3.091 â Introduction to Solid State Chemistry Lecture Notes No. 2 ...

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LN–2The electrons involved are the 1s electron of the hydrogen atom and the unpaired 3p zelectron of the chlorine atom. After spin-pairing, a typical σ bond results in which eachatom attains a noble gas structure: in HCl the hydrogen atom is “involved” with twoelectrons, as in helium, and the chlorine atom with 18 electrons, as in argon. The other(inner) electrons of chlorine do not participate in the bonding; they are termednon-bonding.According to the above considerations, the bonding of HCl may be taken as very similarto that of H 2 . However, in the hydrogen molecule the electrons participating in the σbond formation do not favor the proximity of either of the hydrogen nuclei - instead, theyare equally shared. On the other hand, a preference for one nucleus (chlorine) is shownby the electrons of the σ bond in hydrogen chloride. This is because electrons, forenergetic reasons, favor the environment of the more electronegative atom; inhydrogen chloride this is the chlorine atom and, consequently, the σ bond electronsspend more time in the vicinity of the chlorine atom. This situation results in the chlorineend of the molecule being fractionally negatively charged (δ – ) and the hydrogen endbeing fractionally positively charged (δ + ). We denote such an internal chargeredistribution by the symbolism H δ+ - Cl δ– , where the δ sign indicates a partial electroniccharge; the bond is said to be polar. (The hydrogen molecule does not a priori exhibitsuch an “asymmetric charge distribution”; it may, however, acquire a temporarypolarization.) Molecules with asymmetric electronic charge distribution have apermanent dipole moment, the value of which is given by the product of the fractionalcharge (δ + must be equal to δ – ) and the distance of charge separation (L). Althoughdipole moments can be measured, their values do not allow us to directly calculate thepolarity of bonds between atoms since the detailed internal charge distribution isunknown. The direction of electron drift and, to some extent, the magnitude of its effectcan be estimated from the magnitude of the difference in electron affinity between the18
LN–2reaction partners. Because of a limited data base on electron affinity, L. Paulingintroduced a related term, the relative electronegativity (see below). From the aboveconsiderations it should be clear that no sharp dividing line exists between ionic andcovalent bonding. We might consider a completely ionic bond to result in cases wherethe electron drift is such that one atom (the cation) becomes entirely deficient in one ormore electrons, and the other atom (the anion) becomes correspondingly electron-rich,the bonding electrons being entirely under the influence of the latter. Hydrogen halidesare σ bonded and have dipole moments; the bonds are referred to as polar covalency.The homonuclear diatomic molecules formed by the halogens, viz. F 2 , Cl 2 , Br 2 and I 2 ,are all σ-bonded systems involving spin-pairing of the various p z electrons (2p z , 3p z ,4p z and 5p z respectively). They have no permanent dipole moments.Energetics of Covalent BondingPauling extensively treated the energetics of polar covalencies, encountered in allheteronuclear systems (such as H–Cl) which have permanent dipole moments(H δ+ – Cl δ– ). His approach visualizes the bonding to consist of two components - a purecovalency and an ionic bonding component with attraction resulting from the interactionof the fractional charges on the nuclei involved.Pauling determines the basic covalent bonding component, for example in theformation of HCl, from the experimentally obtained bond energies associated with themolecular species of the components, H 2 and Cl 2 . For this purpose he made the basicassumption that the covalent bond component between the dissimilar atoms is given bythe geometric mean of the pure covalent bond energies associated with these19
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3.091 â Introduction to Solid State Chemistry Lecture Notes No. 2 ...

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3.091 â Introduction to Solid State Chemistry Lecture Notes No. 2 ...