Ch. 6 Bonding 6.2 Covalent Bonding

Molecular Compounds molecule: neutral group of atoms held together by covalent bonds molecular compound: compound whose simplest unit is a molecule

Formulas chemical formula: tells the number of each type of atom in a compound molecular formula: tells the number of each type of atom in a molecular compound ex. H2O, Cl2, C6H12O2

Molecular Compounds diatomic molecule: a molecule containing only 2 atoms usually refers to 2 of the same atoms ex: O2, Br2, F2, etc. 7+1 rule

Formation of Covalent Bond

Formation of Covalent Bond two nuclei and two electron clouds repel each other creating an increase in PE approaching nuclei and electron clouds are attracted to each other to create a decrease in PE

Formation of Covalent Bond a distance between the nuclei is reached in which repulsion and attraction forces are equal potential energy is at the lowest point possible at the bottom of the curve on PE graph

Covalent Bonds Bond Length distance between two bonded atoms at their lowest PE average distance since there are some vibrations measured in pm (1012 pm = 1 m) stronger the bond, shorter the bond

Covalent Bonds Bond Energy energy is released when atoms combine because they have lower PE the same amount of energy must be used to break the bond and form neutral isolated atoms stronger bond, higher bond energy average since varies a small amount based on atoms in entire molecule in kJ/mol

Octet Rule representative elements can “fill” their outer energy level by sharing electrons in covalent bonds Octet Rule- a compound tends to form so that each atom has an octet (8) of electrons in its highest energy level by gaining, losing or sharing electrons Duet Rule- applies to H and He

Octet Rule Less than 8: More than 8: Boron: 6 in outer energy level anything in 3rd period or heavier because may use the empty d orbital ex: S, P, I

Electron Dot Diagrams a way to show electron configuration identifies the number and pairing of valence electrons to show how bonding will occur write the noble gas notation identify the number of valence identify how many are paired and how many are alone do not go by Figure 6-10

N Example Nitrogen Sulfur 1s2 2s2 2p3 5 valence 2 are paired 3 are alone Sulfur 1s2 2s2 2p6 3s2 3p4 6 valence 4 paired (2 pairs) 2 are alone N

Lewis Structures like dot diagrams but for entire molecules atomic symbols represent nucleus and core electrons and dots or dashes represent valence electrons unshared electrons: (lone pairs) pair of electrons not involved in bonding written around only one symbol bonding electrons: written in between 2 atoms as a dash

Types of Bonds single- sharing of one pair of electrons weakest, longest double- sharing of 2 pairs of electrons stronger and shorter triple- sharing of 3 pairs of electrons strongest and shortest multiple bonds include double and triple bonds

Drawing Lewis Structures find the number of valence electrons in each atom and add them up draw the atoms next to each other in the way they will bond add one bonding pair between each connected atoms add the rest of the electrons until all have 8 (consider exceptions to octet rule)

H H C Cl CH3Cl Example 1 methyl chloride C: 4 x 1 = 4 H: 1 x 3 = 3 Cl: 7 x 1 = 7 total = 14 electrons carbon is central H H C Cl duet octet duet octet duet H H C Cl

H N H H NH3 Example 2 ammonia N: 5 x 1 = 5 H: 1 x 3 = 3 total = 8 N is central Example 2 H N H H

Example 3 N2 nitrogen gas N: 5 x 2 = 10 10 electrons N N N N

H C H O Example 4 CH2O formaldehyde C: 4 x 1 = 4 H: 1 x 2 = 2 O: 1 x 6 = 6 total = 12 C is central H C H O

O O O O O O Example 5 O3 ozone O: 6 x 3 = 18 two completely equal arrangements the real structure is an average of these two where each bond is sharing 3 electrons instead of 4 or 2 O O O O O O

O O O O O O Resonance Structures resonance – bonding between atoms that cannot be represented by Lewis structure show all possible structures with double-ended arrow in between to show that electrons are delocalized O O O O O O

NO31- N: 5 x 1 = 5 O: 6 x 3 = 18 total = 23 + 1 = 24 Example 6

Covalent Network Bonding a different type of covalent bonding not specific molecules lots of nonmetal atoms covalently bonded together in a network in all directions example: diamond silicon dioxide graphite