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 Large ∆E Electrons are equally shared Small ∆E

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

5. 4 Bond Length –The distance where the system energy is lowest. –Triple bond < Double Bond < Single Bond

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

7. 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

8. 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

9. 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

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

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

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

13.  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

14. 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.

15. 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

16. 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

17. 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 most metallic atom the central atom..

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

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

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

21. Two possible skeletal structures of formaldehyde (CH 2 O) HCOH H CO H formal charge on an atom in a Lewis structure = total number of valence electrons in the free atom - total number of nonbonding electrons total number of bonds - The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion. Formal Charge

22. 21 HCOH formal charge on C = 4 - 2 -3 = -1 formal charge on O = 6 - 2 -3 = +1 formal charge on an atom in a Lewis structure = total number of valence electrons in the free atom - total number of nonbonding electrons - +1 total number of bonds

23. 22 H CO H formal charge on C = 4 - 0 - 4= 0 formal charge on O = 6 - 4 - 2= 0 formal charge on an atom in a Lewis structure = total number of bonds total number of valence electrons in the free atom - total number of nonbonding electrons - 00

24. 23 Formal Charge and Lewis Structures 1.For neutral molecules, a Lewis structure in which there are no formal charges is preferable to one in which formal charges are present. 2.Lewis structures with large formal charges are less plausible than those with small formal charges. Which is the most likely Lewis structure for CH 2 O? HCOH H CO H HCOH +1 H CO H 00

25. 24 Formal Charge and Lewis Structures 3.Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms. HCOH H CO H

26. Worksheet 8.3 Lewis Structure Practice Skip the following: SO 2 O 3 2 NO 3 -

27. Worksheet 8.4 Lewis Structure Instead of Steric Number Calculate formal charges SO 2 NO 2 C 6 H 6 LiH SiO 2 AlF 3 CO 3 2- NO 3 - SO 4 2- AlH 3 CH 4 PO 4 2- HCN