In his landmark paper, 'The Atom and the Molecule,' G.N. Lewis attempted to describe linkages between the atoms to understand the nature of covalent bonds.
He used dots to represent an atom’s valence electrons and argued that the atoms share their valence electrons to form one, two, or three bonds until they attain a stable octet electron configuration. An exception is the Hydrogen atom that attains a duplet configuration.
Full Charges | Partial Charges | Transient Charges |
|---|---|---|
Atoms attain full charges (+ or -) by losing or gaining electrons and forming ions. So, a full charge on an ion is equivalent to the charge of a proton (in case of electron loss) or an electron (in case of electron gain).
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Learning Objective: To learn about the ion-induced dipole interactions, their strength, and their occurrence.
Skill Level - Intermediate
Prerequisites:
Solution
The compounds in which the central metal atom is linked to ligands (anions or neutral molecules) that donate its pair of electrons to form coordinate covalent bonds with the metal atom.
Example, Ni(CO)4, {Co(NH3)6]Cl3
| sp3 | sp2 | sp |
|---|---|---|---|
Orbitals involved | One s and three p-orbitals of the central atom that are close in energy mix to form four sp3 hybrid orbitals for covalent bond formation. |
1) For comparison of the relative oxidizing and reducing powers of various metals: A Positive number indicates stronger reduction potential of the electrode, and it functions as a stronger oxidizing agent. A negative number means weaker reduction potential of the electrode, and hence it is a more powerful reducing agent.
