For many people, tears flow when they cut raw onions. The tear-inducing agent is thiopropionalde

step 1 All right, we're looking at thio propion aldehyde S. Oxide. This is a bit of a complicated one.

For many people, tears flow when they cut raw onions. The...

Key Concepts

Lewis Structures

Lewis Structures are diagrams that depict the arrangement of valence electrons among atoms in a molecule. They show both shared electron pairs forming covalent bonds and unshared pairs (lone pairs) of electrons, providing insight into the bonding and shape of the molecule.

Valence Electrons

Valence electrons are the electrons in the outermost shell of an atom that participate in chemical bonding. Accurate counting of these electrons is essential when constructing Lewis structures to ensure all atoms achieve stable electron configurations.

Octet Rule

The Octet Rule is a chemical guideline stating that atoms tend to form bonds so that they each have eight electrons in their valence shell, mimicking the electron configuration of noble gases. This rule helps predict the number of bonds and lone pairs in a molecule.

Formal Charge

Formal Charge is a tool used in Lewis structures to assess the distribution of electrons among atoms. It is calculated by comparing the number of valence electrons in an isolated atom to the number assigned to it in the molecule, with a lower overall formal charge indicating a more stable structure.

Covalent Bonding

Covalent Bonding involves the sharing of valence electrons between atoms, resulting in the formation of chemical bonds. Understanding covalent bonds is crucial for building Lewis structures and predicting the geometries of molecules.

Transcript

00:01 All right, we're looking at thio propion aldehyde s.

00:06 Oxide.

00:07 This is a bit of a complicated one.

00:09 Let's start by adding up our valence electrons, and then we'll talk about how you set up what this compound looks like.

00:16 So i've got our periodic table out.

00:17 It's labeled with our valence electrons by column.

00:21 So we're going to start with carbon, which has four, and there's three of those in our compound.

00:28 Hydrogens have one valence electron, and there's six of them.

00:33 Oxygen has six.

00:36 Sulfur also has six.

00:40 Add those all up and you should get a total of 30 valence electrons that we're going to be working with.

00:47 Now for how these connect.

00:49 It is very common in compounds that have a lot of carbons and hydrogens.

00:53 For the carbons to all be linked up into a chain and that's going to be the case here as well.

00:58 So we're going to have three carbons just like that.

01:04 And then the hydrogens are going to be spread among them.

01:09 Part of the reason that we know that is where it says propio.

01:13 Prop means three carbons.

01:16 They're almost always linked together in a chain like this.

01:19 Now, the rest of it is actually a functional group, but it's fairly advanced even for an advanced chemistry class.

01:26 An aldehyde usually means you have an oxygen and a hydrogen, both attached to a carbon.

01:33 But when they put the thio out front, thio is where the sulfur comes in on this.

01:40 So it's an aldehyde made with a sulfur instead of oxygen, and then we tack an oxygen onto the end.

01:48 So the rest of this connection then is going to be a sulfur and then an oxygen.

01:55 After that, we're going to fill in the electrons, or fill in the hydrogens around the carbons.

02:01 Carbons really like to have four bonds.

02:05 So in general, carbons will gain hydrogens will gain hydrogens until they get up to their specified number.

02:13 Now, we only have six hydrogens.

02:16 Technically, there's one more spot open right here for another hydrogen.

02:20 I expect there's going to be a double bond with the sulfur in order to account for that because that carbon isn't fulfilled yet.

02:30 All right, counting our electrons, we've done one, two, three, four, five, six, seven, eight, nine, ten bonds so far.

02:37 Ten bonds times two electrons each, mean we've used up 20 of our electrons.

02:43 We've only got 10 left.

02:46 Now, real quick, the way i knew it was going to be s and then o on the outside, rather than say o and then s, has to do with electronegativity.

02:57 Electron negativity is how well something holds on to electrons.

03:02 The trend is it increases as you go up into the right on the periodic table, as you get closer to fluorine, you become more electronegative, which means you hold on to electrons tighter.

03:15 Now, in this case, sulfur and oxygen are both right next to fluorine.

03:20 Oxygen, if you don't already know, is the second most electronegative atom after fluorine, which is actually something important to hold on to, just so you know how to compare it if you're dealing with sulfur or chlorine.

03:32 So oxygen is more electronegative.

03:34 Now, the atoms on the outside of the, molecule, that tends to be where the electronegative ones go.

03:43 So usually, and it comes down to doing formal charges and things, but usually the formal charge on the outer edges tends to be zero or negative, whereas formal charges on the inner part of the molecule tend to be zeros or positives.

03:59 So since the oxygen wants to be zero or negative, it wants to have more electrons.

04:04 I knew the oxygen that's going to go on the outside of this.

04:09 Okay.

04:10 That being said, let's go ahead and fill it in.

04:12 We care about the octets.

04:13 We'll start on the outside and work our way in.

04:16 Those hydrogens are all fine.

04:18 They only want two electrons.

04:20 One bond equals two electrons, so they're good.

04:22 We can look at that oxygen.

04:24 It only has two electrons right now, and it wants eight.

04:29 So now he's good.

04:31 That takes off six of these.

04:33 We've got four left.

04:35 So next we can put him on the sulfur.

04:39 That drops off four.

04:40 We're down to zero.

04:42 Now, if you had wanted to, you could put two on the sulfur and two on the carbon.

04:47 Either way, we're going to run into the same situation.

04:51 That carbon is not fully, it doesn't have a full octet.

04:58 It only has six.

05:00 We can't add more electrons because we've already used them all up, which means we're going to do a double bond.

05:05 Taking two of the electrons off the sulfur, we're going to move those and create a double bond.

05:13 Sulfur doesn't lose anything by doing this carbon gains two electrons so that fulfills the octets now anytime you start getting a complicated structure like this where there's a lot going on it's wise to check your formal charges okay so to do that we're going to start looking on the we'll start on the left and work our way around remember that formal charge is the number of valence electrons that an atom has just based on the periodic table minus how long many minus half of the bonding electrons attached to that atom.

05:51 The bonding electrons are the ones tied up in bonds.

05:55 And then all of the non -bonding.

05:57 Those are the individual dots.

06:00 An easier way to keep track of this.

06:01 The formal charge of an atom is your valence, minus how many sticks are attached to that atom, minus how many dots are on it.

06:10 It's a little bit easier.

06:12 Okay, starting with, we'll say we'll start with the carbon out here on the edge...

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