1. Chemical Bonding Chapter 7

2. Chemical bonding ionic bond: an electrostatic attraction between ions of opposite charge covalent bond: “sharing” electrons between atoms metallic bond: bonding occurs throughout the substance, electrons flow freely throughout the metal –(metals make good conductors)

3. Ionic bonding ions group together in a lattice formation. There is no “single” molecule of an ionic compound. What kind of energy changes are involved in formation of an ionic compound? Na (g) + Cl (g)  NaCl (g)

5. What trends are there in lattice energies? the potential energy between two particles. –Q=charge on particle 1 or 2, d=distance between the particles, k=constant. So, if magnitude of Q1 and Q2 increases, the energy _______. If d increases, the energy ______, but this change is not as much as when the charges are changed.

6. Crystal lattice energies Determine the order of the crystal lattice energies for the following ionic compounds, from lowest to highest MgO, KBr, ScN KBr (671 kJ/mol) < MgO (3795 kJ/mol) < ScN (7547 kJ/mol)

7. Lewis symbols a pictoral representation of the valence electrons of an atom –valence electrons vs. core electrons examples:

8. Ionic compounds In binary ionic compounds (i.e., no polyatomic ions involved), the atom to be the anion is formed by “swiping” the electrons of the atom to be the cation. Example: NaCl More examples:

9. Covalent bonds Instead of losing or gaining electrons, atoms “share” electrons so that a bond is formed between them.

10. Lewis formulas in covalent bonding Lewis formulas show how atoms “share” electrons between each other to form a bond. Examples: H 2 O, CH 4, NH 3

11. Octet rule Atoms tend to gain, lose, or share electrons until they are surrounded by EIGHT valence electrons. Why eight? –The octet rule is not a law, and there are several exceptions we will discuss later.

12. Rules for drawing Lewis structures 1.Sum all valence electrons from all atoms 2.Arrange the atoms –linear molecules (normally two or three), formula sometimes shows atoms from right to left –central-grouped atoms; center atom normally written first. –Least electronegative element in the center 3.complete octets around all atoms bonded to central atom –remember H has only 2 electrons 4.place all leftover electrons on central atom (even if more than an octet results) 5.try multiple bonds if the central atom doesn’t have an octet using single bonds Examples: H 2 O, CO, NH 3

13. Formal charge A bookkeeping tool used to help determine the best structure between two or more structures that follow the octet rule. (NOT the same as oxidation numbers) valence electrons = assigned electrons = all the unshared electrons (nonbonding electrons) + ½ all the bonding electrons FC = Valence electrons – assigned electrons Sum of FC of all atoms must = charge. Rule: The structure resulting in lowest formal charge on each atom “wins”. Example: structures of N 2 O

14. Resonance structures structures meeting both octet rule and formal charge requirements. Resonant structures are equivalent, but they are also equally wrong. The “actual” structure is in between all resonance structures. Examples: O 3, NO 3 -, SO 3

15. Exceptions to the octet rule 1.Odd number of electrons in molecule. 2.Molecules where an atom has less than an octet. 3.Molecules where atom has more than an octet. Exception 1 examples: NO, ClO 2 Exception 2 examples: BeCl 2, BF 3

16. Exception 3 examples (most common): PCl 5, SF 6, PO 4 3- Central atom expands into its d shell orbitals. For this to occur, central atom must be 3rd row or higher.

17. Polar and non-polar bonds Not all atoms share electrons equally in the atom. Electronegativity is an indicator of how much one atom will “dominate” the possession of electrons. –What is electronegativity? The difference in electronegativities (  EN) lets us know how polar a molecule is…

18. Dipole moments Dipole moments occur in polar molecules. The larger the  EN, the larger the dipole moment. The dipole moment points toward the more electronegative element. Molecules with more than two atoms are more complicated (we’ll talk about that in Chapter 8). Examples: