Electron Deficient Molecules
Electron Deficient Molecules The electron deficient molecule may be defined as a molecule in which there are not enough bonding electrons available to join all the atoms in the molecule together by normal covalent bonds [i. e., electron pair bonds). In other words, electron deficient molecules may also be defined as molecules or substance in which the atoms have more stable orbitals in the valence shell than the electrons. For example, beryllium chloride is collinear in the vapour state and polymerised in solid state in which each beryllium...
read moreMetallic Bond
Metallic Bond Metals exhibit certain characteristic properties which are as follows: (i) High thermal and electrical conductivity. (ii) High melting and boiling points. (iii) High density and mechanical strength. (iv) High ductility and malleability. (v) Metallic lustre. (vi ) Alloy formation. (vii) Emission phenomenon. (viii) Power to replace hydrogen from acids etc. (ix) Crystalline solid. (x) Electropositive character. (xi) Opaque to light. The above said properties of metal cannot be explained on the basis of normal valencies. Hence a...
read moreHydrogen-bond
Hydrogen-Bond [Introduced By Latimer and Rodebush] It may be defined as, “the peculiar weak bond between the hydrogen atom of one molecule and highly electronegative atom (such as N, O, F) of same or another molecule”. For example several H –F molecules are associated by hydrogen bond as: The bond energy of hydrogen bond is of the order of hydrogen bond is represented by dotted line (…….) while the covalent bond is represented by solid line. The hydrogen bond energies of some typical molecules are shown...
read moreHybridisation
Hybridisation Pauling (1931) introduced the revolutionary concept of hybridization. The redistribution of energy of orbitals of individual atoms to give new orbitals of equivalent energy is called hybridisation. The new orbitals formed are known as hybrid orbitals. Different types of hybridisation along with hybrid orbitals and structures are given below: Before discussing the examples, we must mention here the hybridisation rules, which are as follows: (i) Orbitals of a central atom only would undergo hybrodisatoin. (ii) The orbitals of...
read moreVSEPR Theory
Valence Shell Electron Pair Repulsion (VSEPR) Theory This theory was developed by Gillespie and Nyholm. It is based on the effect of electron repulsion on the bond angles. The shape of the molecule or ion depends upon the number of bonding electron pairs (bp’s) and nonbonding electron pairs or lone pairs (lp’s) in the central atom. The central atoms are oriented in such a way that there is minimum repulsion (maximum stability) between them. The molecule has a definite shape because; there is only one orientation of orbitals corresponding...
read moreComparison of VBT and MOT
Similarities between VBT and MOT: (I) They account for the directional nature of the bond. (II) Bond results by the overlapping of two orbitals of minimum energy. (III) Electronic charge persists in between two atomic nuclei in bond formation. (IV) Atomic orbitals of same energy, and same symmetry overlap to produce strong bonds. Dis-similarities between VBT and MOT VBT MOT 1.Ineratomic orbital is produced by multiplying, exchanging and combinations of space wave functions of two electrons. 2.Orbitals of bonded atoms cannot lose their...
read moreBond Order
The bond order (B.O.) in diatomic molecules is half of the difference between the total numbers of the bonding electron ) and antibonding electrons therefore If bond order is zero, the molecule does not exist. Molecular orbital configuration of some homonuclear diatomic molecules and ions with their bond order etc, are given here. 1. Bond energy of and bond length is 2. Bond energy of and bond length is 3. Since bond order of helium molecule is zero hence helium does not exist as 4. The bond energy of molecule is low and...
read moreMolecular Orbital Theory
Molecular Orbital Theory (MOT) [Based on Linear Combination of Atomic Orbitals (L.C.A.O)] Hund and Mulliken have developed an approach to covalent bond formation which is based upon the effects of the various electron fields upon each other and which employs molecular orbitals rather than atomic orbitals. Each such orbital characterizing the molecule as a whole is described by a definite combination of quantum numbers and possesses relative energy value. Comparison between atomic and molecular orbitals Atomic Orbitals Molecular...
read moreSigma and Pi Bonds
Sigma and Pi Bonds According to orbital theory, covalent bond is formed by the result of coupling of electrons with opposite spins belonging to orbitals of outermost orbits of the two atoms. This invariably leads to lowering of potential energy of the system. Such orbitals are said to overlap with each other and the electron pair belongs to both the orbitals. This overlapping takes place in tow different ways and accordingly different types of bonds are formed: 1. Sigma Bond: A single bond is formed between two atoms by overlapping of...
read moreTypes of Overlapping
Types of overlapping (i) S-S overlapping: overlapping between s-s orbital’s of two similar or dissimilar atoms is known as s-s overlapping and forms a single covalent bond. (ii) S-P overlapping: overlapping between s- and p –orbital’s is known as s-p overlapping. is formed by the overlapping between 3 orbital’s of nitrogen ) with 3 orbital’s of 3 hydrogen atoms (s). Strong bond can be formed only when hydrogen electrons approache in the direction of X, Y and Z axes at right angles to each other. (iii) P-P overlapping: p-p...
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