Hydrogen Bonding is the development of hydrogen bonds, which are a type of attractive intermolecular force caused by the dipole-dipole interaction between a hydrogen atom bonded to a strongly electronegative atom and another highly electronegative atom nearby.

You may be surprised to learn that hydrogen bonds hold our body’s basic structure, which contains the genetic information-DNA. The availability of ice in the solid form is also explained by the hydrogen bond. Have you ever wondered why water exists as a liquid? Why does boiling require a temperature of \(100\) degrees Celsius? Why is ethyl alcohol water-soluble yet ethane isn’t? Let’s find out everything about this in detail in this article.

What is Hydrogen Bonding?

Hydrogen bonding is primarily the electrostatic force of attraction due to the dipole-dipole interaction between a hydrogen atom covalently bonded to a highly electronegative atom or group and another highly electronegative bearing a lone pair of electrons that lies in the vicinity of the hydrogen atom.

For example, in water molecules \(\left({{{\text{H}}_2}{\text{O}}} \right),\) hydrogen is covalently bonded to the oxygen atom. As oxygen is highly electronegative, the shared pair of electrons between hydrogen and oxygen are attracted more towards the oxygen atom resulting in the formation of dipoles. It is due to the dipole interactions that give rise to hydrogen bonding in water molecules. The hydrogen atom of one water molecule interacts with the oxygen atom of another \({{\text{H}}_2}{\text{O}}\) molecule.

As there is a large difference in the electronegativities of oxygen and hydrogen, the bonding pair of electrons lie very close to the oxygen atom. This leads to the formation of a partial negative charge \(\left( { – \delta } \right)\) on the oxygen atom and a partial positive charge \(\left({ + \delta } \right)\) on the hydrogen atom. It is due to the electrostatic attraction between the hydrogen atom of one water molecule (with \( + \delta \) charge) and the oxygen atom of another water molecule (with \( – \delta \) charge) that leads to the formation of a hydrogen bond. These attractive intermolecular forces arise only in compounds featuring hydrogen atoms bonded to a highly electronegative atom. Hydrogen bonds are primarily strong in comparison to normal dipole-dipole and dispersion forces. However, they are weak compared to actual covalent or ionic bonds.

Conditions for Hydrogen Bonding

A hydrogen bond is formed only in polar molecules. These polar molecules should contain a hydrogen atom covalently bonded to a strongly electronegative element. Due to the presence of a strongly electronegative element, the shared pair of electrons are attracted more towards the electronegative element.

This results in the formation of a dipole in which one end becomes slightly negative while the other end becomes slightly positive. Hydrogen bond is the result of interaction between the negative end of one molecule and the positive end of the other.

As a result of hydrogen bonding, a hydrogen atom links the two electronegative atoms simultaneously, one by a covalent bond and the other by a hydrogen bond. The conditions necessary for the formation of hydrogen bonding are:

1. Presence of a Highly Electronegative Atom

The compounds containing a hydrogen bond must have a strongly electronegative atom covalently bonded to a hydrogen atom. For example, fluorine bonded to hydrogen as \({\text{HF}}\), oxygen with hydrogen as in water, or nitrogen as in ammonia.

The electronegativity of \({\text{H,F,}}\) and \({\text{N}}\) are in the order \({\text{F}} > {\text{O}} > {\text{N}}\). The higher the electronegativity the more is the polarization of the molecule.

2. Presence of Small Size of an Atom

The size of the electronegative atom should be small. The smaller the size, the greater is the electrostatic force of attraction, higher is the hydrogen bonding.

Consider the molecule of water \(\left({{{\text{H}}_2}{\text{O}}}  \right)\) and hydrogen sulphide  \(\left({{{\text{H}}_2}{\text{S}}} \right).\) The atomic radius of the oxygen atom is \(0.73\,{\text{A}}\), and the atomic radius of sulfur is \(1.02\,{\text{A}}.\) The oxygen atom is smaller in size, with 8 electrons arranged in two energy orbits around the positive nucleus.

The bonding pairs between oxygen and hydrogen are closer to the nucleus of the oxygen atom. Hence, the attractive force to pull the shared pairs of electrons towards the nucleus is more. The bond becomes polar, making the water molecule a dipole. This helps in the formation of hydrogen bonding between the water molecules.

The sulfur atoms in hydrogen sulfide are bigger, with \(16\) outermost electrons arranged in three orbits around its positive nucleus. The bonding pairs between sulfur and hydrogen are away from the nucleus of the sulfur atom. Hence, the attractive force of the nucleus is less in sulfur. This limits the attraction of shared pairs towards the nucleus of sulfur.

No polarity is thus developed within the molecule of hydrogen sulfide. Hence, hydrogen bonding is not formed between the molecules of hydrogen sulfide. The absence of this bonding is responsible for the gaseous state of hydrogen sulfide at room temperature. Whereas the presence of this bonding accounts for the liquid state of water at room temperature.

Effects of Hydrogen Bonding on Elements

The effects of hydrogen bonding on elements are:

Association

The molecular masses of carboxylic acids are found to be double those calculated from their simple formula. This happens due to the presence of hydrogen bonding which makes these molecules exist as dimers rather than discrete molecules. The presence of these hydrogen bonds in alcohols causes the molecules to exist in associated form rather than as discrete molecules. Therefore, a large amount of energy is required to break these bonds. This makes alcohol boil at sufficiently high temperatures.

Dissociation

The molar mass of methane, \({\text{C}}{{\text{H}}_4}\), is \(16\) and a boiling point of \( – {164^ \circ }{\text{C}}\). Whereas water, with a molar mass of \(18\), has a boiling point of \( + {100^ \circ }{\text{C}}\). Although these two compounds have similar molar masses, a significant amount of energy must make the water molecule move into the gas phase. This is in contrast to the non-polar methane molecule. The extra energy is required to break down the hydrogen bonding network present in the water molecule.

Due to the presence of hydrogen bonding in \({\text{HF}}\), it readily dissociates and gives the difluoride ion instead of the fluoride ion in an aqueous solution. Compounds like \( {\text{KHC}}{{\text{l}}_2},\,{\text{KHB}}{{\text{r}}_2},{\text{KH}}{{\text{l}}_{\text{2}}}\) do not exist because \( {\text{Cl,}}\,{\text{Br,}}\) and \({\rm{I}}\) do not form hydrogen bonds.

Why Do Compounds Having Hydrogen Bonding Have High Melting and Boiling Points?

We know that to boil alcohol, the intermolecular forces of attraction responsible for keeping the alcohols in a liquid state must be overcome by supplying a considerable amount of heat. This explains why alcohol boils at high temperatures.

The hydroxyl group present in alcohol is made up of a hydrogen atom and an oxygen atom which is more electronegative of the two. This paves the way for the polar bond between the two. The polarity of this bond eventually builds up a force of attraction between the slightly positive hydrogen atom of one alcohol molecule with the slightly negative oxygen atom of another alcohol molecule.

This force of attraction results in the formation of the intermolecular hydrogen bond. It is said to be intermolecular because the hydrogen bond exists between two different molecules. The number of hydroxyl groups also affects the boiling points of alcohols. The more hydroxyl groups, the higher the boiling point.

1. The high melting and boiling point compounds contain hydrogen bonds because some extra energy is needed to break these bonds.
2. Hydrogen fluoride has an abnormal high boiling point compared to other halogen acids. This happens due to the existence of hydrogen bonding in the \( {\text{H-F}}\) molecule.
3. The presence of hydrogen bonding accounts for the liquid state of water \( { {\text{H}}_2}{\text{O}}.\) On the contrary, due to the absence of hydrogen bonding, \( {{\text{H}}_2}{\text{S,}}\,{{\text{H}}_2}{\text{Se}},\) and \( { {\text{H}}_2}{\text{Te}}\) are all gases at ordinary temperatures. In water, hydrogen bonding causes linkages which makes water boil at a relatively higher temperature.
4. Ammonia has a higher boiling point than \( {\text{P}}{{\text{H}}_3}\). This is because there is hydrogen bonding in \( {\text{N}}{{\text{H}}_3}\) but not in \( {\text{P}}{{\text{H}}_3}.\)
5. Ethanol, compared to its isomer diethyl ether, has a higher boiling point because hydrogen bonding is in the ethanol molecule.

Examples of Hydrogen Bonding

Hydrogen bonds are formed among the polar covalent compounds in which hydrogen is covalently bonded to an electronegative atom. The shared pair of electrons being more inclined towards the electronegative element results in the formation of a dipole. This dipole develops a partial negative charge over the electronegative atom and a partial positive charge over the hydrogen atom.

Hydrogen Bonding in Hydrogen Fluoride

Both hydrogen and fluorine need one electron each to achieve their stable electronic configuration of noble gases in the formation of the HF molecule. Hydrogen and fluorine each share one electron to form a covalent bond. But, fluorine is a highly electronegative atom, and hence it attracts a shared pair of electrons towards itself. As a result, hydrogen develops a partial positive charge \(\left( { + \delta } \right)\) while fluorine develops a partial negative charge \(\left( { – \delta } \right).\) The molecule of \( {\text{HF}}\) thus becomes polar and behaves as a dipole.

Suppose another \( {\text{HF}}\) molecule or any other polar molecule comes near the \( {\text{HF}}\) molecule. The partially positive hydrogen of one molecule will be attracted towards the partially negative atom of the approaching molecule by a weak electrostatic force. This weak electrostatic force of attraction is called a hydrogen bond. It is represented by placing a dotted line between the atoms. Fluorine having the highest value of electronegativity forms the strongest hydrogen bond.

Hydrogen Bonding in Water

The water molecule is comprised of highly electronegative oxygen covalently bonded to the hydrogen atom. The oxygen atom being electronegative attracts the shared pair of electrons more towards itself. This unequal sharing of electrons makes the oxygen end of the water molecule slightly negative, whereas the hydrogen end becomes partially positive.

Hydrogen Bonding in Ammonia

Hydrogen bonding in ammonia contains highly electronegative atom nitrogen linked to hydrogen atoms. The extent of hydrogen bonding in ammonia is limited. This is because each nitrogen only has one lone pair of electrons, which means each ammonia molecule can form one hydrogen bond using its lone pair and the other involving one of its \(\delta + \) hydrogens. The other hydrogens are wasted.

Hydrogen Bonding in Alcohols

Alcohols are the group of organic compounds that contain an \( – {\text{OH}}\) group as the functional group. Alcohols are capable of hydrogen bonding because they have a hydrogen atom attached directly to an oxygen atom. Such molecules will always have higher boiling points than similarly sized molecules which do not have an \( – {\text{O-H}}\) group.

Hydrogen Bonding in Carboxylic Acid

Carboxylic acids are the group of organic compounds which contain the carboxyl group \(\left({{\text{COOH}}} \right)\) as the functional group. The carboxyl group consists of a carbonyl group \(\left({{\text{C=O}}} \right)\) and a hydroxyl group \(\left({{\text{O-H}}} \right)\) attached to the same carbon atom.

Carboxylic acids act as both hydrogen bond acceptors and donors. Due to the presence of the carbonyl group, it acts as a hydrogen bond acceptor. The presence of the hydroxyl group makes it hydrogen bond donors. This helps them to participate in hydrogen bonding. Carboxylic acids usually exist as dimeric pairs in non-polar media because of their tendency to “self-associate.” This tendency to hydrogen bond gives them increased stability as well as higher boiling points relative to the acid in an aqueous solution. Carboxylic acids are polar molecules.

Hydrogen Bonding in Polymers

The double-helical structure of DNA consists of hydrogen bonding between its base pairs. Replication of DNA strands is possible due to the presence of a hydrogen bond. This hydrogen bond links one complementary strand to the other and enables weak hydrogen bond strength replication.

Cellulose and its derived polymers such as cotton and flax also contain hydrogen bonds.

Nylon, a synthetic polymer, has repeated hydrogen bonds. This bond plays a major role in the crystallisation of the material.

Strength of Hydrogen Bond

The strength of the hydrogen bond depends on the electronegativity of the combining atoms. The electronegativity of \( {\text{H}},\,{\text{F}}\) and \( {\text{N}}\) are in the order \( {\text{F}} > {\text{O}} > {\text{N}}\). Hence, the strength of \({\text{N}}{{\text{H}}_3},\,{\text{HF}}\) and \( {{\text{H}}_2}{\text{O}}\) is in the order \({\text{HF > }}{{\text{H}}_2}{\text{O}} > {\text{N}}{{\text{H}}_3}\). The hydrogen bond is weak. The strength of the hydrogen bond is more than the weak van der Waals forces and the strong covalent bonds.

Properties of Hydrogen Bonding

The hydrogen bond is comparatively weaker than the covalent and ionic bond.

1. Solubility: Due to the hydrogen bonding, which the alcohols form with water molecules, the lower alcohols are soluble in water.
2. Volatility: The compounds having hydrogen bonding have high boiling points. As they have high boiling points, the compounds are less volatile in nature.
3. Viscosity and surface tension: The compounds which contain hydrogen bonding exists as an associated molecule and not as a discrete molecule. So their flow becomes comparatively difficult. They have higher viscosity and high surface tension.
4. The lower density of ice than water: The presence of hydrogen bonding in water accounts for the lower density of ice compared to that of water. The hydrogen bonding in solid ice gives rise to a cage-like structure of water molecules. Each water molecule in ice is linked to four other water molecules in a tetrahedral structure. In the solid state of ice, the molecules are not as closely packed as they are in a liquid state. The cage-like structure of ice collapses, when it melts and the molecules come closer to each other. Thus for the same mass of water, the volume decreases and density increases. Therefore, ice has a lower density than water at \(273\,{\text{k}}\). That is why ice floats.

Types of Hydrogen Bonding

The two types of hydrogen bonding are classified as follows-

1. Intermolecular Hydrogen Bonding
2. Intramolecular Hydrogen Bonding

Intermolecular Hydrogen Bonding

These are the hydrogen bonds formed between several molecules of the same substance, for example-water \(\left({{{\text{H}}_2}{\text{O}}} \right)\). These are also formed between several molecules of different substances. For example- Water and acetone. When the formation of a hydrogen bond takes place between different molecules of the same or different compounds, it is called intermolecular hydrogen bonding. For example – hydrogen bonding in water, alcohol, etc.

Intramolecular Hydrogen Bonding

The hydrogen bonds formed between two atoms of the same molecule are called intramolecular hydrogen bonding. When the hydrogen atoms of one group link with the more electronegative atom of the other group within the same molecule, the formation of an intramolecular hydrogen bond take place.

Symmetric Hydrogen Bond

Symmetric Hydrogen bond is a special type of hydrogen bond in which the proton is placed exactly halfway between two identical atoms. It is a three-centre four-electron type of bond. This bond has a strength almost similar to a covalent bond and is also much stronger compared to the “normal” hydrogen bond.

Strength of Hydrogen Bonding

The most important hydrogen bond occurs between hydrogen and highly electronegative atoms. The length of a chemical bond depends on its strength, pressure and temperature. The bond angle depends on the specific chemical species involved in the bond. The strength of hydrogen bonds ranges from very weak (1–2 kJ mol) to very strong (161.5 kJ mol m 1).

Hydrogen Bonding and Water

Hydrogen bonds account for some important properties of water. Even though a hydrogen bond is only 5% as strong as a covalent bond, it is sufficient to stabilise water molecules.

• Hydrogen bonding causes water to remain liquid over a wide temperature range.
• Because it takes additional energy to break hydrogen bonds, water has an unusually high heat of vaporisation. Water has a much higher boiling point than other hydrides.

The effects of hydrogen bonds between water molecules have several important consequences:

• Hydrogen bonding makes ice less dense than liquid water, so ice floats on water.
• The effect of hydrogen bonding helps make sweating an effective means of heat evaporation to lower temperatures for animals.
• The effect on heat capacity means that water protects against extreme temperature changes near large bodies of water or moist environments. Water helps to regulate temperature on a global scale.

Summary

The entire human body is made up of hydrogen bonds, from water comprising \(70\% \) of our body to the DNA responsible for replicating our genes. It is the presence of a hydrogen bond that makes ice float on water. From alcohols, plastics to polymers, all are in some or another way, have hydrogen bonds present.

FAQs

The frequently asked questions on hydrogen bonding are given below:

Q.1. Why does hydrogen bonding increase boiling point?
Ans: The greater the attractions, the more energy is needed, and hence higher will be the boiling point. In water, because of the hydrogen bonding attraction between molecules, greater energy is needed to separate them from their intermolecular attraction; therefore, the higher boiling point

Q.2. Which attractive force is the weakest?
Ans: The London dispersion force is the weakest among all intermolecular forces of attraction. The London dispersion force is a temporary attractive force that results when the electrons in two adjacent atoms occupy positions that make the atoms form temporary dipoles. This force is sometimes called an induced dipole-induced dipole attraction.

Q.3. Which compounds will not form hydrogen bonding?
Ans: Non-polar covalent compounds are not miscible with each other and do not form hydrogen bonds, whereas the others are polar and form \({\text{H}}\)-bonds

Q.4. What are the applications of hydrogen bonding?
Ans: Hydrogen bond has a wide range of applications. It occurs in inorganic molecules, such as water, and organic molecules, such as DNA and proteins. The two complementary strands of DNA are held together by hydrogen bonds between complementary nucleotides (A&T, C&G).

We hope that this detailed article on Hydrogen Bonding proves helpful to you. If you have any queries about hydrogen bonding or in general about this article, ping us through the comment box below and we will get back to you as soon as possible.