What type of intermolecular forces must be overcome in...

Key Concepts

Metallic Bonding

Metallic bonding occurs between metal atoms, where delocalized electrons create a 'sea' of electrons that holds the positively charged ions together. This type of bonding, which contributes to the unique properties of metals, must be disrupted to vaporize a metallic liquid.

Hydrogen Bonding

Hydrogen bonding is a particularly strong type of dipole–dipole interaction that occurs when hydrogen is directly bonded to highly electronegative atoms like oxygen, nitrogen, or fluorine. This interaction significantly elevates boiling points and must be overcome when converting such liquids into gases.

Dipole–Dipole Interactions

Dipole–dipole interactions occur in polar molecules that have permanent dipoles. The electrostatic attractive forces between the partially positive end of one molecule and the partially negative end of another increase the energy required for phase change compared to non-polar molecules.

Intermolecular Forces

Intermolecular forces are the attractions between molecules or atoms that dictate many of a substance’s physical properties, such as boiling and melting points. Overcoming these forces is necessary to convert a substance from a liquid to a gas, as the molecules must separate sufficiently to flow freely in the gaseous state.

London Dispersion Forces

London dispersion forces are the weakest intermolecular attractions and arise from temporary fluctuations in electron density that create instantaneous dipoles. These forces are present in all molecules and atoms, particularly in non-polar substances, and play a major role in determining the boiling points of such liquids.

Transcript

00:01 All right, so we're figuring out what intermolecular forces exist for these molecules when they're going from the liquid state to the gas state.

00:08 Right, now just as a reminder, the big difference between a liquid and a gas is a liquid still has intermolecular forces holding the molecules together.

00:17 A gas does not have intermolecular forces, right? which means that boiling something simply is a matter of adding enough energy to break the intermolecular forces.

00:27 So all we're doing here is identifying what intermolycular forces these have.

00:33 Okay.

00:34 All right.

00:34 Our options here, starting from our weakest and going to our strongest, we have london dispersion forces.

00:42 Okay.

00:42 That is sometimes called induced dipole, induced dipal.

00:50 So whichever term you're familiar with, okay.

00:56 After that, we have diple, dipole, and then hydrogen bonding, and then, ion dipole, which i don't think we're going to have any in that region today.

01:11 All right.

01:12 Now, quick review here.

01:15 London dispersion, everyone does london dispersion.

01:21 If you have electrons, you do london dispersion.

01:24 Okay, so all molecules across the board.

01:26 Dipole, dipole, dipole, the name tries to tell you what the interaction is.

01:31 Die meaning two.

01:32 Pull meaning polar.

01:34 Okay.

01:35 So dipole, dipal means polar molecule.

01:40 Two polar molecules.

01:45 Hydrogen bonding is a type of dipyl dipole dipal.

01:47 It's a very strong type.

01:50 And it requires your molecule to have an oh bond, an nh bond, or an fh bond.

01:59 Okay, so that's your hydrogen bonding.

02:02 Ion dipal is if you actually have an actual ion, which means something with an actual charge and something that's polar.

02:09 I don't think any of these we're going to be.

02:13 Be dealing with.

02:14 Okay, so right off the bat, absolutely everybody in our list is going to have london dispersion.

02:20 So i'm just going to put an ld for london dispersion.

02:23 Now, if induced dipole, induced dipole is the term that your professor uses, go with that.

02:31 Okay, everybody does that.

02:33 All right.

02:33 Next, we have to decide if these molecules are polar or not.

02:37 Okay, so let's start by drawing these out, looking at them.

02:41 Okay, first off, we've got our oxygen.

02:44 Oxygen, if we draw the lewis structure of this, we would have six electrons, six valence electrons in each oxygen.

02:55 So that would be 12 valence electrons.

02:59 So if we draw this out, each bond has two electrons in it.

03:04 So that uses up two electrons so far, 10 left to go.

03:08 Put them around our atoms.

03:09 One, two, three, four, five, six, seven, eight, nine, ten.

03:16 We've used up all our electrons, but that oxygen on the left isn't satisfied yet.

03:22 So we're going to take one of those electron pairs from the right, turn that into a double bond.

03:29 Okay.

03:30 Which effectively erases that guy.

03:34 So here is our oxygen atom here.

03:39 Whether or not it's polar.

03:40 Will it be polar? your atoms have to have different electronegativities.

03:44 Remember, electronegativity is how well an atom can hold on to its electrons.

03:52 Well, we have two atoms that are the same thing.

03:55 So they're sharing equally.

03:56 So this guy is what we call non -polar.

04:01 And what that means for intermolecular forces is the only thing he does is those london dispersions.

04:07 Okay, so he's done.

04:09 We've done everything for that.

04:11 All right.

04:13 B, mercury.

04:17 Okay.

04:17 What's interesting is they didn't give you the formula.

04:19 Mercury is just an element.

04:21 Okay.

04:22 When it's by itself, it's generally hg2.

04:26 There's usually two of them.

04:28 Hg is your symbol for mercury.

04:30 And it's going to have the same situation as the oxygen.

04:34 Okay.

04:34 And i'm not going to bother to draw the whole thing out, but we've got two atoms.

04:39 They're the same thing.

04:40 There's no electronegativity difference.

04:42 So this is also non -polar.

04:47 So again, london dispersion is the only thing he can do...

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