2. Quick Review Covalent bond – two atoms held together by sharing electrons -- Usually occurs between nonmetals. Octet Rule – chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest energy level. Valence electrons – the electrons in an atom’s outer energy level

3. Electron-dot diagrams shows an atom’s valence electrons Draw electron-dot diagrams for: NFNaO

4. New Terms Molecule – a neutral group of atoms held together by covalent bonds Molecular formula – shows the types and numbers of atoms combined in a single molecule (ex: H 2 O, CO 2, H 2 SO 4 ) Diatomic molecule – molecule containing two atoms (Examples: H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2, HCl, NO)

5. Using Electron Dot Notation diagrams H 2 molecule (nonpolar covalent = EQUAL sharing) H··H H : H Structural formula: H-H (Use single line for e- pair) the hydrogen orbitals overlap · · · ·

6. F 2 molecule ·· ·· ·· ·· : F· ·F : : F : F : ·· ·· ·· ·· Structural formula: F-F Each fluorine atom has 8 electrons in the outer energy level.

7. HCl molecule H Cl H Cl H – Cl Each atom now has a filled outer energy level.

8. Usually there will be more than two atoms bonding together. Water H 2 O Each hydrogen atom is bonded to oxygen atom H O H H Structural formula H O H

9. Try these molecules: NH 3, ClI, H 2 O 2 NH 3 molecule Each hydrogen atom is bonded to the nitrogen atom. H N H H

10. ClI molecule (chlorine and iodine) Cl I Structural formula Cl – I

11. H 2 O 2 molecule H O H O O H O H Structural FormulaH O O H

12. Bond Length and Energy Bond length – the average distance between bonded atoms Bond energy – energy required to break a chemical bond and form neutral isolated atoms.

13. Single Bonds Single bonds – a covalent bond produced by the sharing of one pair of electrons between two atoms. All of the examples so far have been of single bonds.

14. Double Bonds Double bonds – a covalent bond produced by the sharing of two pairs of electrons between two atoms. O 2 molecule O OStructural formula O=O

15. Triple Bonds Triple bonds – a covalent bond produced by the sharing of three pairs of electrons between two atoms. N 2 molecule N NStructural formulaN≡N

16. Relationships Between bond length and # of bonds: As the number of bonds INCREASES, the bond length DECREASES Between bond energy and # of bonds: As the number of bonds INCREASES, the energy required to break those bonds INCREASES

17. How to draw Lewis Structures (AKA electron dot diagrams) Two Types of Electron Pairs: Shared Pair – a bonded pair; exists between 2 atoms in the same molecule. Represented by a straight line between the bonded atoms. Unshared Pair – a lone pair; belongs entirely to one atom. Represented by a pair of dots on that atom.

18. Rules for Drawing Lewis Structures: 1) Draw a skeleton structure for the compound by joining the atoms with single bonds (single straight lines). The central atom is usually : a) The one with the highest number of valence electrons b) The largest atom c) The least electronegative atom

19. Rules for Drawing Lewis Structures: d) Hydrogen will NEVER be a central atom e) Oxygen will only be central if bonded to H or F

20. Rules for Drawing Lewis Structures: 2) Count/Tally the number of total valence electrons 3) Determine the number of electron pairs by dividing the total number of valence electrons (step #2) by 2 4) Determine the number of available electron pairs by subtracting the number of pairs already used in the skeleton structure

21. Rules for Drawing Lewis Structures: 5) Place LONE PAIRS around the terminal/end atoms. If any pairs are left, put them as lone pairs on the central atom 6) Check the OCTET RULE. If the central atom is not yet surrounded by 8 electrons, form multiple bonds.

22. Rules for Drawing Lewis Structures: Example: draw the Lewis Structure for Carbon Dioxide

23. Rules for Drawing Lewis Structures: Draw the Lewis Structure for Methanal (Formaldehyde) CH 2 O

24. Rules for Drawing Lewis Structures: CF 4 PCl 3 OCl 2 O 2