1. Unit 6

2.  Atoms that are held together by sharing electrons  Usually 2 non-metals  Forms a molecule (or molecular compound)  Tend to have low melting and boiling points  Described by a molecular formula

3.  What is electronegativity?  How attracted are electrons to the atom  (how “likely” is the atom to become more negative by gaining an electron)

5.  Atoms tend to form bonds to acquire a total of 8 electrons  Single bond=1 pair of e -  Double bond=2 pairs of e -  Triple bond =3 pairs of e -  Electron dot structures are used to represent the shared pair of electrons

6.  Single Bond H H + H H H Each dash indicates a pair of shared e - Triple Bond Double Bond O + O OO OO N + N NN N N

7. Water, H 2 O

8. Ammonia, NH 3

9. Methane, CH 4

10. Propane, C 3 H 8

11. Propene, C 3 H 6

12. Propyne, C 3 H 4

13.  Step 1  count total valence e - involved  Step 2  connect the central atom (usually the first in the formula) to the others with single bonds  Step 3  complete valence shells of outer atoms  Step 4  add any extra e - to central atom IF the central atom has 8 valence e - surrounding it.. YOU’RE DONE!

14. Given below is an outline of how to determine the "best" Lewis structure for NO 3 -. 1. Determine the total number of valence electrons in a molecule 2. Draw a skeleton for the molecule which connects all atoms using only single bonds. In simple molecules, the atom with the most available sites for bonding is usually placed central. N (1) = 5 O (3) = 18 1 neg charge = 1

15. 3. Of the 24 valence electrons in NO 3 -, 6 were required to make the skeleton. Consider the remaining 18 electrons and place them so as to fill the octets of as many atoms as possible (start with the most electronegative atoms first then proceed to the more electropositive atoms). 4. Are the octets of all the atoms filled? If not then fill the remaining octets by making multiple bonds (make a lone pair of electrons, located on a more electronegative atom, into a bonding pair of electrons that is shared with the atom that is electron deficient).

16.  You only have two atoms, so there is no central atom, but follow the same rules.  Check & Share to make sure all the atoms are “happy”. Cl 2 Br 2 H 2 O 2 N 2 HCl

17. 1) CO 2 2) SiO 2 3) PCl 3 4) NO 2 -1 5) CH 3 F

18. 1) CO 2 2) SiO 2 3) PCl 3 4) NO 2 -1 5) CH 3 F

19. Resonance structures: Occur when it is possible to write 2 or more structural formulas for a compound Ex: NO 3 -

20. You must memorize these!! H 2 N 2 O 2 F 2 Cl 2 Br 2 I 2 Magnificent 7—

21.  Similar to ionic bonding  CO  CO 2

22. We need to say how many of each element we have! Use the prefixes! 1- mono6- hexa 2- di7- hepta 3- tri8- octa 4- tetra9- nona 5- penta 10- deca Examples: NO SiCl 4

23.  First Element: If you have more than one of the first element then you use a prefix. If there is only one then you just state the element  Second Element: Always has a prefix

24.  Formulas to names 1. SO 3 2. ICl 3 3. PBr 5 4. CO 5. CO 2  Names to formulas 1. Carbon tetrachloride 2. Dinitrogen monoxide 3. Dinitrogen tetroxide 4. Phosphorus triiodide 5. Sulfur heptafluoride

25. Hydrochloric- HCl Acetic Acid- HC 2 H 3 O 2 Nitric Acid- HNO 3 Sulfuric Acid- H 2 SO 4 Carbonic Acid- H 2 CO 3 Phosphoric Acid- H 3 PO 4

26. NON-Polar bonds  Electrons shared evenly in the bond  E-neg difference is zero Between identical atoms Diatomic molecules

27. Polar bond  Electrons unevenly shared  E-neg difference greater than zero but less than 2.0 closer to 2.0 more polar more “ionic character”

28.  HCl  CH 4  CO 2  NH 3 N2N2  HF a.k.a. “ionic character”

29.  Sometimes the bonds within a molecule are polar and yet the molecule is non-polar because its shape is symmetrical. H H HHC Draw Lewis dot first and see if equal on all sides

30.  Not equal on all sides  Polar bond between 2 atoms makes a polar molecule  asymmetrical shape of molecule

31. HCl -- ++

32. H -- ++