Hydrogen Bonding in Chemistry

Chemistry is not limited to chemical bonds, such as ionic or covalent bonds. But there are many other forces between and within molecules that shape various aspects of chemistry. One of them is the hydrogen bond.

The hydrogen bond (aka H-bond) is an electrostatic force arising from the polar nature of the molecule. Hydrogen bonding involves a hydrogen atom attached to an electro-negative atom in a molecule, which creates partial polarization. Since the hydrogen atom is less electro-negative, it acquires a partial positive charge. And the other, more electro-negative one, gains a partial negative charge. Under the presence of any other electro-negative species bearing a lone pair of electrons, partial-positive hydrogen shows a high affinity toward the electrons of the species. Thus, resulting in a weaker electrostatic bond, which is what we call a hydrogen atom.

Hydrogen bonds should not be confused with ionic or covalent bonds. They do not result in chemical changes in substance, unlike ionic or covalent. However, they affect the physical properties of the substance, such as the boiling point. One well-known phenomenon is the elevation in the boiling point of water, which we will discuss later in this article.

Since they do not cause permanent chemical bonds, they are weaker compared to ionic and covalent bonds. But they are stronger compared to van der Waal forces.

Hydrogen bonds can exist between different molecules and within the same molecules. We will discuss this too with an example later in the article.

Explanation

From what we discussed above, it is clear that for the formation of a hydrogen bond, we require three things: hydrogen, H-bond donor, and H-bond acceptor.

Hydrogen: Hydrogen is the foremost requirement for hydrogen bonding. The tendency of hydrogen to acquire a partial positive charge when bonded to electronegative species is what makes hydrogen special.

H-bond donor: The hydrogen-bond donor is an electro-negative species, usually made of oxygen (O), nitrogen (N), or fluorine (F). When hydrogen is covalently bonding to O, N, or F, which are far more electro-negative than hydrogen, the bonding electrons are shared biasedly. And the electronic cloud is shifted more toward the electronegative atom. Thus, we see a partial positive charge on hydrogen and a partial negative charge on O, N, or F.

The diagram below explains the same.

H-bond acceptor: In order for a hydrogen bond to occur, we also need an H-bond acceptor. An acceptor has to be an electronegative species holding one or more lone pairs of electrons. The lone pair of electrons of an acceptor attracts partial-positive hydrogen, resulting in a hydrogen bond.

Usually, acceptors are chemical species containing O, N, or F. Since all three are electro-negative and possess one or more lone pairs of electrons on them.

The dotted line in the above diagram represents an H-bond. As mentioned earlier, the H-bond is weaker than covalent and ionic bonds. So, the distance between hydrogen and the acceptor is larger than hydrogen and the donor, which reflects in the previous illustration.

Examples

Hydrogen bonding is seen in simple chemical structures like water (H2O), ammonia (NH3), etc., and complex structures like proteins, DNA, and other bio-molecules.

Water

A classical example to study hydrogen bonding in water. Water consists of one central oxygen atom bearing two lone pairs of electrons. And two hydrogen atoms are covalently bonded to oxygen. The geometry of a water molecule is bent (or angular), having two hydrogens on two vertices of a tetrahedron. And the two lone pairs of electrons rest on the remaining two vertices.

A water molecule fulfills the requirements of both H-bond donor as well as H-bond acceptor. Both hydrogens are bonded to oxygen. Because of the electro-negative difference between oxygen and hydrogen, the bonding electrons get closer toward oxygen’s nucleus creating a dipole. The partial polarization of a water molecule is its intrinsic property, not the result of an external environment.

When a water molecule comes in the vicinity of another electron-enriched species (acceptor), hydrogens of a water molecule form H-bonds with it. In our example, oxygen in water acts as a donor and acceptor. Therefore, a water molecule forms hydrogen bonds with other water molecules. The below figure illustrates the same.

As we see from the previous figure, the two lone pairs of water molecules attract hydrogen from other water molecules and vice versa.

Hydrogen bonding does not change the chemical nature of water molecules but rather alters the physical behavior of water. In the previous figure, hydrogen bonding restricts the free movement of water molecules and brings individual molecules closer. It also explains why water exists as a liquid at room temperature and its boiling point of 100°C, which is relatively high in comparison to other molecules with similar molecular weight.

You can think of hydrogen bonding in the water as a strong inter-molecular force, which holds the individual molecules of water together.

In a water molecule, we have two lone pairs of electrons and two hydrogens; each of these is utilized in hydrogen bonding. Thus, a single water molecule can make up to four H-bonds.

Solvation (hydration)

The polar nature of water makes it an excellent solvent for ionic compounds and partial-polar covalent compounds. For example, when NaCl is dissolved in water. Sodium cations (Na+) attract oxygen atoms of water molecules, while hydrogen atoms cluster chloride anions (Cl−). The energy exchange during the hydration is called hydration energy.

Ammonia

Ammonia has a central nitrogen atom, three hydrogen atoms covalently bonded to nitrogen, and a lone pair of electrons. Hydrogen bonding in ammonia is illustrated below:

In ammonia, the number of hydrogens exceeds the number of a lone pair of electrons. Thus, hydrogen bonding in ammonia is not as effective as in water since the shortage of lone pairs cannot satisfy all hydrogen atoms.

Alcohols

Alcohols (R–OH) are defined by the presence of a hydroxyl group (–OH) attached to an alkyl group (–R). This hydroxyl group is similar to –OH of water. The oxygen of hydroxy is more electro-negative than hydrogen and creates the –OH group suitable for hydrogen bonding.

The effects of hydrogen bonding in alcohols is pronounced in their boiling and melting points and also in their water solubility.

The table below displays the boiling point of alcohols and their closest mate hydrocarbons.

Alcohol	Mol Wt	BP (°C)	Alkane	Mol Wt	BP (°C)	Ether	Mol Wt	BP (°C)
Ethanol	46	78	Ethane	30	−89	Methoxymethane	46	−24
Propan-1-ol	60	97	Propane	44	−42	Methoxyethane	60	7
Butan-1-ol	74	118	Butane	58	−1	Methoxypropane	74	39

From the table, it is clear alcohols have higher boiling points than equal-carbon alkanes and ethers.

o-nitrophenol

Intramolecular hydrogen bonding happens in higher molecules. One such example is ortho-nitrophenol.

In nitrophenol, we have a nitro group (–NO2) and hydroxyl group (–OH). When the nitro group is located in an ortho position, we see the hydrogen bonding between both functional groups.

In this case, hydrogen bonding makes o-nitrophenol more volatile than p-nitrophenol, which experiences intermolecular hydrogen bonding.

DNA

DNA is a pretty complex and enthralling chemical structure that is responsible for passing genetic information from one generation to another.

The two strands of DNA are connected with the help of hydrogen bonding. Nitrogen bases on one strand connect nitrogen bases on the other strand with H-bonds.

Image source: Boumphreyfr/ CC-BY-SA-3.0