Hydrogen bonding

The interactions described so far are universal in the sense that they are possessed by all molecules independent of their specific identity. However, there is a type of interaction possessed by molecules that have a particular constitution. A hydrogen bond is an attractive interaction between two species that arises from a link of the form A-H · · · B, where A and B are highly electronegative elements and B possesses a lone pair of electrons. Hydrogen bonding is conventionally regarded as being limited to N, O, and F but, if B is an anionic species (such as Cl−), it may also participate in hydrogen bonding. There is no strict cutoff for an ability to participate in hydrogen bonding, but N, O, and F participate most effectively. The formation of a hydrogen bond can be regarded either as the approach between a partial positive charge of H and a partial negative charge of B or as a particular example of delocalized molecular orbital formation in which A, H, and Beach supply one atomic orbital from which three molecular orbitals are constructed (Fig. 18.8). Thus, if the A-H bond is regarded as formed from the overlap of an orbital on A, ψ_A, and a hydrogen 1s orbital, ψ_H, and the lone pair on B occupies an orbital on B, ψ_B, then, when the two molecules are close together, we can build three molecular orbitals from the three basis orbitals:

ψ=c₁ψ_A+c₂ψ_H+c₃ψ_B

One of the molecular orbitals is bonding, one almost nonbonding, and the third anti-bonding. These three orbitals need to accommodate four electrons (two from the original A-H bond and two from the lone pair of B), so two enter the bonding orbital and two enter the nonbonding orbital. Because the antibonding orbital remains empty, the net effect—depending on the precise location of the almost nonbonding orbital—may be a lowering of energy. In practice, the strength of the bond is found to be about 20 kJ mol⁻¹. Because the bonding depends on orbital overlap, it is virtually a contact-like interaction that is turned on when AH touches B and is zero as soon as the contact is broken. If hydrogen bonding is present, it dominates the other intermolecular interactions. The properties of liquid and solid water, for example, are dominated by the hydrogen bonding between H2O molecules. The structure of DNA and hence the transmission of genetic information is crucially dependent on the strength of hydrogen bonds between base pairs. The structural evidence for hydrogen bonding comes from noting that the internuclear distance between formally non-bonded atoms is less than their van der Waals contact distance, which suggests that a dominating attractive interaction is pre sent. For example, the O-O distance in O-H···O is expected to be 280 pm on the basis of van der Waals radii, but is found to be 270 pm in typical compounds. Moreover, the H···O distance is expected to be 260 pm but is found to be only 170 pm. Hydrogen bonds may be either symmetric or unsymmetric. In a symmetric hydro gen bond, the H atoms lies midway between the two other atoms. This arrangement is rare, but occurs in F-H···F−, where both bond lengths are 120 pm. More common is the unsymmetrical arrangement, where the A-H bond is shorter than the H···B bond. Simple electrostatic arguments, treating A-H···B as an array of point charges (partial negative charges on A and B, partial positive on H) suggest that the lowest energy is achieved when the bond is linear, because then the two partial negative charges are furthest apart. The experimental evidence from structural studies support a linear or near-linear arrangement.

Fig. 18.8 The molecular orbital interpretation of the formation of an A H···B hydrogen bond. From the three A, H, and B orbitals, three molecular orbitals can be formed (their relative contributions are represented by the sizes of the spheres. Only the two lower energy orbitals are occupied, and there may therefore be a net lowering of energy compared with the separate AH and B species.

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Units

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Working galvanic cells

Electrolysis

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The rate laws

The electric potential at the interface

The structure of the interface

Catalytic activity at surfaces

Mechanisms of heterogeneous catalysis

Mobility on surfaces

The rate of desorption