Section 8.1 Types of Chemical Bonds 2. Covalent Bonding ▪ Electrons are shared by nuclei ▪ Usually formed by bonding nonmetals ▪ Non-polar covalent – electrons are shared equally ▪ Polar covalent – electrons are shared unequally ▪ Results in a partial positive and partial negative charge. Section 8.1 Types of Chemical Bonds 2 Section 8.1 Types of Chemical Bonds Ionic and Covalent bonding animation http://www.youtube.com/watch?v=QqjcCvzWwww Section 8.2 Electronegativity ▪ Hydrogen forms stable molecules where it shares two electrons. (Duet rule) Octet Rule Almost all other elements form stable molecules when surrounded by eight electrons. F + 7e- F F F 7e- 8e- 8e- Lewis structure of F2 single covalent bond lone pairs F F lone pairs single covalent bond lone pairs F F lone pairs 5 Section 8.2 Electronegativity ▪ Single covalent bond: two atoms share one pair of electrons. H–H ▪ Example: All halogens form single covalent bonds Copyright © Cengage Learning. All rights reserved All th Section 8.2 Electronegativity ▪ A double covalent bond: two atoms share two pairs of electrons. Example: O=O ▪ A triple covalent bond: two atoms share three pairs of electrons. Example: Section 8.2 Electronegativity Bond Dissociation Energy ▪ Energy required to break a covalent bond ▪ Large dissociation energy means strong bond ▪ The more bonds located between 2 atoms, the shorter the bond. ▪ The shorter the bond, the stronger the bond. H–H Single bond Weaker O=O Double bond Stronger N≡N Triple Bond Strongest Section 8.2 Bond Energies Electronegativity ▪ To break bonds, energy must be added to the system (endothermic, energy term carries a positive sign). ▪ To form bonds, energy is released (exothermic, energy term carries a negative sign). ∆H = ΣD(bonds broken) – ΣD(bonds formed) D represents the bond energy per mole of bonds (always has a positive sign). Copyright © Cengage Learning. All rights reserved Section 8.2 Electronegativity CONCEPT CHECK! Predict ∆H for the following reaction: Given the following information: Bond Energy (kJ/mol) C–H 413 C–N 305 C–C 347 891 [3(413) + 305 + 891] – [3(413) + 347 + 891] = –42 kJ ∆H = –42 kJ Copyright © Cengage Learning. All rights reserved Section 8.2 How to draw a Lewis structure Electronegativity 1. Count the total number of valence electrons. For ions, add 1 for each negative charge. Subtract 1 for each positive charge. 2. Choose the central atom. The atom in front is often the central. The central atom is usually the least electronegative, or the atom with the highest valence. 3. Draw a skeletal structure – connect the atoms. No ring around the rosy! 4. Arrange the remaining electrons to satisfy the octet rule. If there are not enough electrons, form double or triple bonds. Section 8.2 Electronegativity Example: Draw a Lewis structure for water. Notice that a pair of valence electrons are not shared. They are ‘lone’ pairs. Draw a Lewis structure for each of the following molecules: 1. NH3 5. NH4+ 2. HCN 6. CO2 3. CCl4 7. SO3 4. SO428. NO2- Copyright © Cengage Learning. All rights reserved Section 8.2 Electronegativity End lesson 2