CHM 2103 Lecture Notes on Inorganic Chemistry - I (PDF)

# CHM - 2103 ## MODULE: ### PART: ## Inorganic Chemistry - I ### Lecture No. 7 - Driving Forces for Chemical Interaction - Electronegativity - Octet Rule ## Professor Brij B. Tewari # 1. Driving Forces for Chemical Interactions - Coulombic attraction and repulsion forces between nucleus an...

# CHM - 2103 ## MODULE: ### PART: ## Inorganic Chemistry - I ### Lecture No. 7 - Driving Forces for Chemical Interaction - Electronegativity - Octet Rule ## Professor Brij B. Tewari # 1. Driving Forces for Chemical Interactions - Coulombic attraction and repulsion forces between nucleus and electrons of two or more atoms are main driving forces for "Chemical Interaction". - Detailed nature of chemical interaction is a complex matter. One of the simplest, correct and broadly applicable ideas involved in such description is that a chemical bond can exist when outer orbitals on different atoms overlap so as to concentrate electron density between the atomic cores. - The overlap has a positive sign when superimposed regions of the two orbitals have the same sign, both + or both - - Overlap has a negative sign when superimposed region of two orbitals have opposite signs. - Precisely zero overlap results when there are precisely equal regions of overlap with opposite signs. - In a region where two orbitals $1 and $2 have positive overlap, the electron density is higher than the sum of the electron densities of two separate orbitals. That is ($1 + $₂) ²) is greater than ![Image of a graphic showing electron density distribution for H₂ ion. ] - $1² + $2² < 2 $1 $2. More electron density is shared between the two atoms. The attraction of both nuclei for these electrons is greater than the mutual repulsion of the nuclei and a net attractive force of bonding interaction therefore results. This is shown in the figure above. - The line (1) shows electron distribution in the 1 s orbitals for each atom $1² and $2², separately. - The line (2) shows the simple sum of these ($1² + $B²) 12 - If these two orbitals brought together with the same sign they give a positive overlap and the electron density will be given by ($a + $b)². This is shown as line (3), which lies above (2) through out the region between the nuclei, where it is simultaneously attracted to both of them and the H₂+ ion is more stable than H+ H or H + H+. - The same is the case for negative overlap, the shared electron density reduced by 2 $a $b and internuclear repulsion increases. This causes a net repulsion or anti-bonding interaction between the atoms. This is illustrated by Curve 4. The electron density is now also much lower between the nuclei, actually reaching zero at the mid point and nuclei repel each other strongly. - When the net overlap is zero, there is neither an increase nor a decrease in shared electron density and therefore neither a repulsive nor an attractive interaction. This is described as non-bonding interaction. # 7. Electronegativity - "Electronegativity of an element is the power of an atom of the element to attract electrons to itself when it is a part of a compound." - Electronegativity has been defined by several researchers as follows: (I) According to Pauling, electronegativity difference between two atoms is equal to 0.180 Δ. where Δ = resonance energy in kcal/mol. Resonance Energy: (Actual bond energy) - (Energy for 100% covalent bond) - Pauling introduced the idea that ionic character of a band varies with the difference in electronegativity. ![Image of a graph showing the relationship between electronegativity difference and ionic character.] Electronegativity difference - < 1.7 = More covalent nature of bond - > 1.7 = More ionic character of bond - 1.7 = 50% ionic and 50% covalent character (II) Robert Mulliken proposed a definition of electronegativity using data from atomic spectra. According to Mulliken, Xm as the average value of ionization energy and electron affinity Xm = 1/2 (IP + EA) There is a relationship between Pauling electronegativity Xp and Mulliken electronegativity Xm. Xp = 1.35√Xm - 1.37 Elements near to fluorine (F) have high IP and appreciable electron affinity. These elements have highest Mulliken electronegativities. (III) According to A.L. Allred and E. Rochow, electronegativity is determined by the electric field at the surface of an atom and is given by the expression: XAR = 0.744 + 0.3590 Zeff/(r/Å)² Zeff = Effective Nuclear Charge r = Covalent radius The value obtained by the above equation are parallel to Pauling electronegativities and are useful to discussing the electron distribution in the compounds. # 11. Applications of electronegativity - Some of the applications of electronegativities are as follows: 1. Description of bond energies. 2. Predicting of polarity of bonds and molecules. 3. Rationalization of the types of reactions that substances undergo. # 12. Octet Rule - Each atom acquires share in electrons until its valence shell achieves eight electrons. - Representative elements achieve a noble gas configuration in most of their compounds. - The rule is called octet rule because these configurations have 8e- in their outermost shell. - [Except for H₂, Li+, Be²+ which have 2e-] # Examples. - H: O: 0:: C:: 0: H H₂O CO₂ - H: N: H+ H NH₃+ - In the water molecule, the oxygen has a share in eight outer shell electrons giving it the neon configuration, while hydrogen shares two electrons, attaining the helium configuration. - Similarly, carbon and oxygen of CO₂ and nitrogen of NH₃ and NH₄+ ion each have share in eight e- in their outer shell, the neon configuration. The hydrogen atom in NH₃ and NH₄+ each share two electrons, the helium configuration. # Exceptions of Octet Rule - Octet rule is simple and useful in introducing the basic concept of bonding. Limitations are noted in dealing with ionic compounds of transition metals. Octet rule also fails in many situations involving covalent bonding. - (I) Molecules with an odd number of electrons. - (II) Molecules in which an atom has less than an octet. - (III) Molecules in which an atom has more than an octet. # (I) Odd Number of electrons - In few molecules such as ClO₂, NO, and NO₂ the number of electrons is odd. Complete pairing of these electrons is impossible, and an octet around each atom can not be achieved. For example, NO contains 5 + 6 = 11 valency e-. # (II) Less than an Octet - This type of exception occurs when there are fewer than eight electrons around an atom in a molecule or polyatomic ion. This situation is encountered in the compound of B and Be for example, BF₃. - There are only 6 e around the boron atom. # (III) More than an Octet - The third and largest class of exceptions consists of molecules or ions in which there and more than eight electrons in the valency shell of an atom. For example, PCl₅: - We force to expand the valency shell and place 10 e- around the central phosphorus atom. Other similar examples are SF₆, AsF₆, ICl₅...