Introduction to Covalent bonding Lewis Theory of Covalent

Introduction to Covalent bonding Lewis Theory of Covalent Bonding- The driving force of bond formation is the desire of each atom in a molecule to attain an octet of electrons in its valence shell. A. Covalent Bond- A chemical bond that is formed from a pair of “shared” electrons between two atoms 1. Caused by a small difference in electronegativities of the bonded atoms General Rule for covalent bonds: a. If Delectronegativities < 2. 0 (1. 7) b. If Allen Factor is < 0. 4 2. General Characteristics a. typically between 2 nonmetals b. occur from an “overlapping” of orbitals where shared electrons are attracted by both nuclei c. very strong bonds.

Introduction to Covalent bonding

Lewis structures – for CH 2 Cl 2 1. Draw a simple structural formula usingle covalent bonds. The least electronegative atom serves as the central atom. Note: Never use Hydrogen as the central atom 2. Determine the total number of valence electrons in the molecule or ion. a. For molecules- sum of the valence electrons on each atom b. For cations- sum of the valence electrons – the charge of the cation c. For anions- sum of the valence electrons + the charge of the anion 3. Deduct 2 valence electrons for each bond from 2. , then distribute the remaining about each atom as lone pairs to fill the octet (8). If too few electrons exist, convert single bonds to multiple bonds. Most multiple bonds occur between Carbon(s), Oxygen(s), Nitrogen(s), Sulfur, & Phosphorus.

Lewis structures – for CH 2 Cl 2 1. Draw a simple structural formula usingle covalent bonds. The least electronegative atom serves as the central atom. Note: Never use Hydrogen as the central atom 2. Determine the total number of valence electrons in the molecule or ion. a. For molecules- sum of the valence electrons on each atom b. For cations- sum of the valence electrons – the charge of the cation c. For anions- sum of the valence electrons + the charge of the anion 3 & 4. Deduct 2 valence electrons for each bond from 2. , then distribute the remaining about each atom as lone pairs to fill the octet (8). If too few electrons exist, convert single bonds to multiple bonds. -Most multiple bonds occur between Carbon(s), Oxygen(s), Nitrogen(s), Sulfur, & Phosphorus.

Lewis structures – for CH 2 Cl 2 1. Draw a simple structural formula usingle covalent bonds. The least electronegative atom serves as the central atom. Note: Never use Hydrogen as the central atom 2. Determine the total number of valence electrons in the molecule or ion. a. For molecules- sum of the valence electrons on each atom b. For cations- sum of the valence electrons – the charge of the cation c. For anions- sum of the valence electrons + the charge of the anion 3 & 4. Deduct 2 valence electrons for each bond from 2. , then distribute the remaining about each atom as lone pairs to fill the octet (8). If too few electrons exist, convert single bonds to multiple bonds. -Most multiple bonds occur between Carbon(s), Oxygen(s), Nitrogen(s), Sulfur, & Phosphorus. 5. Calculate formal charge for each element -the number of valence electrons in the free atomic state minus the number of assigned electrons of the atom in a molecule.

Lewis structures- practice CCl 2 O SCN 1 - NHF 31+