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Warmup P-H N-H F-H

Warmup P-H N-H F-H a. Use your EN chart to calculate the EN differences of the bonds shown above b. Rank the bonds in order of decreasing polarity (most polar to least polar): c. Draw the distribution of the electron cloud around each bond and the dipoles. Lewis Structures.

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Warmup P-H N-H F-H

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  1. Warmup P-H N-H F-H a. Use your EN chart to calculate the EN differences of the bonds shown above b. Rank the bonds in order of decreasing polarity (most polar to least polar): c. Draw the distribution of the electron cloud around each bond and the dipoles

  2. Lewis Structures

  3. Lewis structures are drawings that show HOW atoms are covalently bonded. Atoms in a covalent bond share valence e- to achieve completely satisfied valence shells. Bonding (Shared) Pair Lone (Unshared) Pair Lone (Unshared) Pair Bonding (Shared) Pair

  4. I Ex. Draw the Lewis Structure for ICl Hmmmmm how many val e- does iodine have? What about chlorine?

  5. I Cl

  6. I Cl

  7. I Cl

  8. I Cl

  9. I Cl

  10. I Cl

  11. I Cl *They each share 1 e- with each other

  12. How many val e- does each atom have? • Both end with “full octets”. All elements MUST end up with a full octet (few exceptions) I Cl 8 Valence electrons 8 Valence electrons

  13. What about harder structures? Let’s try like…..a method… Ex. Draw the Lewis Structure for H2CO 1) Calculate total # of valence e- (4 e-)(1 atom) = 4 C = (1 e-)(2 atoms) = 2 H = O = (6 e-)(1 atom) = 6 12 valence e-

  14. Draw the Lewis Structure for H2CO 2) Plan the molecular skeletonHints: C is often central, H and halogens are never central, O is rarely central 3) Place 1 pair e- between each of the atoms 4&5) Add in e- to create double/triple bond or lone pairs so that: 6) Check that each atom is surrounded by 8 valence e- and that total valance e- are all used 6 val e- H C O H

  15. 7. Final structure: replace each bonding e- pairs with a line Clarification: the structural formula shows bonds as lines and the lone pairs aren’t ALWAYS included. A dot diagram represents bonding electrons pairs as dots. A Lewis structure can technically be either, we are not too particular. H C = O H

  16. Simple Structures Practice CH4 HF H2 H2O HF and H2O are polar molecules (uneven distribution of electrons)

  17. More Practice • SF6 • BCl3 • WEIRDNESS!!!! • CH3Br • NH3

  18. expanded octet Sulfur usually makes 2 bonds but can make up to 6 Phosphorus can also expand it’s octet. Nitrogen cannot.

  19. Incomplete Octet BCl3 BeF2 on

  20. More Practice • OH- • Total e- • 6 + 1+ 1 • NH4+ • Total e- • 5 + 4– 1 • N2 • 10 e- total, but double or triple bond? [ ]-1 H O

  21. Sometimes there are various Lewis structures possible! To figure out which is the most plausible, we calculate formal charges on each atomin these structures: f.c.= [valence – belonging] • Take valence e- of that element • Subtract the e- in lone pairs on that element • Subtracthalfthe number e- shared by that atom in that structure : N = N : : N = N : 0 0 +1 : -1 try niTRATE(NO3-)

  22. Most likely Lewis structure : formal charge 0 on all atoms. If not zero:-formal charges as small as possible-negative formal charges on most electronegative atoms-adjacent atoms should not carry formal charges of same sign-add up to 0 (molecule) or the net charge (polyatomic ion)formal charge will help you draw the very best structure, but we’re not going to get too crazy with this ok? -1 Formal charges add up to -1 0 +2 +1 -1 -1 -1 -1

  23. This is not the most likely structure for SO42- Draw a more stable Lewis structure for the sulfate ion, with a lower formal charge on the sulfur atom. There are several options! Hint: try (a) double bond(s)?) Total electrons: S: 6 O:4 x 6 ion: 2 extra e- = 32 e-

  24. A bond is not ACTUALLY 2 dots (or 4 or 6) in between two atoms!!!!!!!!!!!!!!!!!!!! A bond is formed when one orbital (containing an electron) overlaps with an orbital from another atom. The electrons are shared and travel within the two orbitals….they are SHARED. Some orbitals are full and don’t have to overlap (lone pairs!) O C O 1e- You will not need to draw something like this!!! 2e- 1e- 1e- 1e- 1e- 2e- 1e-

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