01:146:295 Lecture Notes - Lecture 1: Molecular Orbital Theory, Valence Bond Theory, Orbital Hybridisation

• School: Rutgers University
• Department: Cell Biology and Neuroscience
• Course: 01:146:295
• Professor: schjott

Document Summary

Chapter 1 - covalent bonding and shapes of molecules. To draw lewis structures, we should know: number of valence electrons of each atom, covalent bonds formed between the atoms sharing of electrons to complete the octet & to attain nearest noble gas configuration. H atom can share one electron to attain the configuration of he. 2nd row elements share electrons to complete the octet. If the number of bonds formed is more or less as given above, this gives formal charge to the atom. To determine formal charge, draw the lewis structure: Fc = valence e - (no of shared e--) - no of unshared e s. Note: in polyatomic ions/molecules, the sum of the formal charge on individual atoms is equal to the total charge on the ion/molecule. Try: lewis structure of hn3, showing formal charges on all nitrogen. (in class) Shapes & bond angles can be explained using vsepr (valence shell e pair repulsion) approximation: