Lecture 9: Hydrogen Bonding

Hydrogen's placement in the periodic table is unique due to its position in period 1, possessing only one electron and one proton. This unique configuration influences its bonding behavior and ability to participate in hydrogen bonds.
The boiling points of hydrides indicate a special type of intermolecular bonding for oxygen, nitrogen, and fluorine hydrides, known as hydrogen bonding. This is essential for understanding the unusual properties of water, ammonia, and hydrogen fluoride compared to other hydrides.

A hydrogen bond is a special weak attractive force that exists through space between a hydrogen atom and the lone pair electrons on atoms like oxygen and nitrogen. These are highly electronegative atoms, causing a significant dipole moment in the bond with hydrogen.
It is a particularly important form of bonding in biology and earth sciences, influencing protein structure, DNA base pairing, and numerous other biological processes. For example, hydrogen bonds stabilize the double helix structure of DNA.
An intermolecular hydrogen bond exists between two separate molecules. Example: H-O-H…O-H where the dots represent the hydrogen bond between the molecules. This type of bonding explains water's high surface tension, cohesion, and adhesion properties.
An intramolecular hydrogen bond exists between two parts of the same molecule, forming a ring-like structure within the molecule itself. This can stabilize the conformation of large biomolecules such as proteins and contribute to their unique three-dimensional structures.

A hydrogen bond forms when a hydrogen atom covalently bonded to an appropriate donor atom (N, O) is weakly connected through space to one or more acceptor atoms (N, O, F). Fluorine is included as it is also highly electronegative and can act as a strong hydrogen bond acceptor.
The donor atom must act as an electron-withdrawing group to polarize the hydrogen atom, creating a dipole. This polarization is essential for the hydrogen atom to have a partial positive charge and be attracted to the acceptor atom.
The acceptor atom must have an available lone pair of electrons to interact with the partially positive hydrogen. These lone pairs provide a region of high electron density for the hydrogen bond to form.
Representation: X-H…Y, where X is the donor, H is the hydrogen atom, Y is the acceptor, the line between X and H represents a strong covalent bond, and the dots between H and Y represent the hydrogen bond. This notation helps to visualize the hydrogen bond and the atoms involved.

The strength of a hydrogen bond is weak compared to covalent and ionic bonds, but significant compared to other intermolecular forces like Van der Waals forces. Hydrogen bonds are stronger than typical dipole-dipole interactions due to the high polarity of the X-H bond.
Bond energies for a hydrogen bond are approximately 5 - 25 kJ/mol, while for ionic or covalent bonds, they are approximately 200 - 1000 kJ/mol. This difference dictates the ease with which these bonds can be broken or formed, influencing reaction kinetics and physical properties.
Remembering 15 kJ/mol is a good approximation for hydrogen bonding energy. This average value can be used for estimations in various calculations and comparisons.
Bond energy (B.E.) is the amount of energy required to dissociate (break) the bond and dissociate the molecule into atoms. It's a measure of bond stability, with higher bond energies indicating stronger bonds.
Example: For H_2(g) -> 2H(g), the B.E.(H-H) = 436 kJ/mol. This indicates the energy needed to break the bond in a hydrogen molecule, illustrating the strength of a covalent bond.

The strength of a hydrogen bond is influenced by three factors:

(i) A stronger H-bond results from a more electronegative donor atom (X) and a more electron-rich acceptor atom (A). Greater electronegativity differences lead to stronger bonds due to increased polarization of the X-H bond.
(ii) The H---A bond length (1.1 - 1.9 Å). This parameter responds to factor (i) and steric constraints. Shorter distances generally indicate stronger bonds because the electrostatic attraction is greater at shorter distances, but steric hindrance can affect this.
(iii) The X-H---A bond angle should ideally be close to 180^o. Bond strength decreases with decreasing bond angle. The angle is also influenced by factor (i) and steric constraints. Deviations from linearity weaken the bond due to reduced orbital overlap and electrostatic interactions.

Example of a weak H-bond: H…F ilda 1.8 Å, NHF ilda 120^o. Note: 1Å = 1 x 10^{-10}m. This example illustrates how deviations from optimal conditions weaken the bond. The longer bond length and non-linear angle result in a weaker interaction.

A very polar covalent bond exists between H and an electronegative atom with lone pairs, such as O, F, N. This polarity is crucial for H-bonding because it creates partial charges on the atoms involved.
The polar covalent bond results from large differences in electronegativity. For example: F (4.0) and H (2.1). The substantial difference creates a strong dipole moment, with fluorine carrying a partial negative charge and hydrogen carrying a partial positive charge.
The small size of H means it has a high positive charge density, allowing close approach and strong interaction with lone pairs. This high charge density enhances the electrostatic attraction between the hydrogen and the acceptor atom.
The positive H of one molecule is attracted to the partially negative lone pair on the next molecule, resulting in an intermolecular H-bond. The attraction is electrostatic in nature, similar to the attraction between oppositely charged ions, but weaker.

Comparing the physical properties of water (which exhibits H-bonding) to methane (with a similar molecular weight but no H-bonding) reveals significant differences. Water's properties are dramatically affected by H-bonding, making it essential for life.
Melting point, boiling point, heat of vaporization, and heat of fusion are all significantly increased in water due to H-bonding holding the H_2O molecules together. These properties require more energy to overcome the intermolecular forces, resulting in higher values compared to similar-sized molecules without H-bonding.

Water is considered the environmental and biological solvent because of its ability to form hydrogen bonds with many substances. This allows it to dissolve a wide