Covalent Bonding

Covalent Bonding Covalent bond –formed from the sharing of electrons Usually between 2 non metallic elements

Molecules Molecule – two or more atoms covalently bound together Diatomic molecule – two of the same atom bound together

Diatomic Molecules Br I N Cl H O F or the Magnificent 7 These atoms never exist alone. They always come in pairs For example: Br  Br2 I  I2 N  N2 Cl  Cl2 H  H2 O  O2 F  F2

Binary Molecular Compounds Binary Compounds consist of 2 elements Binary covalent compounds can be recognized by containing 2 nonmetals This is different from ionic compounds that contain a metal & nonmetal, metal & a polyatomic ion, or 2 polyatomic ions

Lewis Structure Lewis Structures – shows how the valence electrons are arranged among the atoms of a molecule There are rules for Lewis Structures that are based on the formation of a stable compound Atoms want to achieve a noble gas configuration

Octet & Duet Rules Octet Rule – atoms want to have 8 valence electrons Duet Rule – H is the exception. It wants to be like He & is stable with only 2 valence electrons

Steps for drawing Lewis Structures Sketch a simple structure with a central atom and all attached atoms Add up all of the valence electrons for each individual atom If you are drawing a Lewis structure for a negative ion add that many electrons to create the charge If you are drawing a Lewis structure for a positive ion subtract that many electrons to create the charge

Steps for drawing Lewis Structures Subtract 2 electrons for each bond drawn Complete the octet on the central atom & subtract those electrons Complete the octet on the surrounding atoms & subtract those electrons Get your final number If 0  you are done! If +  add that many electrons to the central atom If -  need to form multiple bonds to take away that many electrons

Examples CCl4 Sketch a simple structure with a central atom and all attached atoms Cl │ Cl – C – Cl

Examples Add up all of the valence electrons for each individual atom 4 + 4(7) = 32 Subtract 2 electrons for each bond drawn 32-8 = 24 Complete the octet on the central atom & subtract those electrons Done

Examples Complete the octet on the surrounding atoms & subtract those electrons 24 – 24 = 0 Final number = 0…DONE! Final structure is… __ │Cl │ __ │ __ │Cl – C – Cl │ │

Examples N2

Examples NH3

Examples PO4 -3

Bond Types Sigma bonds () – single covalent bond Pi bonds () – occur when multiple bonds are formed Single bond – sigma Double bond – 1 sigma & 1 pi Triple bond – 1 sigma & 2 pi

Bond length & Strength As the number of bonds increases, the bond length decreases The shorter the bond, the stronger the bond