How do hydrogen bonds occur in HFCs

Hydrogen bond

Snapshot of a computer simulation of liquid water. The individual water molecules are shown, oxygen is shown in red, hydrogen in white. The proportions of the atoms are Not to scale, only the distances between the atomic nuclei. The dashed blue lines represent hydrogen bonds. In addition, the length of the individual hydrogen bonds is noted in Ångström.

Hydrogen bonds, also briefly Hydrogen bonds or H bridges, bonds are electrostatic in nature. They are a form of the minor valence bond and their strength is usually well below that of the covalent atomic bond and ionic bonds. In water it is based on the fact that the hydrogen atoms and the oxygen atom in the water molecule are not on a straight line (H – O – H), but form an angle, which means that the electrical charges are distributed asymmetrically: negative on the oxygen atom, positive on the hydrogen atoms. These charges of neighboring molecules attract each other and so chains and larger groups (clusters) are formed. The individual water molecules are therefore not free to move and need z. B. a lot of energy to transition into the gaseous state, d. that is, the boiling point of water is relatively high.

In proteins, hydrogen bonds maintain the important three-dimensional structures of the molecules. Finally, in DNA, hydrogen bonds connect the individual strands to form a characteristic double helix. This principle is therefore a necessary prerequisite for all known living beings.

discovery

The concept of hydrogen bonds was first described in 1920 by Wendell Mitchell Latimer and Worth H. Rodebush to explain the high dielectric constant of water.[1]

Creation of the bond

Hydrogen bonds are formed when two functional groups interact via hydrogen atoms. It does not matter whether the bond is created between two molecules or two separate sections of a macromolecule. A distinction is made between the functional groups Proton donor (also: Donor, Donor) and the Proton acceptor. The donor is an electronegative atom (e.g. nitrogen, oxygen, or fluorine) to which a hydrogen atom is covalently bonded. Due to the difference in electronegativity, two partial charges are formed here. One negative on the electronegative atom, and one positive on the bonded hydrogen atom. One speaks here of a polar bond. The acceptor is any other electronegative atom with lone pairs of electrons.

If the donor (–X – H) and acceptor (Y–) are geometrically aligned in such a way that the acceptor's lone pair of electrons points in the direction of the hydrogen atom, a hydrogen bond is formed. In general, such a system is represented as follows:

$ \ mathrm {R ^ 1 \ mathord-X \ mathord-H \ cdots Y \ mathord-R ^ 2 \} $

The dotted line symbolizes the hydrogen bridge bond.

It should be noted that certain functional groups can act as donors and acceptors at the same time. A simple example is hydroxyl groups or the hydrogen bonds between water molecules.

Classification of hydrogen bonds

George A. Jeffrey[2] introduced a classification of hydrogen bonds. A distinction is made here between the strength of the bond.

  • Strong bonds (63–167 kJ / mol): For example in hydrofluoric acid H – F ··· H – F
  • Medium ties (17–63 kJ / mol): For example in water or in carbohydrates.
  • Weak ties (<17 kJ / mol): C − H ··· O interactions, for example in proteins.[3]

Effects of hydrogen bonds

Hydrogen bonds in biomolecules

Hydrogen bonds are responsible for the special properties of many molecules that are important for living things:

Hydrogen bonds in hydrogen compounds in general

Hydrogen compounds of elements in the 6th main group have a tetrahedral geometry and two possibilities to act as acceptor and donor. Each molecule can form four hydrogen bonds, creating a particularly dense network of hydrogen bonds. These compounds include water (H.2O), hydrogen sulfide (H.2S) and hydrogen selenide (H.2Se).

In the event of a phase transition, these hydrogen bonds must be broken. Therefore, compared to the hydrogen compounds of the other elements in their respective periods (which cannot form as many hydrogen bonds) they have increased melting and boiling temperatures.

Within this main group, the hydrogen compound of oxygen, i.e. water (H.2O), the highest melting and boiling temperatures, because oxygen has the highest electronegativity and can therefore also form the strongest hydrogen bonds. The hydrogen bonds of the elements in the course of the group (“below”) are less electronegative and form weaker hydrogen bonds. Therefore the molar mass is the more decisive factor for the melting and boiling temperatures (from H.2They rise again).

element 6
C.
7
N
8
O
9
F.
15
P.
16
S.
17
Cl
35
Br
Hydrogen compound CH4NH3H2OHFPH3H2S.HClH2SeHBr
Melting point in ° C −182 −77 0 −83 −134 −85 −114 −66 −87
Boiling point in ° C −162 −33 100 19 −88 −60 −85 −41 −67

Hydrogen bonds of water in particular

The ideal tetrahedral structure of the water cluster: each oxygen points to 4 hydrogens and each hydrogen to two oxygens
Spatial networking of the water molecules

Due to the higher electronegativity of oxygen with 3.5 compared to that of hydrogen with 2.1, the water molecule has a dipole character. The oxygen is therefore partially negative (δ−), the hydrogen partially positive (δ +). The hydrogen bonds form between the different partial charges.

Hydrogen bonds are responsible for a number of important properties in water. Among them are the liquid state of aggregation under normal conditions, the cohesion, the relatively high boiling point and the density anomaly of the water

The typical bond length of hydrogen bonds in water is 0.18 nm. There are two types of bonds. So-called linear bonds with a bond angle of 180 ° and non-linear 180 ° ± 20 °, with the linear bond predominating. Whereas a purely tetrahedral network (bond angle 180 °) would have to lead to 4 nearest neighbors (coordination number 4) the coordination number measured (by X-ray scattering) under normal conditions is 4.5.

With decreasing density, this measure of order decreases (in contrast to an increase in the coordination number in most other liquids) to 4 and thus to the value for an ideal tetrahedral structure.

Mainly 2, 4 or 8 water molecules are connected to each other in water. When evaporating, these must be separated; This also explains the high energy expenditure (compared to other substances) to convert liquid water at 100 ° C into steam at 100 ° C (see heat of vaporization).

Intramolecular hydrogen bonds

pseudocyclic structure of ricinoleic acid
If several donors / acceptors occur within a molecule, hydrogen bonds can form within the molecule, as is the case with ricinoleic acid. Both a hydroxyl group and a carboxy group are present there. The hydroxyl group should actually increase the melting or boiling point. However, the boiling point of ricinoleic acid is even lower than that of oleic acid, which only differs in the lack of a hydroxyl group, since the hydroxyl group forms hydrogen bonds with the carboxyl group. The carboxy group can therefore no longer enter into intermolecular hydrogen bonds to the same extent as if the intramolecular hydroxyl group were not present.

The spatial structure is also changing; pseudocyclic structures arise due to the forces of attraction between the polar groups.

literature

General textbooks

Special books

  • George A. Jeffrey: An Introduction to Hydrogen Bonding. Oxford University Press, 1997, ISBN 978-0195095494.
  • George C. Pimentel, A.L. McClellan: Hydrogen Bond. W. H. Freeman & Co Ltd., San Francisco 1960, ISBN 978-0716701132.
  • Anthony J. Stone, A. J. Stone: The Theory of Intermolecular Forces. Oxford University Press, Oxford 1997, ISBN 978-0198558835.

Web links

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Individual evidence

  1. ↑ Wendell M. Latimer, Worth H. Rodebush: Polarity and Ionization from the Standpoint of the Lewis Theory of Valence. In: Journal of the American Chemical Society. 42, 1920, pp. 1419-1433, doi: 10.1021 / ja01452a015 (full text).
  2. ↑ George A. Jeffrey: An Introduction to Hydrogen Bonding. Oxford University Press, 1997, ISBN 978-0195095494.
  3. ↑ Lin Jiang, Luhua Lai: CH ⋯ O Hydrogen Bonds at Protein-Protein Interfaces. In: J. Biol. Chem.. 277, No. 40, 2002, pp. 37732-37740, doi: 10.1074 / jbc.M204514200.
  4. ↑ L. Pauling, R. B. Corey, H. R. Branson: The Structure of Proteins: Two Hydrogen-Bonded Helical Configurations of the Polypeptide Chain. In: Proceedings of the National Academy of Sciences of the United States of America. 1951, pp. 205-211.
  5. ^ L. Pauling, R. B. Corey: Configurations of polypeptide chains with favored orientations of the polypeptide around single bonds: Two pleated sheets. In: Proceedings of the National Academy of Sciences of the United States of America. 37, 1951, pp. 729-740.
  6. ↑ M. A. Williams, J. E. Ladbury: Hydrogen Bonds in Protein-Ligand Complexes. In: H.-J. Böhm, G. Schneider (Ed.): Protein-ligand interactions. Wiley-VCH, Weinheim 2005, ISBN 978-3-52730521-6, pp. 137-161, doi: 10.1002 / 3527601813.ch6.