1. Chemical Bonds Modern Chemistry: Chapter 6 Why? How? What? Where?

2. Why? Electromagnetism is 1 of the 4 universal forces Balance between repulsion & attraction – Protons repel protons; electrons repel electrons – VSEPR = valence shell electron pair repulsion – Positive nuclear charges attract electrons Electrons are from self & nearby atom(s) Compounds are more stable than free atoms/ions – Lower potential energy at the optimal bond length – Energy is released upon bond formation, typically Heat &/or Light Sound Movement

3. Octet Rule answers “Why?” Bonds typically form to give 8 e - in outer shell. This provides a “noble gas configuration.” There are, however, exceptions. – BF 3 & AlCl 3 only have 6 electrons for B & Al – PF 5 & SF 6 have more than 8 electrons for atoms – H 2 & He only have 2 electrons in outer shell

4. What? Types of bonds – Ionic chemical bonds Oppositely charged ions attract one another – Covalent chemical bonds Nonpolar = Equal sharing of electrons Polar = Unequal sharing of electrons – Metallic bonds Mobile sea of electrons surround cations

5. How? Chemical bonds are determined by differences in the degrees of electronegativity (0 – 4.0) – Ionic = difference > 2.0 Alkali & alkaline earth metals with halogens/ nonmetals Some transition metals with nonmetals – Polar covalent = difference of 0.6 – 1.9 Nonmetals with one another Some transition metals with nonmetals – Nonpolar covalent = difference of 0 – 0.5 Diatoms (I, Br, Cl, F, O, N, H) Nonmetals with very similar nonmetals

6. How else? Molecular orbitals form around atoms Size of atoms/ ions influence bond formation Unshared valence electrons affect the shape of the entire molecule (molecular geometry) Bond length: minimum potential energy – Single bonds are longest Two electrons are involved Smallest bond energy (least repulsion) – Triple bonds are shortest Six electrons (3 pairs) are shared

7. What about metallic bonds? Vacant orbitals in outer energy levels overlap Overlapping orbitals allow outer electrons to roam freely throughout the entire metal – Malleability, Ductility, & Conductivity Many orbitals spaced by incremental energy levels allow absorption of many frequencies – Luster

8. Where? Covalent bonds form in the electron cloud –  bonds: symmetrical along nuclei’s axis –  bonds: side by side overlap of p orbitals in sausage-shaped regions above & below axis Ionic bonds form in the charged space of the electron cloud Metallic bonds form between cations of the same element as a mobile sea of electrons

9. Electron-Dot Notation Illustrates only the valence electrons – Nucleus & inner-shell electrons = element symbol – Valence electrons shown as dots: E, N, W, & S Compounds shown as Lewis structures – Shared valence electrons = dot-pairs or dash(es) – Unshared valence electrons = dot(s) Structural formulas don’t show unshared pairs  F-F, H-Cl, K-I, etc…

10. Resonance Structures A single representation is inadequate Molecule may constantly alternate between bonding structures Molecule may form an average of 2 structures – Ozone (O 3 ) forms identical O-O bonds that are between a single & a double bond

11. Ionic Compounds Formula unit = simplest collection of atoms Crystal lattice = 3-D arrangement of ions – Cubic- Monoclinic-Triclinic – Tetragonal- Hexagonal – Orthorhombic- Rhombohedral Lattice energy (kJ/mol) = energy released upon crystal formation from gaseous ions Polyatomic Ions: NH 4 +, MnO 4 -, SO 4 =, etc… – Ions held together by covalent bonds.

12. VSEPR Theory Molecules form to lessen e - pair repulsion Diatoms  form linear (180 o ) Group III/13  form trigonal-planar (120 o ) Group IV/14  form tetrahedral (109.5 o ) Group V/15  form trigonal-pyramidal (107 o ) Group VI/16  form bent or angular (105 o ) SF 6 types  form octahedral (90 o )

13. Hybridization Atomic orbitals mix & form equal hybrid orbitals on the same atom – s & p orbitals  sp orbital (180 o ) BeF 2 – s, p, & p orbitals  sp 2 orbital (120 o ) BF 3 – s, p, p, & p orbitals  sp 3 orbital (109.5 o ) CCl 4

14. Intermolecular Forces Dipole-dipole forces – Separated equal but opposite charges  Dipole – Represented by arrow with head toward (-) pole Forces of attraction between polar molecules Hydrogen bonding – H of 1 molecule pulled to (-) charge on another London dispersion forces – Attractions from creation of instantaneous dipoles