Can dimethyl ether form H-bonding?

Question

CH3OCH3, its non-polar.... it has london, but what else forces are there?

Answer 1

dimethyl ether can NOT form H-bonding - only molecules where H is bonded directly to O, N or F can you have H-bonding as in intermolecular force - in the case of dimethyl ether, the O is bonded to each C and all of the H's are bonded to C

H3C - O - CH3

the C-O bonds are polar bonds - if you arrange the two CH3 groups to form a right angle with one on the left of the oxygen and one below the oxygen and draw a diagonal line through the structure from top left to bottom right, you will get the oxygen on the right side of the line and everything else on the left side - this means that this molecule will form a dipole and there will be dipole-dipole forces present as well as London dispersion forces

Answer 2

The strength of bonds within the molecule isnt related to boiling point, what is important is the strength of the attractions between the molecules. Because the alcohol has an OH group, you get what is known as hydrogen bonding, a relatively strong type of intermolecular attraction. The ether doesnt have this and thus the interaction between molecules is weaker.

Answer 3

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Viscosity depends on intermolecular forces. Ethanol is an alcohol with H-bonding so will have larger intermolecular forces and greater viscosity than the ether that does not have H-bonding. Answer 3.

Answer 4

hydrogen bonding only occurs between water molecules (H2O) or HF or HN.
what is the INTERMOLECULAR BONDING - HYDROGEN BONDS
The hydrogen is attached directly to one of the most electronegative elements, causing the hydrogen to acquire a significant amount of positive charge.

Each of the elements to which the hydrogen is attached is not only significantly negative, but also has at least one "active" lone pair.

Lone pairs at the 2-level have the electrons contained in a relatively small volume of space which therefore has a high density of negative charge. Lone pairs at higher levels are more diffuse and not so attractive to positive things.

The + hydrogen is so strongly attracted to the lone pair that it is almost as if you were beginning to form a co-ordinate (dative covalent) bond. It doesn't go that far, but the attraction is significantly stronger than an ordinary dipole-dipole interaction.

Hydrogen bonds have about a tenth of the strength of an average covalent bond, and are being constantly broken and reformed in liquid water. If you liken the covalent bond between the oxygen and hydrogen to a stable marriage, the hydrogen bond has "just good friends" status. On the same scale, van der Waals attractions represent mere passing acquaintances!

Water as a "perfect" example of hydrogen bonding

Notice that each water molecule can potentially form four hydrogen bonds with surrounding water molecules. There are exactly the right numbers of + hydrogens and lone pairs so that every one of them can be involved in hydrogen bonding.

This is why the boiling point of water is higher than that of ammonia or hydrogen fluoride. In the case of ammonia, the amount of hydrogen bonding is limited by the fact that each nitrogen only has one lone pair. In a group of ammonia molecules, there aren't enough lone pairs to go around to satisfy all the hydrogens.

In hydrogen fluoride, the problem is a shortage of hydrogens. In water, there are exactly the right number of each. Water could be considered as the "perfect" hydrogen bonded system.

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Answer 5

Methoxymethane (to give it its proper name) is a polar molecule which cannot hydrogen bond with itself, but can certainly hydrogen bond with a solvent such as water.