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 (HONC rule as well).  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. The HONC Rule HH Hydrogen (and Halogens) form one covalent bond O Oxygen (and sulfur) form two covalent bonds One double bond, or two single bonds N Nitrogen (and phosphorus) form three covalent bonds One triple bond, or three single bonds, or one double bond and a single bond C Carbon (and silicon) form four covalent bonds. Two double bonds, or four single bonds, or a triple and a single, or a double and two singles

16. 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 atom wanting the most bonds central..

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

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

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