Hydrogen Bonding

Hydrogen Bonding

Hydrogen bonding is a special type of dipole-dipole interaction that exists between a hydrogen atom covalently bonded to a highly electronegative atom (like Nitrogen, Oxygen, or Fluorine) and another highly electronegative atom located nearby. It is a strong intermolecular force, significantly stronger than typical dipole-dipole forces but weaker than covalent or ionic bonds.

Hydrogen bonding plays a crucial role in determining the physical properties of many substances, most notably water, and is fundamental to the structure and function of biomolecules like DNA and proteins.

Cause Of Formation Of Hydrogen Bond

The formation of a hydrogen bond is a direct consequence of the significant electronegativity difference between hydrogen and certain highly electronegative atoms (N, O, F).

1. Polarity of the H-X Bond:

• When hydrogen is covalently bonded to a highly electronegative atom like Nitrogen (N), Oxygen (O), or Fluorine (F), the shared electrons in the covalent bond are pulled more strongly towards the electronegative atom.
• This creates a highly polar bond, resulting in a significant partial positive charge ($\delta^{+}$) on the hydrogen atom and a significant partial negative charge ($\delta^{-}$) on the electronegative atom (X, where X = N, O, or F).
$$ \stackrel{\delta+}{\text{H}} - \stackrel{\delta-}{\text{X}} $$
• The hydrogen atom, with its small size and concentrated positive charge, becomes a strong electrophilic center.

2. Lone Pair on Another Electronegative Atom:

• For a hydrogen bond to form, there must be another highly electronegative atom (usually N, O, or F) in a nearby molecule (or in the same large molecule) that possesses at least one lone pair of electrons.
• This lone pair acts as a Lewis base and is attracted to the partially positive hydrogen atom.

3. The Hydrogen Bond Interaction:

• The partially positive hydrogen atom ($\delta^{+}$) of one molecule is attracted electrostatically to the lone pair of electrons ($\delta^{-}$) on the highly electronegative atom of another molecule.
• This attraction constitutes the hydrogen bond. It is often represented by a dotted or dashed line:
$$ \stackrel{\delta+}{\text{H}}^{\text{a}} - \stackrel{\delta-}{\text{X}}^{\text{a}} \quad \cdots \quad \stackrel{\delta+}{\text{H}}^{\text{b}} - \stackrel{\delta-}{\text{Y}}^{\text{b}} $$ Here, H^a is bonded to X, and the hydrogen bond is formed between H^a and Y (where Y is usually N, O, or F). X and Y are electronegative atoms with lone pairs.

Strength of Hydrogen Bond:

• The strength of the hydrogen bond depends on the electronegativity of the atom X and Y, and the distance between them.
• The electronegativity order for hydrogen bonding is F > O > N. Thus, HF molecules exhibit stronger hydrogen bonds than $$H_2O$$ molecules, which in turn have stronger hydrogen bonds than $$NH_3$$ molecules.

Types Of H-Bonds

Hydrogen bonds are broadly classified into two main types based on their strength and the involvement of electronegative atoms:

1. Intermolecular Hydrogen Bonds:

• These bonds occur between molecules.
• They are responsible for many of the unusual physical properties of substances like water.
• Examples:
  • In Water ($$H_2O$$): Each water molecule can form up to four hydrogen bonds with its neighbors. The partially positive hydrogen of one water molecule is attracted to the lone pair on the oxygen of another water molecule. This extensive hydrogen bonding network explains water's high boiling point, high surface tension, and its ability to act as a solvent.
  • In Ammonia ($$NH_3$$): Nitrogen is electronegative and has a lone pair, and hydrogen bonded to it is partially positive. Ammonia molecules can form intermolecular hydrogen bonds, though they are weaker than those in water due to nitrogen's lower electronegativity compared to oxygen.
  • In Hydrogen Fluoride (HF): Fluorine is the most electronegative element, and the H-F bond is highly polar. HF molecules form strong intermolecular hydrogen bonds, leading to a high boiling point for HF compared to other hydrogen halides like HCl.
  • In Alcohols (R-OH): The presence of the -OH group allows alcohols to form intermolecular hydrogen bonds.

2. Intramolecular Hydrogen Bonds:

• These bonds occur within the same molecule.
• They are generally weaker than intermolecular hydrogen bonds because the distance between the participating electronegative atoms is usually larger, and the interaction involves different parts of the same molecule.
• Intramolecular hydrogen bonds can occur when a molecule contains both a hydrogen atom bonded to a highly electronegative atom (like -OH, -NH, -FH) and another highly electronegative atom with a lone pair, positioned such that a ring structure can be formed.
• Examples:
  • Ortho-nitrophenol ($$o-NO_2C_6H_4OH$$): The hydrogen atom of the hydroxyl (-OH) group forms a hydrogen bond with the oxygen atom of the nitro (-NO₂) group, which is located ortho to the -OH group. This forms a six-membered ring.
  • In Biomolecules: Intramolecular hydrogen bonds are crucial for maintaining the secondary and tertiary structures of proteins (e.g., alpha-helices and beta-pleated sheets) and the double helix structure of DNA.

Comparison of Strengths:

The strength of hydrogen bonds can be categorized:

• Strong H-bonds: Typically formed between F atoms and H atoms bonded to F or O (e.g., in HF, $$H_2O$$). These are often considered to have some degree of covalent character.
• Moderately strong H-bonds: Formed between O or N atoms and H atoms bonded to O or N (e.g., in $$H_2O, NH_3, alcohols, amines$$).
• Weak H-bonds: Formed when the electronegative atom is less electronegative, like Chlorine (Cl) or Sulfur (S), or when hydrogen is bonded to Carbon in specific contexts (e.g., in the C-H bond of chloroform, $$CHCl_3$$). These are often referred to as weak hydrogen bonds or sometimes as dipole-dipole interactions.

Hydrogen bonding significantly impacts properties like boiling point, melting point, viscosity, solubility, and the structure of complex molecules.