Each of the following molecules contains at least one...

Key Concepts

Lewis Structures

Lewis structures are diagrams that show the bonding between atoms of a molecule and the lone pairs of electrons that may exist. They provide a visual representation of how valence electrons are distributed among the atoms and are fundamental in predicting the arrangement of atoms in a molecule.

Valence Electron Counting

Valence electron counting involves summing the number of electrons available in the outer shells of the atoms involved in a molecule. This is a critical step for constructing Lewis structures as it determines how many electrons need to be placed as bonding pairs or lone pairs to satisfy the octet rule.

Octet Rule

The octet rule is a guideline which states that atoms tend to form bonds in such a way that each atom has eight electrons in its valence shell, achieving a stable electron configuration similar to that of noble gases. This rule is used in drawing Lewis structures, especially for second period elements.

Multiple Bonds

Multiple bonds refer to the presence of more than one pair of electrons shared between two atoms, such as double or triple bonds. These bonds are used in Lewis structures to satisfy the octet rule when single bonds are insufficient, particularly in molecules with less electronegative central atoms or when there are electron deficiencies.

Formal Charge

Formal charge is a bookkeeping tool used to assign charges to atoms in a molecule based on the assumption that electrons in all chemical bonds belong equally to the bonded atoms. Calculating formal charges helps in determining the most stable and plausible Lewis structure by minimizing charge separation.

Resonance Structures

Resonance structures are different possible arrangements of electrons that represent the same molecule. In cases where a single Lewis structure is not sufficient to accurately depict the electron distribution, resonance structures are used to convey the delocalization of electrons, thus providing a more accurate model of the molecule's electronic structure.

Transcript

00:01 When you're drawing lewd structures, some of the molecules will contain double or triple bonds.

00:05 So let's work through some examples to see what this would look like.

00:08 And so with hcn, i separated this molecule into its separate atoms, hcn, and drew their lewis symbols.

00:16 And so this is based on valence electrons.

00:18 And if you're not sure how to do this, you can view my last video.

00:23 And so anyway, it's moving forward.

00:25 So for this atom that, or for the electron that is on the h atom, this one will create a bond with the one electron on the carbon atom.

00:36 And so this is creating this single equivalent bond here.

00:41 And then in between the c and the end, a single bond is also created through these two electrons.

00:47 And so next we'll notice that we have some electrons left over on the carbon and the nitrogen.

00:54 And so these are going to come and form a bond as well on the top and the bottom, giving us a triple bond between the c and the n.

01:05 And so this will be our final answer for part a.

01:09 And it is okay that if we leave the loan pair on the nitrogen atom, and that's because this does not participate in bonding.

01:18 And so next moving on to our part b, i went ahead and separated all of the different.

01:25 Atoms out in this molecule.

01:27 And so there's, i mean, quite a few molecules that we're going to be looking at here.

01:31 So to begin with, you'll notice that carbon is in the middle, and this is because there are more opportunities for carbon to use as electrons to create bonds with the other atoms.

01:42 And so to start with, let's go ahead and do the obvious one, which is between the sulfur and the carbon right here on top.

01:50 And so moving forward, we'll notice that there are.

01:55 It's going to participate in bonding with the two nitrogen atoms.

02:00 But then there's going to be this leftover atoms.

02:03 So don't forget that that's there because that's going to be super important.

02:06 But we'll go ahead and keep working through the obvious one.

02:10 So we'll leave this one right here.

02:13 And then so between the nitrogen and the hydrogen, go ahead and highlight it, there are going to be two bonds formed here, leaving a lone pair on the top nitrogen and this will look exactly the same on the right side.

02:28 And so now you'll notice that we have a single electron on the sulfur along with two lone pairs.

02:37 And so what we can do here to stabilize this and complete it because it would not be correct just leaving those lone electrons there is we're going to form a bond between the sulfur and the carbon using these two electrons.

02:50 And so we can get rid of these and put a double bomb.

02:55 Here.

02:57 And so this answer shown will be our final answer for this part b.

03:01 Now moving on to f2co, i went ahead and drew out each of the atoms lewis symbol.

03:09 And so now let's kind of work through what we're going to do to draw the lewis structure.

03:16 And so creating the bond between the fluorine and the carbon can be our first step and then also doing that for the other fluorine.

03:25 And then you'll notice that we will have a quite obvious one between the c and the o.

03:36 So we can go ahead and draw this out for our, and then keeping in mind that we still have these two lone electrons on the top.

03:53 And so what we're going to do to get rid of those is we are going to create another bond on top there...

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