Structures and Properties Chemistry 11

Covalent Bonds and Molecules When we consider covalent compounds, it is often important to consider the molecular geometry This is vary with the number of lone and shared pairs as well as single, double and triple bonds We will use Lewis Structures to model simple molecules

Pairs – Lone or Shared? There are two types of electron pairs that we consider in molecules Lone – these are pairs that belong entirely to one atom Shared – these are pairs that are shared between two of the atoms in the molecule

Drawing Lewis Structures – Simple Molecules Determine the total number of valance electrons in all atoms Draw a skeleton structure where the atom with the largest number of unpaired electrons is the central atom Place lone pairs around all of the atoms except the central atom to form octets (except hydrogen which has one pair) If all valance electrons have not been used, add long pairs to central atom If all valance electrons have been used but the central atom does not have an octet, move lone pairs to form double or triple bonds as required

Example - Formaldhyde O O H H C C H H CH2O Valance Electrons H – 2x1 O – 6 Total - 12 O O H H C C H H

Co-ordinate Covalent Bonds Normally in a covalent bond, each atom contributes one electron to the shared pair However, it is possible that one atom contributes both electrons; this is known as a co-ordinate covalent bond This behaviour explains how ammonium is formed

[ ] Example - Ammonium H + H N H H NH4+ Valance Electrons Total - 8 H – 4x1 + – -1 Total - 8 H + [ ] H N H H

Resonance Structures These occur when there is more than one possible Lewis Structure These models will have the same chemical formula but different placing of bonding and lone pairs

Practice Problems Page 189 Questions 1-5