i am supposed to explain the order of boiling points from intramolecular forces. the order i ended up coming up with was this:
diethyl ether 34.6 degrees C
hexane 69
ethanol 78.3
pyridine 115.2
acetic acid 117.9
octane 126
I know all of them have
-London/Van der Waal forces,
and that
-diethyl ether has a dipole moment,
-ethanol and acetic acid have H-bond forces,
and that
-pyridine is in a nice "stackable" planar ring.
According to intramolecular forces alone, shouldn't acetic acid and ethanol (because they have H-bond AND London forces) have the highest boiling points?
And according to geometry shouldn't pyridine? I'm so confused!
i realized just now that acetic acid also has a dipole moment! why does octane (which looks so similar to hexane) have a higher boiling point?
Acetic acid has a dipole moment AND it has H bonding. Diethyl ether has no dipole moment. Look at the molecule:
CH3CH2OCH2CH3
Octane has a higher boiling point than hexane because the molar mass of octane is higher. You must factor in the molar mass in addition to van der Waals forces and H bonding and dipole moment.
thank you and happy halloween!
Ah, the mysterious world of boiling points! Let's uncover the humor behind it.
First, we have diethyl ether at a toasty 34.6 degrees C. It's got some London forces and a dipole moment, making it quite an interesting molecule. But don't let its dipole fool you, it's not the hottest one on the block.
Next up, we have hexane, feeling a bit more confident at 69 degrees C. Like diethyl ether, it rocks those London forces, but lacks any fancy H-bonding. Still, hexane knows how to warm things up.
Now, let's shake things up with ethanol at a steamy 78.3 degrees C. It not only has London forces but also boasts some charming H-bonding. It's like a molecular matchmaker, bringing molecules closer together with its H-bonding powers.
But wait, here comes pyridine strutting in at 115.2 degrees C. Pyridine's secret weapon is its planar ring, making it "stackable" like a neat tower of blocks. Its unique geometry gives it an advantage in the boiling point game.
Hold on tight because acetic acid is heating things up at 117.9 degrees C. With its dipole moment and H-bonding, it's double trouble for other molecules. Acetic acid knows how to create strong connections and keep them boiling, heating up the competition.
Finally, we have octane, the heavyweight champion at 126 degrees C. Octane may look like hexane's long-lost twin, but its high molar mass gives it an edge. It's like the big brother of the hydrocarbon family, making it harder to break free from the clutches of boiling.
So there you have it, my confused friend. Boiling points are a blend of molecular forces, geometry, and molar mass. Each molecule brings its own flair to the boiling point party, making it a hot contest. Happy Halloween, and may your chemistry experiments be a real scream!
Happy Halloween to you too!
You've made some good observations regarding the boiling points of the given compounds and the corresponding intramolecular forces at play. Let's break it down step by step:
1. London/Van der Waals forces: These forces are present in all the compounds you listed. They arise due to temporary fluctuations in electron density and induce temporary dipoles in adjacent molecules. London forces generally increase with increasing molecular size.
2. Dipole moment: Diethyl ether and acetic acid have dipole moments. A dipole moment is a result of differences in electronegativity between atoms in a molecule, leading to a separation of charge and the presence of polar bonds. As a result, molecules with dipole moments generally have stronger intermolecular forces compared to those without.
3. Hydrogen bonding: Both ethanol and acetic acid can form hydrogen bonds due to the presence of hydrogen attached to highly electronegative atoms (oxygen and nitrogen). Hydrogen bonding is stronger than London forces and dipole-dipole interactions and can significantly affect the boiling point.
4. Molecular shape and stacking ability: Pyridine's planar ring structure allows for efficient stacking of its molecules, resulting in stronger van der Waals forces and potentially higher boiling points.
Now, let's address your confusion:
You are correct that based on hydrogen bonding alone, ethanol and acetic acid would be expected to have higher boiling points than the other compounds. However, boiling points are not solely determined by one type of intermolecular force. In this case, the molecular size/mass also plays a significant role.
Octane, despite its similarity to hexane, has a higher boiling point due to its larger molecular size. The increase in molecular size leads to stronger London forces, which contribute to increased boiling points.
So, in summary, the boiling points of these compounds are determined not only by the type of intramolecular forces (London/Van der Waals, dipole-dipole, and hydrogen bonding) but also by factors such as molecular size and molar mass.
I hope this clarifies your confusion. If you have any more questions, feel free to ask!
To explain the order of boiling points based on intramolecular forces, you have correctly identified the following forces present in each compound:
1. Diethyl ether: London/Van der Waals forces and a dipole moment.
2. Hexane: Only London/Van der Waals forces.
3. Ethanol and acetic acid: Both have London/Van der Waals forces and hydrogen bonding.
4. Pyridine: Planar ring structure with London/Van der Waals forces.
5. Octane: Only London/Van der Waals forces.
Based solely on intramolecular forces, one might expect that compounds with both London/Van der Waals forces and hydrogen bonding (such as ethanol and acetic acid) would have the highest boiling points. However, other factors like molar mass also play a role.
Acetic acid, with a dipole moment in addition to hydrogen bonding and London/Van der Waals forces, has a higher boiling point than ethanol, which only has hydrogen bonding and London/Van der Waals forces. Acetic acid's higher boiling point can be attributed to its stronger intermolecular forces.
Now, regarding octane and hexane, although they have similar molecular structures and exhibit only London/Van der Waals forces, octane has a higher boiling point. This is because octane has a larger molar mass (114.22 g/mol) compared to hexane (86.18 g/mol). Molecules with higher molar mass generally have stronger van der Waals forces, leading to higher boiling points.
In summary, the boiling point order considering intramolecular forces, dipole moments, hydrogen bonding, and molar mass is as follows:
Diethyl ether < Hexane < Ethanol < Acetic acid < Pyridine < Octane
I hope this clears up your confusion! Feel free to ask if you have any more questions. Happy Halloween!