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Class 11 Chemistry (India)
Course: Class 11 Chemistry (India) > Unit 4
Lesson 8: Hydrogen bondingHydrogen bonding
AP.Chem:
SAP‑5 (EU)
, SAP‑5.A (LO)
, SAP‑5.A.4 (EK)
, SAP‑5.A.5 (EK)
, SAP‑5.B (LO)
, SAP‑5.B.1 (EK)
Hydrogen bonding is a special type of dipole-dipole interaction that occurs between the lone pair of a highly electronegative atom (typically N, O, or F) and the hydrogen atom in a N–H, O–H, or F–H bond. Hydrogen bonds can form between different molecules (intermolecular hydrogen bonding) or between different parts of the same molecule (intramolecular hydrogen bonding). Created by Sal Khan.
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Well, I still didn't get the difference between dipole-dipole forces and hydrogen bonds. Could anyone explain?(12 votes)
Hydrogen bonds are a special type of dipole-dipole forces. In hydrogen bonding, the dipole-dipole attraction occurs between hydrogen and a highly electronegative atom(O,F,Cl,etc). Hydrogen bonds are much stronger than a general dipole-dipole force.These are the only differences, otherwise everything is same. I hope this helps.(24 votes)
why are there differences between the boiling points of compounds with hydrogen bonds?(7 votes)
Compounds that have hydrogen bonding in general have higher boiling points than those that don't. This graph gives a good idea of the effect:
https://ch301.cm.utexas.edu/svg/H-bond-trend.svg(4 votes)
Why do we consider only hydrogen bonds? Why not something like Cesium or Francium bonds? Shouldn't attraction between the δ+ and δ- parts be more?
I mean, why don't we consider a bond between a more electro+ve atom (wrt H) and a highly electro-ve atom? Wouldn't the δ+ and δ- (as well as the net dipole moment)gained be greater as the change in electro-vety increases?
Is this because H is smaller and can hence form more stable bonds?(2 votes)
Well atoms like cesium or francium have very small electronegativity values. And so pairing them with the electronegative atoms commonly associated with hydrogen bonding (nitrogen, oxygen, and fluorine) will have a large enough electronegativity difference that their bonding is no longer considered polar covalent, but rather ionic. So we'll start having ionic interactions instead of dipole-dipole interactions which includes hydrogen bonding.
The reason we focus on hydrogen specifically is a result of several reasons. First, it's simply a commonplace element in many crucial molecules like water. Second the electronegativity difference between hydrogen and small electronegative atoms is large enough to be considered polar covalent, but not so much that it becomes ionic. Third, since these atoms all quite small, the hydrogen atom on one molecule can approach the small electronegative atom (again N, O, and F) on another molecule very closely. So fundamentally it's the same interaction which occurs in dipole-dipole, but because of the atom's sizes this attraction is enhanced and is stronger than other conventional dipole-dipole interactions.
Hope that helps.(6 votes)
My textbook shows HCl under examples of dipole-dipole Interaction but not under Hydrogen Bonds. So do HCl molecules not form Hydrogen bonds? If so, why?(2 votes)- Well for hydrogen bonding to occur you want the hydrogen atom bonded to an electronegative atom (acting as the hydrogen bond donor) interacting with another electronegative atom (acting as the hydrogen bond acceptor). You also want these electronegative atoms to be small the that donor and acceptor can approach each other more closely adding to the strength of the hydrogen bond. Usually this means atoms like fluorine, oxygen, and nitrogen, but this can also include atoms like carbon, chlorine, and sulfur. Part of the reason why these other atoms aren't commonly recognized as engaging hydrogen bonding is that compared to the traditional atoms, their hydrogen bonding is quite weak. And this is due to the nontraditional atoms either being not as electronegative (in the case of carbon), or not as small (in the case of chlorine and sulfur) compared to the traditional atoms.
But nowadays it is recognized more elements than just fluorine, oxygen, and nitrogen engage in hydrogen bonding (would would include hydrochloric acid). The strengths of these hydrogen bonds just vary widely.
Hope that helps.(6 votes)
- what about dipole induced dipole forces......can u please explain deeply about that(1 vote)
Dipole-induced dipole forces arise between polar sites in a molecule and non-polar sites in neighboring molecules. The polar site induces the opposite charge in the non-polar sites creating relatively strong electrostatic attractions. Generally, this is the strongest intermolecular force between gaseous molecules.(4 votes)
- Isn't Chlorine more electronegative than Nitrogen? Why is nitrogen usually seen in a hydrogen bond, but chlorine isn't?
Also, what makes hydrogen special? Wouldn't francium or another element be better because of its low electronegativity? I'm having trouble understanding why hydrogen bonds are significant and why they are so strong.(2 votes)- I believe Nitrogen and Chlorine bear the same value for electronegativity, and what makes Nitrogen a better atom for Hydrogen to H-Bond with is the size of its highest occupied molecular orbital. Nitrogen has a smaller highest occupied molecular orbit than Chlorine. A greater difference between the lowest occupied molecular orbital (which belongs to hydrogen in both cases) and highest occupied energy orbital will lead to less mixing between orbitals. Since Nitrogen carries a smaller unit area than Chlorine, it has a greater charge density which then leads to a greater attraction.(2 votes)
At2:52Sal says "In fact all the molecules have similar molar mass" but the molar mass of all the molecules are not same
For instance, NH3 has a molar mass of 17g/mol while HF has a molar mass of 20g/mol
What is the thing thatI am missing?(1 vote)- Similar molar masses means that they have approximately equal molar masses, not that they are exactly equal. Their values are almost the same, but are not completely the same.(2 votes)
Why does pentanone not have hydrogen bonds even though it contains oxygen and hydrogen?(1 vote)- For hydrogen bonding to occur, there not only needs to be a small electronegative atom like oxygen, but also have that small electronegative atom must be directly bonded to hydrogen.
Pentanone (both the 2-pentanone and 3-pentanone isomers) are known as ketones in organic chemistry. This means there is a carbonyl group where a carbon is double bonded to an oxygen and that carbon has two additional carbon groups bonded to it. So the sole oxygen in a ketone is bonded to a carbon and only a carbon. So since the oxygen is not directly bonded to a hydrogen, it doesn’t display hydrogen bonding.
Hope that helps.(2 votes)
- Why hydrogen bond is much stronger than other dipole-dipole force? The electronegativity of hydrogen is 2.1, which is pretty large compared to most of the elements, why there is still a highly polarized bond between H and O.(1 vote)
- A polar covalent bond arises when there is a large electronegativity difference between two bonding atoms. There isn’t a single agreed range of what constitutes a polar covalent bond, but approximately an electronegativity difference between 0.4 – 2.0 is considered polar covalent. At 0.4 and below is considered nonpolar covalent, while greater than 2.0 is considered ionic.
Hydrogen bonding is most common with electronegative small atom in the second period, namely nitrogen, oxygen, and fluorine. Using Pauling electronegativity values, hydrogen has an electronegativity of 2.20, nitrogen is 3.04, oxygen is 3.44, and fluorine is 3.98. The electronegativity differences are therefore: N-H = 0.84, O-H = 1.24, F-H = 1.78, which would place them all within the polar covalent range. Since all three are polar bonds they all make molecules with the correct geometry polar and therefore experience dipole-dipole attraction.
So the reason for hydrogen bonding is not solely due to a large electronegativity difference since there are other polar molecules with dipole-dipole, but not hydrogen bonding. The key is that we have hydrogen, a very small atom, bonded to other small atoms. This small size of both atoms means molecules with these N-H, O-H, and F-H groups can approach each other quite close which results in an even stronger attraction than simply dipole-dipole. In that sense we can think of hydrogen bonding as a stronger case of dipole-dipole.
There is a range of strengths for dipole-dipole and hydrogen bonding depending on the chemical environment and the types of molecules involved, but generally hydrogen bonding is stronger than dipole-dipole. If we quantify strength as the amount of energy required to pull molecules apart, then hydrogen bonding ranges from 4-50 kJ/mol while dipole-dipole only ranges from 2-8 kJ/mol.
Hope that helps.(2 votes)
I don't get why H2S has a higher boiling point than HCl. They have very similar atomic mass(Similar LDF), and Cl has a greater electronegativity than S, so it is supposed to also have more dipole moment(Dipole-dipole forces). And I also don't understand if HCL can be considered an hydrogen bond, as hydrogen is bonded with an halogen.(1 vote)- The primary factor for the difference in boiling points between hydrogen sulfide and hydrogen chloride is their hydrogen bonding. Both molecules experience hydrogen bonding, and individually a hydrogen bond in hydrogen chloride is stronger than one in hydrogen sulfide. However hydrogen sulfide has more hydrogen atoms and can therefore make more hydrogen bonds per molecule compared to hydrogen chloride. Stronger hydrogen bonding in hydrogen sulfide translates to a higher boiling point as compared to hydrogen chloride. You can see this same trend between water and hydrogen fluoride for the second period elements.
There's no reason why hydrogen chloride would be excluded from performing hydrogen bonding considering a molecule like hydrogen fluoride experiences it (and fluorine is a halogen). The most common and strongest hydrogen bonding occurs with nitrogen, oxygen, and fluorine but we can see examples of it occurring with elements like sulfur, chlorine, and carbon even.
Hope that helps.(2 votes)
2:52
Video transcript
- [Instructor] Let's talk
about hydrogen bonds. Depicted here, I have three
different types of molecules. On the left, I have ammonia. Each ammonia molecule
has one nitrogen bonded to three hydrogens. In the middle, I have something you're
probably very familiar with, in fact, you're made up
of it, which is water. Each oxygen is bonded to two hydrogens. And then here on the right,
I have hydrogen fluoride. Each fluorine is bonded to one hydrogen. Now, why are these types
of molecules interesting? And what does that have
to do with hydrogen bonds? And the simple answer is,
in each of these cases, you have hydrogen bonded to a much more electronegative atom. Even though these are covalent bonds, they're going to be polar covalent bonds. You are going to have a
bond dipole moment that goes from the hydrogen to the
more electronegative atom, from the hydrogen to the
more electronegative atom, from the hydrogen to the
more electronegative atom. The more electronegative atom
is going to hog the electrons. The electrons are gonna
spend more time around that. So that end of the molecule is going to have a partial negative charge. And then the ends with the hydrogens, those are gonna have
partial positive charges. Another way to think about it is, if you added these dipole moments, you would have a net dipole
for the entire molecule that would look something like that. So we are dealing with polar molecules. And the polarity comes
from both the asymmetry, and you have a very electronegative
atom bonded to hydrogen, oxygen, very electronegative
atom, bonded to hydrogen. So this end of the molecule
is partially negative. This end of the molecule or these ends of the molecule
are partially positive. For hydrogen fluoride, this
end is partially positive. This end is partially negative. And so what do you think could happen when these molecules
interact with each other? The nitrogen end right
over here, of this ammonia, could be attracted to
one of these hydrogens that has a partially positive
charge right over there. Or this hydrogen, the
partial positive charge, might be attracted to that nitrogen that has a partial negative charge. And this attraction between the partial positive hydrogen end and the partially negative
end of another molecule, those are hydrogen bonds. And they are an intermolecular
force that will be additive to the total intermolecular force from, say, things like
London dispersion forces, which makes you have
a higher boiling point than you would have if you just thought about London dispersion forces. And to make that clear,
you can look at this chart. You can see all of these
molecules are formed between period two elements and hydrogen. In fact, all of these molecules
have similar molar masses, methane, ammonia, hydrogen
fluoride, and water. If we were just thinking about
London dispersion forces, London dispersion forces are proportional to the polarizability of a molecule, which is proportional to
the electron cloud size, which is proportional to the molar mass. And generally speaking, as
you go from molecules formed with period two elements
to period three elements to period four elements
to period five elements, you do see that as the molar mass of those molecules increase, there is that general upward
trend of the boiling point, and that's due to the
London dispersion forces. But for any given period,
you do see the separation. And in particular, you see a lot of separation
for the molecules formed with oxygen, fluorine, and nitrogen. These molecules, despite
having similar molar masses, have very different boiling points. So there must be some other
type of intermolecular forces at play above and beyond
London dispersion forces. And the simple answer is yes. What you have at play
are the hydrogen bonds. Now, some of you might be wondering, well, look at these molecules formed with period three elements and hydrogen or period four elements and hydrogen, they also don't have
the same boiling point, even though you would expect similar London dispersion forces because they have similar molar masses. And the separation that you
see here in boiling points, this, too, would be due to other things, other than London dispersion forces. In particular, dipole-dipole
forces would be at play. But what you can see is
the spread is much higher for these molecules formed
with nitrogen and hydrogen, fluorine and hydrogen,
and oxygen and hydrogen. And that's because hydrogen
bonds can be viewed as the strongest form
of dipole-dipole forces. Hydrogen bonds are a special
case of dipole-dipole forces. When we're talking about hydrogen bonds, we're usually talking about
a specific bond dipole, the bond between hydrogen and
a more electronegative atom like nitrogen, oxygen, and fluorine. And so we're specifically talking about that part of the molecule, that hydrogen part that has
a partially positive charge being attracted to the
partially negative end of another molecule. So it's really about a bond
dipole with hydrogen bonds versus a total molecular dipole when we talk about dipole-dipole
interactions in general. And so you could imagine, it doesn't even just
have to be hydrogen bonds between a like molecule. You could have hydrogen bonds
between an ammonia molecule and a water molecule or
between a water molecule and a hydrogen fluoride molecule. And I mentioned that these are
really important in biology. This right over here is a closeup of DNA. You can see that the base pairs in DNA, you can imagine the rungs of the ladder, those are formed by hydrogen
bonds between base pairs. So those hydrogen bonds are strong enough to keep that double helix together, but then they're not so strong that they can't be pulled apart when it's time to replicate
or transcribe the DNA. Hydrogen bonds are also
a big deal in proteins. You learn in biology class
that proteins are made up of chains of amino acids, and the function is heavily influenced by the shape of that protein. And that shape is
influenced by hydrogen bonds that might form between the amino acids that make up the protein. So hydrogen bonds are everywhere. There are many hydrogen bonds
in your body right now mainly, not just because of the DNA, mainly because you're mostly water. So life, as we know it, would not exist without hydrogen bonds.