2. Covalent Bonding Bonding models for methane, CH 4. Models are NOT reality. Each has its own strengths and limitations.

3. Polar-Covalent bonds Nonpolar-Covalent bonds Covalent Bonds  Electrons are unequally shared  Electronegativity difference between.3 and 1.7  Electrons are equally shared  Electronegativity difference of 0 to 0.3

4. Covalent Bonding Forces  Electron – electron repulsive forces  Proton – proton repulsive forces  Electron – proton attractive forces

5. Bond Length and Energy Bonds between elements become shorter and stronger as multiplicity increases.

6. Bond Energy and Enthalpy D D = Bond energy per mole of bonds Energy requiredEnergy released Breaking bonds always requires energy Breaking = endothermic Forming bonds always releases energy Forming = exothermic

7. The Octet Rule sharing Combinations of elements tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level. Monatomic chlorineDiatomic chlorine

8. The Octet Rule and Covalent Compounds  Covalent compounds tend to form so that each atom, by sharing electrons, has an octet of electrons in its highest occupied energy level.  Covalent compounds involve atoms of nonmetals only.  The term “molecule” is used exclusively for covalent bonding

9. The Octet Rule: The Diatomic Fluorine Molecule F F 1s 2s 2p seven Each has seven valence electrons FF

10. The Octet Rule: The Diatomic Oxygen Molecule O O 1s 2s 2p six Each has six valence electrons O O

11. The Octet Rule: The Diatomic Nitrogen Molecule N N 1s 2s 2p five Each has five valence electrons N N

12.  Lewis structures show how valence electrons are arranged among atoms in a molecule.  Lewis structures Reflect the central idea that stability of a compound relates to noble gas electron configuration.  Shared electrons pairs are covalent bonds and can be represented by two dots (:) or by a single line ( - ) Lewis Structures

13. Comments About the Octet Rule  2nd row elements C, N, O, F observe the octet rule.  2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive.  3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals.  When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals.

14. Show how valence electrons are arranged among atoms in a molecule. Reflect the central idea that stability of a compound relates to noble gas electron configuration. Lewis Structures

15. C H H H Cl.. Completing a Lewis Structure - CH 3 Cl Add up available valence electrons: C = 4, H = (3)(1), Cl = 7 Total = 14 Join peripheral atoms to the central atom with electron pairs. Complete octets on atoms other than hydrogen with remaining electrons Make the least electronegative atom central (Never Hydrogen)..

16. Multiple Covalent Bonds: Double bonds Two pairs of shared electrons Ethene

17. Multiple Covalent Bonds: Triple bonds Three pairs of shared electrons Ethyne

18. Acetic Acid Two electrons (one bond) per hydrogen Eight electrons (four bonds) per carbon Eight electrons (two bonds, two unshared pairs) per oxygen