Halogens form a variety of covalent compounds with each other. For example, chlorine and fluorin

Halogens form a variety of covalent compounds with each...

Key Concepts

Hypervalency and Octet Expansion

Hypervalency refers to the ability of certain atoms, typically those in the third period and beyond, to accommodate more than eight electrons in their valence shell. This concept is important when considering compounds like ClF3 and ClF5, where chlorine acts as a central atom with more than four substituents. In contrast, atoms in the second period, like fluorine, are generally limited by the octet rule and cannot serve as central atoms in hypervalent compounds, which explains why a molecule like FCl3 would be unstable.

Lone Pair Repulsion and Bond Angle Deviations

Lone pairs of electrons occupy more space than bonding pairs, leading to greater repulsive forces between electron groups. This repulsion can cause deviations from ideal bond angles in the molecular geometry. In the case of molecules such as ClF3, the placement of lone pairs in equatorial positions influences the observed bond angles, reducing them from the idealized angles of the electron domain geometry.

Electronegativity and Atomic Size in Central Atom Selection

The choice of the central atom in a polyatomic molecule is influenced by both its electronegativity and atomic size. Larger atoms with lower electronegativity, such as chlorine, are able to accommodate more electron pairs around them, making them suitable central atoms for hypervalent compounds. In contrast, extremely electronegative and small atoms like fluorine cannot efficiently accommodate additional substituents without destabilizing the electron configuration, which is why a compound with fluorine as the central atom (FCl3) would be unlikely to be stable.

VSEPR Theory

The VSEPR (Valence Shell Electron Pair Repulsion) theory is used to predict the three?dimensional shapes of molecules based on the repulsions between electron pairs (bonding and lone pairs) around a central atom. This theory provides guidance on how to arrange electron pairs in space to minimize repulsion, which in turn determines the molecular geometry and bond angles.

Electron Domain Geometry and Molecular Geometry

Molecular geometry considers the spatial arrangement of both bonding pairs and lone pairs around a central atom. The overall electron domain geometry is determined by the number of electron groups (bonding pairs plus lone pairs), while the molecular geometry is the arrangement observed when the lone pairs are not counted. For example, in molecules like ClF3 and ClF5, different numbers of lone pairs on the central atom lead to T-shaped and square pyramidal geometries respectively, with bond angles adjusted from ideal values due to lone pair repulsion.

Transcript

00:03 Okay, so here we're given different halogen compounds with chlorine and fluorine.

00:11 And we are asked to give the bond angles and the geometry of these compounds.

00:17 And then to discuss our opinion on the stability of whether fcl3, whereby the fluorine is the central atom.

00:27 Okay, let's start by drawing the lewis structures of each of these.

00:32 So, clf, let's start there.

00:37 C -l bond to f.

00:41 Okay.

00:42 We know that halogens have seven valence electrons.

00:47 From the periodic table, we can see that.

00:50 So chlorine will have seven and fluorine will have seven.

00:53 So that should be 14 valence electrons.

00:58 In forming this single bond, this single covalent bond, we used up two minus.

01:05 Two.

01:06 So we have 12 valence electrons left.

01:09 So now we start putting lone pairs around the most electronegative atom.

01:15 So let's go ahead and add some one, two, three, four, five, six.

01:21 So that's six.

01:23 We have six left.

01:26 And we can help chlorine complete its octet by putting six lone pairs around it.

01:32 Okay.

01:34 Let's erase this.

01:36 So this is the structure of, our first compound.

01:40 Now, it asks us the molecular structure.

01:44 And since there's only two things bound together, the only geometry it can have is linear.

01:53 Okay.

01:54 There's really no other possibility.

01:56 And so since it's linear, it'll be about 180 degrees.

02:03 Next one, clf3.

02:06 So chlorine is our central atom.

02:08 Let's draw that in.

02:10 F, f, f.

02:13 Okay? and i'm drawing.

02:15 This flat, so don't pretend like this is going to give you any information about the geometry.

02:20 This is not that it's going to be trigonal planar.

02:24 Okay? so we need to count the valence electrons.

02:27 So 7, 14, 21, 28.

02:31 There should be a total of 28 valence electrons for four halogens.

02:40 So by forming these three covalent bonds, we've already used two for six.

02:48 Okay, so we have 22 valence electrons left.

02:54 Let's go ahead and give the more electronegative atom, fluorine, a full octet with lone pairs.

03:00 So 1, 2, 3, 4, 5, 6, 7, 8, 9, 10, 11, 12, 13, 14, 15, 16, 17, 18.

03:08 Okay.

03:09 So we used 18 electrons.

03:13 Okay.

03:14 And we have four valence electrons left.

03:18 Well, chlorine currently has six electrons around it, so we can add a lone pair.

03:22 Pair to give it an octet as well.

03:27 But now we have two valence electrons left.

03:30 And we have to decide where to put them.

03:32 We have two options.

03:33 We could put them on a fluorine or a chlorine.

03:37 You may know that the larger an atom becomes, the more able it is to accept more electrons and break the octet rule.

03:46 So chlorine is going to have two lone pairs because it can handle it.

03:53 It has an expanded octet.

04:00 Now, we need to look here and see that we have three bonded pairs, so three bonds, and two lone pairs.

04:12 So if you look in a chart, you will see that the structure will actually be.

04:28 I'm going to draw the loan pairs on wedges, just to show.

04:33 So these loan pair is away from you and this one is toward you.

04:36 So the geometry we have here is t -shaped.

04:48 And t -shaped geometry, it's usually about 90 degrees.

04:54 However, lone pairs take up a lot of space, and so they tend to make the other angles slightly smaller than the ideal number.

05:07 So slightly less than 90 degrees is what we would expect for our angles.

05:12 Okay, i rewrote that a little bit...

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