Unit 9 ~ Bonding (Chapter 12)

9-1 Introduction Key Concepts 1) All chemical bonds, regardless of type, are the result of attractions between opposite electrical charges: + (nucleus) and – (electrons). 2) The type of chemical bonding in a compound determines the chemical and physical properties of the compound. Example: allotropes of carbon: diamond, graphite, Buckyballs. 3) Chemical bonds are not static (stationary); bonds can stretch, bend, and in some cases, twist or rotate. The most common analogy to represent a chemical bond is a spring.

Chemical bonds form because the formation of a bond lowers the energy (more stable) of a group of atoms. Bonded atoms are more stable than the individual atoms.

The Na atom has 1 valence e- (3s1) The Cl atom has 7 valence e-s (3s23p5) By transferring the 1 electron from Na to Cl, both atoms reach the stable or inert configuration of s2p6 (lower energy!).  

Two Driving Forces for Bonding: 1. Individual atoms seek to optimize their valence electron configurations by LOSING, GAINING, or SHARING electrons. 2. Bonding electrons are simultaneously attracted to two positively charged nuclei.  

In the graph we see how energy is released when a bond is formed In the graph we see how energy is released when a bond is formed. Two H atoms, at the far right of the graph are brought closer together. The energy drops and reaches a minimum at the bond length of 74 pm. This lower energy reflects more stability for the bonded pair of atoms and a release of energy. The opposite would be true for breaking a bond.

9-2 Ionic Bonds (Sections 12.1, 12.4, 12.5) Ionic Bonds are formed by a complete transfer of one or more electrons between two atoms, typically between a metal and a nonmetal.

After the transfer of the electron(s) from the metal to the nonmetal, the + metal cation and the – nonmetal anion are attracted due to the opposite charges. The ions pack tightly in a crystal “lattice” to maximize the electrostatic attractions. Ionic bond formation is an example of a REDOX reaction: Review here

Ox # 0 0 +1 -1 2Na(s) + Cl2(g) → 2NaCl(s)   Oxidized (LEO) = Na Reduced (GER) = Cl Oxidizing agent (reduced) = Cl2 Reducing agent (oxidized) = Na # of electrons lost = # electrons gained

Lewis Structures for Ionic Bonds More review here Lewis Structures are convenient shorthand representations that illustrate the distribution of valence electrons in molecules. For each element, indicate the number of valence electrons with dots around the symbol:

Li one valence e - Be 2 valence e - B 3 valence e - C 4 valence e -

N O F Ne

Next, draw an ionic Lewis Structure for each of the following compounds. Sodium chloride is given as an example Na + Cl - * Formula Unit = lowest ratio of ions in crystal lattice structure (but recall they are not found this way – as are molecules)

MgCl2 AlBr3 Li2O Demonstrated on the board.....