The Big Picture1

1. The importance of Coulombs Law:• Atomic attraction• Relative electronegativity• Electron repulsion model for shapes of molecules• Choice of dominant resonance contributors.

2. The tendency of electrons to spread out (delocalize):• Resonance forms• Bonding overlap

3. The correlation of the valence electron count with the Aufbau principle.• Associated stability of the elements in noble gas-octet-closed-

shell configurations obtained by bond formation.

4. The characteristic shapes of atomic and molecular orbitals:• Provides a feeling for the location of the “reacting” electrons

around the nuclei.

5. The overlap model for bonding:• Allows a judgment of energetics, directions and overall

feasibility of reactions.

Important Concepts1

1. Organic Chemistry – Chemistry of carbon and its compounds.

2. Coulomb’s Law – Relates attractive or repulsive force between charges to the distance between them.

3. Ionic Bonds – Result from coulombic attraction of oppositely charged ions.

4. Covalent Bonds – Result from electron sharing between two atoms.

5. Bond Length – Average distance between two covalently bonded atoms

6. Polar Bonds – Formed between atoms of differing electronegativity

Important Concepts1

7. Shape of Molecules – Strongly Influenced by electron repulsion.

8. Lewis Structures – Describe bonding using valence electron dots. Hydrogen receives a duet while other atoms receive an octet. Charge separation should be minimized but may be enforced by the Octet Rule.

9. Resonance Forms – When a structure is described by two or more Lewis structures differing only in their electron positions. The actual molecule is an average of the resonance forms. Some resonance structures may be more important that others.

10. De Broglie Relation – Relates wavelength of an electron to its mass and velocity.

Important Concepts1

11. Wave Equations – Describe motions of electrons about the nucleus. Solutions are called orbitals. These describe probabilities of finding the electrons in particular regions of space.

12. s Orbital – Spherical. P-orbital – Figure Eight. Each orbital can hold two electrons of opposite spin. With increasing energy, the number of nodes in an orbital increases.

13. Aufbau Principle – Building electronic configurations by adding one electron at a time to the atomic orbitals, starting with those of lowest energy. (Pauli exclusion principle, Hunds’ Rule).

Important Concepts1

14. Molecular Orbital – Two overlapping atomic orbitals form either a bonding or an antibonding molecular orbital. The number of molecular orbitals equals the number of atomic orbitals overlapped.

15. Bonds – Formed when atomic orbitals overlap along the bond axis. bonds – Formed from p-orbitals overlapping perpendicular to the bond axis.

16. Hybrid Orbitals - Formed by mixing of orbitals on the same atom. sp: 2 linear orbitals, sp2: 3 trigonal orbitals , sp3: 4 tetrahedral orbitals. Atomic orbitals not hybridized remain unchanged. Hybrid orbitals can contain either bonding or lone pair electrons.

Important Concepts1

17. Elemental Analysis – Determines ratios of types

of atoms in a compound. Molecular Formula – Actual number of atoms of each type.

18. Constitutional Isomers – (Structural Isomers) Same molecular formula but different connectivity of atoms. Different properties.

19. Condensed and Bond-Line Formulas – Abbreviated representations of molecules. Dashed-Wedged Line Drawings – Illustrate molecules in three dimensions.