Have you ever jumped into a pool and hit the water with your belly or back? Your skin ends up stinging a bit. The effect is even worse if you jumped from high up on a diving board. So why does jumping into the water hurt?
That is because of hydrogen bonding.
In this article, we will be learning about what hydrogen bonds are and why they make water so strong (and painful for some unsuspecting jumpers).
- This article is about hydrogen bonds.
- First, we will define what a hydrogen bond is
- Next, we will explain why and how hydrogen bonds form
- Then, we will look at the two types of hydrogen bonds: intermolecular and intramolecular
- Lastly, we will look at how hydrogen bonding affects different properties of molecules and learn why hydrogen bonds are so important
Hydrogen Bond Definition
Let's start by looking at the definition of hydrogen bonds.
Hydrogen bonding is the attractive force between the lone pair(s) of a very electronegative atom (normally oxygen, nitrogen, or fluorine) and hydrogen when it is bonded to another very electronegative atom (again, usually O, N, or F)
Electronegativity is the tendency for an atom to pull electrons/electron density towards itself. The closer an atom is to fluorine (top right) on the periodic table, the more electronegative it is and the closer it is to francium (bottom left), the less electronegative it is.
"Hydrogen bonding" is a bit of a misnomer, since it is mostly an attraction, not a bond. More specifically, it is an electrostatic (due to electrons/charges) attraction. However, there is some covalency happening (charge transfer through orbital overlap), but it is significantly weaker than a fully covalent bond.
Hydrogen Bond Formation
Hydrogen bonding is able to occur because of a bond dipole.
A bond dipole is a separation of opposite charges across a bond.
When hydrogen is bonded to N, O, or F, there is a significant difference in electronegativity between hydrogen and the other atom. Because of this, electron density is being pulled towards that atom, leaving hydrogen with a lack of electron density. This causes a dipole.
In this dipole, the other atom has an excess of electron density, so it has a partial negative charge (δ-). Hydrogen, on the other hand, is lacking electron density, so it has a partial positive charge (δ+). This partial positive charge is the key to hydrogen bonding.
Since hydrogen now really wants some electron density, it will go after lone pairs.
Lone pairs are pairs of electrons that don't participate in bonding. These electrons are valence electrons, meaning they are in an atom's outermost shell/highest energy level.
Since the other atom's electrons are participating in bonding, the lone pairs make the perfect target for hydrogen.
Hydrogen bonding describes the attraction between these electrons and the (partial) positive hydrogen.
Here is what hydrogen bonding looks like in water:
Fig.1 Hydrogen bonding in water
Hydrogen is attracted to the lone pairs on the nearby oxygen. Oxygen contains two lone pairs, so each molecule of water can form four hydrogen "bonds".
While hydrogen bonds are one of the strongest intermolecular forces (forces between molecules), they are still pretty weak compared to actual covalent bonds (bonds where electrons are shared). Hydrogen bonds have ~1/10th the strength of a covalent bond, and are constantly broken and reformed in (liquid) water.
Hydrogen Bond Types
There are four major types of hydrogen bond:
Intermolecular hydrogen bond
Intramolecular hydrogen bond
Salt bridges
π-hydrogen bonding with aromatic compounds
Intermolecular hydrogen bonds are between molecules, like in water, as we saw earlier, whereas intramolecular hydrogen bonds are within a molecule. The other two types of hydrogen bonding (salt bridges and π-hydrogen bonding with aromatic compounds) are not as prevalent at the AP chemistry levels. So, we will touch on these topics in a bit less detail.
Hydrogen can be attracted to a highly electronegative atom as long as it is close enough and contains a lone pair. Fifgure 2 shows an example of intramolecular hydrogen bonding:
Fig.2 An example of intramolecular hydrogen bonding
The hydrogen is close enough to oxygen that it can "bond". While it isn't pictured, hydrogen is still attacking a lone pair.
The key factor here is distance. Think of it like putting 2 magnets on opposite sides of a table. When they are far away, they won't affect each other. However, once they get close enough, they will be pulled together.
As an aside, be careful when you look at structures like these. Lone pairs aren't usually shown, but you have to remember when an atom will actually have them. For example, nitrogen will have a lone pair when it has three bonds (like in NH3), but it doesn't have a lone pair (usually) when it has four bonds (like NH4), so hydrogen bonding won't occur.
Other types of Hydrogen Bonding
As I mentioned earlier, there are two other types of hydrogen bonding that we are going to touch on briefly: salt bridges and π-hydrogen bonding with aromatic compounds.
Salt bridges are a type of hydrogen bonding where the donor atom (atom hydrogen is bonded to) and the acceptor atom (what hydrogen is attracted to) are "fully charge" (i.e. ions)
The types of hydrogen bonding we saw earlier were with neutral molecule dipoles. In this case, the atoms involved (besides hydrogen) are ions. Below is an example of a salt bridge:
Fig.3-An example of a salt bridge
Here, the nitrogen has a full positive charge, while oxygen has a full negative charge. In our previous examples, the charges were only partial due to dipole formation in neutral molecules/atoms.Next, we have π-hydrogen bonding with aromatic compounds.
An aromatic ring is a carbon ring structure with alternating single and double bonds. The partial charges present on each carbon causes the ring to be a hydrogen bond acceptor (i.e. hydrogen is attracted to it).
Below is an example:
Fig.4-Hydrogen is attracted to the aromatic ring, forming a hydrogen bond
Hydrogen bonding affects different properties of molecules. One such property is the boiling point. Let's look at the different boiling points for hydrogen-group 16 compounds:
Fig. 5- Water is an outlier in the boiling point trend
Usually, a higher molecular mass means the boiling point should be higher. While the other compounds follow this trend, water is a clear outlier. This is because it is the only species here, in the graph, that has hydrogen bonding.
Molecules are held together by different intermolecular forces. The weaker the forces, the easier it is for them to be broken, so the species can become a gas. The hydrogen bonding in water is a stronger intermolecular force than those that exist for the other compounds, so water has a much higher boiling point (this is also true for melting point for the same reasons)
Another property that hydrogen bonds affect is viscosity. Viscosity is essentially the "thickness" of a liquid. When a compound has strong forces between its molecules, they are more closely held together. This means they are stronger and "thicker".
The last property we are going to talk about is surface tension. We brought up an example of surface tension in the intro when we talked about belly-flopping
Surface tension is the ability for a liquid to resist an external force (in this case, someone jumping into a pool).
The hydrogen bonding in the water keeps that surface layer strong and hard to "break" which is why it can hurt to belly flop. This is also why insects such as water striders are able to skate across the water without breaking the surface.
Hydrogen Bond Importance
Hydrogen bonding is an incredibly important force. Here are just a few ways hydrogen bonding is essential to life:
Hydrogen bonds stabilize the structures of DNA and proteins
Water can hydrogen bond to the cell walls of plants. This is why water can "climb" up the roots of plants
Antibodies use hydrogen bonding to target and "fit" into a specific species that is attacking the body
Hydrogen bonds keep the boiling point of water high, so our oceans, lakes, and seas don't rapidly boil away
Based on these few reasons alone, you can see why hydrogen bonds are so important!
Hydrogen Bonds - Key takeaways
- Hydrogen bonding is the attractive force between the lone pair(s) of a very electronegative atom (normally O, N, or F) and hydrogen when it is bonded to another very electronegative atom (again, usually O, N, or F)
- Electronegativity is the tendency for an atom to pull electrons/electron density towards itself. The closer an atom is to fluorine (top right) on the periodic table, the more electronegative it is.
- A bond dipole is a separation of opposite charges across a bond.
- Hydrogen bonds form when hydrogen is bonded to a very electronegative atom (N, O, or F) and then is attracted to the lone pair of electrons of another very electronegative atom. Hydrogen bonds can exist between molecules (intermolecular) or within molecules (intramolecular)
- Hydrogen bonding increases properties such as boiling point, viscosity, and surface tension.
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