Homework Problem 3F.5

[Talia Dini - 3I]

Hi! Could someone please explain to me how I'm supposed to approach this problem, I'm having some trouble figuring it out.

3F.5 Suggest, giving reasons, which substance in each of the following pairs is likely to have the higher normal melting point (Lewis structures may help your arguments): (a) HCl or NaCl; (b) C2H5OC2H5 (diethyl ether) or C4H9OH (butanol); (c) CHI3 or CHF3; (d) C2H4 or CH3OH.

[Bailey Giovanoli 1L]

So, first start by recognizing that as the number of bonds increases and shortens, the strength of the bonds become stronger. A stronger bond is harder to break and takes more energy to break or melt. Also take a look at the types of bonds, for example, an ionic bond is much stronger and harder to break than a hydrogen bond for example. Hope this helps. Look at your intermolecular forces:)

[David Chibukhchian 2G]

For this problem, it's important to keep in mind which intermolecular forces these molecules will have. For example, the London forces in a molecule of NaCl are going to be stronger than the London forces in HCl, because larger atoms are more polarizable (since their increased electrons make it easier to distort their electron cloud). Because of this, the increased intermolecular forces will cause NaCl to have a higher boiling point.

Meanwhile, for some problems it is a good idea to draw the Lewis Structures to get an idea about which intermolecular forces will be present. For example, in part b, it's helpful to draw the structure of both molecules to show that one has hydrogen bonding while the other doesn't, meaning it will have a higher boiling point due to more intermolecular forces. That's how I approached this problem, I hope that helps!

[Melody Wu 2L]

You want to look at intramolecular and intermolecular forces.
1. Ionic < Covalent character (compare differences in electronegativity)
2. London Forces < Dipole-Dipole < H-Bonding (drawing out lewis structures can help)
Also for two molecules with only London dispersion forces, see which has a larger magnitude (larger atom = more LDF)

[Brendan Duong 1I]

You want to look at intramolecular and intermolecular forces.
1. Ionic < Covalent character (compare differences in electronegativity)
2. London Forces < Dipole-Dipole < H-Bonding (drawing out lewis structures can help)
Also for two molecules with only London dispersion forces, see which has a larger magnitude (larger atom = more LDF)

Larger magnitude would suggest higher boiling point right? Why is this so, bc I thought longer bonds are easier to break...

[Sunny Wu 3A]

For this problem, it's important to keep in mind which intermolecular forces these molecules will have. For example, the London forces in a molecule of NaCl are going to be stronger than the London forces in HCl, because larger atoms are more polarizable (since their increased electrons make it easier to distort their electron cloud). Because of this, the increased intermolecular forces will cause NaCl to have a higher boiling point.

Meanwhile, for some problems it is a good idea to draw the Lewis Structures to get an idea about which intermolecular forces will be present. For example, in part b, it's helpful to draw the structure of both molecules to show that one has hydrogen bonding while the other doesn't, meaning it will have a higher boiling point due to more intermolecular forces. That's how I approached this problem, I hope that helps!

I think a better explanation for NaCl vs HCl is that the strongest IMFs in NaCl is ion-ion while the strongest IMFs in HCl is dipole dipole. Ion-ion forces are significantly stronger than dipole-dipole.

[Brendan Duong 1I]

For this problem, it's important to keep in mind which intermolecular forces these molecules will have. For example, the London forces in a molecule of NaCl are going to be stronger than the London forces in HCl, because larger atoms are more polarizable (since their increased electrons make it easier to distort their electron cloud). Because of this, the increased intermolecular forces will cause NaCl to have a higher boiling point.

Meanwhile, for some problems it is a good idea to draw the Lewis Structures to get an idea about which intermolecular forces will be present. For example, in part b, it's helpful to draw the structure of both molecules to show that one has hydrogen bonding while the other doesn't, meaning it will have a higher boiling point due to more intermolecular forces. That's how I approached this problem, I hope that helps!

I think a better explanation for NaCl vs HCl is that the strongest IMFs in NaCl is ion-ion while the strongest IMFs in HCl is dipole dipole. Ion-ion forces are significantly stronger than dipole-dipole.

How is NaCl an ion ion? I thought ion ion means both atoms have charges or something? I wouldve just thought it was dipole dipole...

[Sunny Wu 3A]

For this problem, it's important to keep in mind which intermolecular forces these molecules will have. For example, the London forces in a molecule of NaCl are going to be stronger than the London forces in HCl, because larger atoms are more polarizable (since their increased electrons make it easier to distort their electron cloud). Because of this, the increased intermolecular forces will cause NaCl to have a higher boiling point.

Meanwhile, for some problems it is a good idea to draw the Lewis Structures to get an idea about which intermolecular forces will be present. For example, in part b, it's helpful to draw the structure of both molecules to show that one has hydrogen bonding while the other doesn't, meaning it will have a higher boiling point due to more intermolecular forces. That's how I approached this problem, I hope that helps!

I think a better explanation for NaCl vs HCl is that the strongest IMFs in NaCl is ion-ion while the strongest IMFs in HCl is dipole dipole. Ion-ion forces are significantly stronger than dipole-dipole.

How is NaCl an ion ion? I thought ion ion means both atoms have charges or something? I wouldve just thought it was dipole dipole...

NaCl is a salt formed by a positively charged Na+ and a negatively charged Cl-