Write correct Lewis structures and assign a formal charge to each atom. (a) CH4 CHO (b) N3^- (c)

Write correct Lewis structures and assign a formal charge to...

Key Concepts

Lewis Structures

Lewis structures are diagrams that show the bonding between atoms in a molecule and the lone pairs of electrons that may exist. They are constructed by arranging the atoms, drawing bonds between them, and placing electrons as dots around the atoms. This diagrammatic approach helps determine the overall connectivity in a molecule and predict molecular geometry.

Formal Charges

Formal charge is a bookkeeping method used to indicate the distribution of electrons in a molecule. By comparing the number of valence electrons in the neutral atom with the number assigned in the Lewis structure, one can calculate the formal charge for each atom. This helps in assessing the stability of different resonance forms and in identifying the most plausible structure among them.

Valence Electrons Counting

Counting valence electrons is a fundamental step in constructing Lewis structures. It involves summing the number of electrons available for bonding from each atom, accounting for any extra electrons due to ionic charges. This total is then used to distribute electrons in bonds and lone pairs, ensuring that the structure conforms to known chemical rules.

Octet Rule

The octet rule is a guideline in chemistry that atoms tend to form structures in which they have eight electrons in their valence shell, similar to the electron configuration of noble gases. While there are exceptions, this rule is central to predicting the arrangement of electrons in molecules and thus plays a vital role in drawing accurate Lewis structures.

Transcript

00:01 We're doing some loose structures and then doing formal charges to go with them.

00:04 The first step we have is number a, which is ch4, cho.

00:09 Now, there's actually a typo here.

00:12 By referencing the book, this is actually ch3, cho.

00:17 Hopefully the website will get it fixed, but last time i checked it, it was still ch4.

00:21 So the book says ch3, cho, which will make this low structure, make a lot more sense.

00:27 Okay.

00:28 Now, to draw our image, we've got to decide how many electrons are going to be in it.

00:31 Of course, with lewis structures, we only care about valence electrons, which are the outermost atoms of each bond, or each atom.

00:40 I've already set up the number of valence electrons on the periodic table based on column.

00:45 All right, so first column, all those atoms would have one.

00:49 Second column, they don't have two.

00:51 Right, so we can quickly figure out how many valence electrons are going to be in our picture.

00:55 So first up, we have a carbon, which has four valence, plus hydrogen, which has one, and there's three of those right there.

01:06 I'm just doing it in order.

01:08 Then we have another carbon, so another four, another hydrogen.

01:13 And then lastly, we have an oxygen, which has six.

01:17 Add those up, and you should get a total of 18.

01:22 So next we go ahead and draw our molecule.

01:24 Now, the way that this formula is written, where it kind of seems out of order instead of just like c2, h4o, the reason they've done that is they're trying to tell you what the structure of this compound is.

01:34 In general, the carbons will be connected to each other.

01:37 So these two carbons will be side by side.

01:40 And then whatever follows after a carbon is what is connected to that carbon.

01:46 So ch3 means that first carbon has three hydrogens on it.

01:52 And then cho means that that second carbon has a hydrogen and it has an oxygen.

01:58 So this is our generic structure.

02:02 All right, to track our electrons, let's see how many we use so far.

02:05 If there's two electrons in each bond, and we've drawn one, two, three, four, five, six, that means we've used 12 electrons so far.

02:15 So we're going to subtract off 12.

02:17 We've got six more to go, right? so first we want to check if our terminal atoms or our outer atoms all have their octets.

02:26 All the hydrogens just want to have two electrons.

02:29 They have one bond.

02:30 They're happy.

02:31 All the hydrogens are good.

02:32 The oxygen, though, currently has one bond.

02:35 Attached to it.

02:37 So it thinks it owns two electrons and it wants to have eight.

02:42 So we need to add six more to that oxygen.

02:47 So we'll subtract six off of here.

02:49 And that's the rest of our electrons.

02:51 But we're not quite done yet because one of our inner atoms, our carbons, does not have an octet.

02:57 The carbon on the left has four bonds.

03:00 So he's fine.

03:01 Four times two is eight.

03:02 The carbon on the right only has three bonds.

03:06 Three times two is six.

03:07 Doesn't have enough electrons yet.

03:10 Now we can't add, we cannot, for example, add a lone pair on that carbon because we're out of electrons.

03:19 Can't make them up.

03:20 So this is where double bonds come into play.

03:23 What a double bond does is it takes a lone pair from an outer atom and it moves it into a bond.

03:32 So instead of being on that oxygen, it's going to create a double bond.

03:36 So we're going to put a double bond right here.

03:39 And then i'll go ahead and erase.

03:45 Right.

03:46 So we have not added new electrons at all.

03:48 We just took a lone pair that was already there and we moved it.

03:52 Now that oxygen still believes it owns both of those.

03:55 So the oxygen hasn't lost anything.

03:56 It still has eight.

03:57 And the carbon now believes it has eight.

04:00 So we're good.

04:00 This is your final structure.

04:05 All right.

04:05 Next we've got to figure out the formal charges.

04:08 Formal charges have a bit of a long definition to them, a little bit of a wordy definition, and then they have a simplified one.

04:16 And so i'll give you both.

04:18 The long definition is a formal charge for an atom.

04:22 Each atom will have its own formal charge, first of all.

04:25 A formal charge for an atom equals how many valence electrons that atom has based on the periodic table.

04:36 Okay, so based on how many valence electrons it had.

04:39 Coming into the molecule before it ever connected, minus the one half of the bonding electrons attached to that atom.

04:51 Bonding electrons are the electrons tied up in bonds.

04:55 And then subtract off all of the non -bonding electrons.

04:59 Those are, oh, goodness.

05:02 I went to write and i hit the scroll bar.

05:06 Non -bonding electrons.

05:07 I'm out of room.

05:07 Non -bonding electrons.

05:09 These are the dots.

05:11 Okay, these are what go into the loan pairs.

05:14 A lot of teachers don't give you this equation on exams or anything like that.

05:20 And so a really simplified way of doing this would be, i'll write it right here.

05:28 Formal charge equals your valence electrons, minus for a single atom that's minus the sticks attached to it, minus how many dots are on that so let's walk through that.

05:43 Let's see if i can get a good scroll right there.

05:46 Alright, so to start with, we'll look at one of the hydrogens.

05:50 Hydrogen comes into the game with one valence electron.

05:53 Same number that we use to add up the valence at the beginning.

05:57 Okay? that electron has one bond attached to it, one stick.

06:02 So minus one.

06:04 And then it doesn't have any dots.

06:06 So his formal charge is 1 minus 1, which is 0.

06:10 So his formal charge is 0, which is 0, which is good.

06:14 Good.

06:15 Zeros are ideal.

06:16 That's what you want.