which of the molecules listed below can form hydrogen bond? For wich of the molecules would dispersion forces be the only intermolecular force? Give reasons for answer.
A.H2
B.NH3
C.HCl
D.HF
HF is the hydrogen bond in the listed below because hydrogen must be bonded to either a fluorine,oxygen or nitrogen atom.
H bonds are formed in molecules with N, O, or F as the central atom and containing H atoms attached. So NH3, HCl, and HF are capable of forming H bonds. H2 is a homomolecule and will have only dispersion forces.
Hcl
Well, let's break it down and have some fun with it!
A. H2 - Ah, good ol' hydrogen. It's a bit of a lone ranger and can't form hydrogen bonds with itself. It doesn't have a readily available lone pair of electrons to participate in hydrogen bonding.
B. NH3 - Nitrogen gets a little fancy here. It not only has a lone pair of electrons, but it's also surrounded by those delightful hydrogen atoms. With its nitrogen-hydrogen bond, NH3 can form hydrogen bonds. Time to party with hydrogen bonding!
C. HCl - Now, hydrogen has a hot date with chlorine. While they have a polar bond, it's not strong enough to form hydrogen bonds. They just have dispersion forces. Sorry, HCl! No hydrogen bonding for you.
D. HF - Ah, the charismatic duo of hydrogen and fluorine! They may be a small pair, but they make up for it with the ability to form hydrogen bonds. Their electronegativity difference creates a strong and polar bond, allowing them to engage in the wonderful world of hydrogen bonding.
So, to summarize:
- NH3 and HF can form hydrogen bonds because they have a lone pair of electrons and are surrounded by hydrogen atoms.
- H2 and HCl can't form hydrogen bonds. H2 doesn't have a lone pair, while HCl's polarity isn't strong enough. They rely on good old dispersion forces.
Hope that brings a smile to your face!
To determine which of the molecules listed can form hydrogen bonds and for which molecules dispersion forces are the only intermolecular force, we need to understand the nature of these interactions.
Hydrogen bonding occurs when a hydrogen atom is covalently bonded to a highly electronegative atom (like nitrogen, oxygen, or fluorine), and it is attracted to a lone pair of electrons on another electronegative atom. Hydrogen bonds are stronger than dispersion forces.
On the other hand, dispersion forces (also called London forces) are the weakest intermolecular forces that exist between all molecules. These forces arise due to temporary fluctuations in electron distribution, leading to temporary dipoles.
Now let's analyze each molecule from the provided options:
A. H2: Hydrogen gas (H2) consists of two hydrogen atoms bonded by a covalent bond. It cannot form hydrogen bonds since there is no highly electronegative atom present. However, it can exhibit dispersion forces, which are the only intermolecular forces present in this case.
B. NH3: Ammonia (NH3) contains one nitrogen atom and three hydrogen atoms. The nitrogen atom is highly electronegative, and it has a lone pair of electrons. Therefore, NH3 can form hydrogen bonds. In addition to hydrogen bonding, dispersion forces can also exist between NH3 molecules.
C. HCl: Hydrogen chloride (HCl) comprises a hydrogen atom and a chlorine atom. While hydrogen bonding is not possible in this molecule, it can exhibit dispersion forces due to temporary fluctuations in electron distribution.
D. HF: Hydrogen fluoride (HF) contains a hydrogen atom bonded to a fluorine atom. Fluorine is highly electronegative, and it has a lone pair of electrons, allowing HF to form hydrogen bonds. Dispersion forces can also exist between HF molecules.
To summarize:
- Hydrogen bonds can form in NH3 and HF due to the presence of highly electronegative atoms (nitrogen and fluorine) with lone pairs of electrons.
- Dispersion forces are present in all the molecules listed (H2, NH3, HCl, and HF). However, for H2, dispersion forces are the only intermolecular forces since hydrogen bonding is not possible.
Remember, identifying hydrogen bonding and dispersion forces requires an understanding of the electronegativity and presence of lone pairs on atoms within the molecules.