms. And our terminal
atoms are of course the fluorines, which we know are
going to follow the octet rule. And so, once again,
each fluorine has two electrons around it. So that means I'm
going to give each of the other fluorines six more. So I'm going to go ahead and
put six more valence electrons on each of my fluorines. And so when I'm thinking about
how many valence electrons I have represented
now, right, so I have six more on five fluorines. So 6 times 5 is 30. So I've represented
30 more electrons. And so that means I
have two left over. So two valence
electrons leftover. We haven't had this
situation before. Let's go back up
to the guidelines to refresh our
memory, what we do, when we have leftover
electrons after step three. So we get to step four here. So, if necessary,
in this case it is, we're going to assign any
leftover electrons back to the central atom this time. And if the central
atom has an octet or exceeds an octet, which is
what we have in this example, you are usually done. And we're not going to
worry about formal charges in this video, I'll talk
about them in the next video. So we're going to take
those two electrons and assign them into our xenon. So let's go back
down here, and assign those two extra
electrons to the xenon. And so we would
draw it like that. And so now we are done,
we've represented all 42 of the valence electrons
that we were supposed to. So you can count all
those up if you want to. Now most people will
represent this dot structure by putting brackets here. And putting a positive
charge outside of it. So there's your xenon
pentafluoride cation. So we're going to
a lot more examples for drawing dot structures
in the next several videos, and see how drawing
dot structures allows you to predict the shapes
of different molecules.