A hydrogen bond (often informally abbreviated H -bond) is a primarily electrostatic force of attraction between a hydrogen (H) atom which is covalently bound to a more electronegative atom or group, particularly the second-row elements nitrogen (N), oxygen (O), or fluorine (F)—the hydrogen bond donor (Dn)—and another electronegative atom bearing a lone pair of electrons—the hydrogen bond acceptor (Ac). Such an interacting system is generally denoted Dn–H···Ac, where the solid line denotes a fully covalent bond, and the dotted or dashed line indicates the hydrogen bond. The use of three centered dots for the hydrogen bond is specifically recommended by the IUPAC.[4] There is general agreement that there is actually a minor covalent component to hydrogen bonding, especially for moderate to strong hydrogen bonds (> 5 kcal/mol), although the importance of covalency in hydrogen bonding is debated. At the opposite end of the scale, there is no clear boundary between a weak hydrogen bond and a van der Waals (e.g., dipole-dipole) interaction.
Hydrogen bonds can be intermolecular (occurring between separate molecules) or intramolecular (occurring among parts of the same molecule).[5][6][7][8] Depending on the nature of the donor and acceptor atoms which constitute the bond, their geometry, and environment, the energy of a hydrogen bond can vary between 1 and 40 kcal/mol.[9] This makes them somewhat stronger than a van der Waals interaction, and weaker than fully covalent or ionic bonds. This type of bond can occur in inorganic molecules such as water and in organic molecules like DNA and proteins.
The hydrogen bond is responsible for many of the anomalous physical and chemical properties of compounds of N, O, and F. In particular, intermolecular hydrogen bonding is responsible for the high boiling point of water (100 °C) compared to the other group 16 hydrides that have much weaker hydrogen bonds.[10] Intramolecular hydrogen bonding is partly responsible for the secondary and tertiary structures of proteins and nucleic acids. It also plays an important role in the structure of polymers, both synthetic and natural.
Weaker hydrogen bonds[11] are known for hydrogen atoms bound to elements such as sulfur (S) or chlorine (Cl); even carbon (C) can serve as a donor, particularly when the carbon or one of its neighbors is electronegative (e.g., in chloroform, aldehydes and terminal acetylenes).[12][13] Gradually, it was recognized that there are many examples of weaker hydrogen bonding involving donor Dn other than N, O, or F and/or acceptor Ac with electronegativity approaching that of hydrogen (rather than being much more electronegative). Though these "non-traditional" hydrogen bonding interactions are often quite weak (~1 kcal/mol), they are also ubiquitous and are increasingly recognized as important control elements in receptor-ligand interactions in medicinal chemistry or intra-/intermolecular interactions in materials sciences. The definition of hydrogen bonding has gradually broadened over time to include these weaker attractive interactions. In 2011, an IUPAC Task Group recommended a modern evidence-based definition of hydrogen bonding, which was published in the IUPAC journal Pure and Applied Chemistry. This definition specifies:
The hydrogen bond is an attractive interaction between a hydrogen atom from a molecule or a molecular fragment X–H in which X is more electronegative than H, and an atom or a group of atoms in the same or a different molecule, in which there is evidence of bond formation.[14]
As part of a more detailed list of criteria, the IUPAC publication acknowledges that the attractive interaction can arise from some combination of electrostatics (multipole-multipole and multipole-induced multipole interactions), covalency (charge transfer by orbital overlap), and dispersion (London forces), and states that the relative importance of each will vary depending on the system. However, a footnote to the criterion recommends the exclusion of interactions in which dispersion is the primary contributor, specifically giving Ar---CH4 and CH4---CH4 as examples of such interactions to be excluded from the definition.[4]
Nevertheless, most introductory textbooks still restrict the definition of hydrogen bond to the "classical" type of hydrogen bond characterized in the opening paragraph.
A hydrogen bond which is usually written as H-bond is a primarily electrostatic force of attraction between a hydrogen atom which is covalently bound to more electro negative atom or group, particularly second row elements like nitrogen, oxygen or fluorine.
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Answer:
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A hydrogen bond (often informally abbreviated H -bond) is a primarily electrostatic force of attraction between a hydrogen (H) atom which is covalently bound to a more electronegative atom or group, particularly the second-row elements nitrogen (N), oxygen (O), or fluorine (F)—the hydrogen bond donor (Dn)—and another electronegative atom bearing a lone pair of electrons—the hydrogen bond acceptor (Ac). Such an interacting system is generally denoted Dn–H···Ac, where the solid line denotes a fully covalent bond, and the dotted or dashed line indicates the hydrogen bond. The use of three centered dots for the hydrogen bond is specifically recommended by the IUPAC.[4] There is general agreement that there is actually a minor covalent component to hydrogen bonding, especially for moderate to strong hydrogen bonds (> 5 kcal/mol), although the importance of covalency in hydrogen bonding is debated. At the opposite end of the scale, there is no clear boundary between a weak hydrogen bond and a van der Waals (e.g., dipole-dipole) interaction.
Hydrogen bonds can be intermolecular (occurring between separate molecules) or intramolecular (occurring among parts of the same molecule).[5][6][7][8] Depending on the nature of the donor and acceptor atoms which constitute the bond, their geometry, and environment, the energy of a hydrogen bond can vary between 1 and 40 kcal/mol.[9] This makes them somewhat stronger than a van der Waals interaction, and weaker than fully covalent or ionic bonds. This type of bond can occur in inorganic molecules such as water and in organic molecules like DNA and proteins.
The hydrogen bond is responsible for many of the anomalous physical and chemical properties of compounds of N, O, and F. In particular, intermolecular hydrogen bonding is responsible for the high boiling point of water (100 °C) compared to the other group 16 hydrides that have much weaker hydrogen bonds.[10] Intramolecular hydrogen bonding is partly responsible for the secondary and tertiary structures of proteins and nucleic acids. It also plays an important role in the structure of polymers, both synthetic and natural.
Weaker hydrogen bonds[11] are known for hydrogen atoms bound to elements such as sulfur (S) or chlorine (Cl); even carbon (C) can serve as a donor, particularly when the carbon or one of its neighbors is electronegative (e.g., in chloroform, aldehydes and terminal acetylenes).[12][13] Gradually, it was recognized that there are many examples of weaker hydrogen bonding involving donor Dn other than N, O, or F and/or acceptor Ac with electronegativity approaching that of hydrogen (rather than being much more electronegative). Though these "non-traditional" hydrogen bonding interactions are often quite weak (~1 kcal/mol), they are also ubiquitous and are increasingly recognized as important control elements in receptor-ligand interactions in medicinal chemistry or intra-/intermolecular interactions in materials sciences. The definition of hydrogen bonding has gradually broadened over time to include these weaker attractive interactions. In 2011, an IUPAC Task Group recommended a modern evidence-based definition of hydrogen bonding, which was published in the IUPAC journal Pure and Applied Chemistry. This definition specifies:
The hydrogen bond is an attractive interaction between a hydrogen atom from a molecule or a molecular fragment X–H in which X is more electronegative than H, and an atom or a group of atoms in the same or a different molecule, in which there is evidence of bond formation.[14]
As part of a more detailed list of criteria, the IUPAC publication acknowledges that the attractive interaction can arise from some combination of electrostatics (multipole-multipole and multipole-induced multipole interactions), covalency (charge transfer by orbital overlap), and dispersion (London forces), and states that the relative importance of each will vary depending on the system. However, a footnote to the criterion recommends the exclusion of interactions in which dispersion is the primary contributor, specifically giving Ar---CH4 and CH4---CH4 as examples of such interactions to be excluded from the definition.[4]
Nevertheless, most introductory textbooks still restrict the definition of hydrogen bond to the "classical" type of hydrogen bond characterized in the opening paragraph.
Answer:
A hydrogen bond which is usually written as H-bond is a primarily electrostatic force of attraction between a hydrogen atom which is covalently bound to more electro negative atom or group, particularly second row elements like nitrogen, oxygen or fluorine.
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