Lewis Dot Structures. Team Chemistry Lanier H.S.

Transcription

1 Lewis Dot Structures Team Chemistry Lanier H.S.

2 Part 1: Review of Lewis Dot Symbols

3 To Draw a Lewis Dot Symbol: 1. Write the symbol for the atom 2. Find the number of valence electrons (use Periodic Table) 3. For every valence electron, draw dot around the symbol Example #1: Sodium Na *Sodium has 1 valence electron, so we draw one dot.

4 Arrangement of Electrons Pretend there s a box around the symbol Draw the first e- on one side of the box, then rotate to the next side and draw another Keep rotating until you ve drawn them all Example #2: Carbon C *Carbon has 4 valence electrons

5 Arrangement of Electrons Up to two electrons can be on each side Valence e- prefer to be in pairs (one of the reasons atoms bond with other atoms is to pair up their valence e-) Example #3: Sulfur S *Sulfur has 6 valence electrons Can use drawing to determine the charge! Sulfur will gain 2 e- to get to 8, so charge is -2.

6 Part 2: Lewis Dot Structures Lewis Dot Structures are used to depict basic structures of covalent compounds

7 Steps to Writing Lewis Dot Structures Step 1: Figure out the skeletal structure the least electronegative atom goes in the middle (the central atom ) Hydrogen and halogens will occupy end positions (only one bond will go to them) Example 1: Methane (CH 4 )

8 Steps to Writing Lewis Dot Structures Step 2: Total the number of valence electrons for all atoms

9 Steps to Writing Lewis Dot Structures Step 3: Draw a single bond connecting the atoms. For each bond you draw, subtract 2 valence electrons from your total

10 Steps to Writing Lewis Dot Structures Step 4: Use the remaining electrons to complete octets. Remember, hydrogen only needs 2 ve - to have a full outer energy level -- a single bond to H is enough! Example 2: Ammonia (NH 3 )

11 Steps to Writing Lewis Dot Structures Step 5: Check for octets. If every atom now has an octet. You re done. If not, go to step 6. Octets! Octets!

12 Steps to Writing Lewis Dot Structures Step 6: Use double or triple bonds to complete octets for any atoms that don t have them. Example 3: Carbon Dioxide (CO 2 )

13 Practice! 1. CF 4 2. Cl 2 Try each of these, then compare your structure with your lab partner s. 3. SO 2 4. N 2

14 Resonance Structures

15 Try this: Draw the Lewis Dot Structure of nitrate ion (NO 3-1 ) Look at your structure. Could it have been drawn another way?

16 A slight problem... The original Lewis Dot structure that you drew for nitrate ion is not entirely correct The correct Lewis Dot structure for nitrate ion can only be achieved by a supposition of all three seemingly correct structures Nitrate ion, however, only has one correct structure it is an AVERAGE of all three bond lengths and strengths

17 Resonance Resonance occurs when more than one valid LDS exists for a molecule A correct structure for a molecule with resonance is an average of all of the bond lengths and angles The electrons are actually not locked into positions they are delocalized, they move all around the molecule

18 More Practice: 1. Ozone (O 3 ) 2. SO 3 and SO 3-2 (sulfur trioxide vs. sulfite ion)

19 Exceptions to the Octet Rule

20 Octet Rule Basics A full outer energy level (valence shell) will cause most atoms to be stable Most atoms need 8 valence electrons to be stable H and He only have one energy level, and can only hold 2 valence electrons total

21 General Rules C, N, O, and F will obey the octet rule Second row elements NEVER exceed the octet rule

22 Exceptions B and Be often have fewer than 8 electrons around them very reactive (electrondeficient) 3 rd row and heavier elements sometimes exceed the octet rule by using their d orbitals

23 What should you do... In the Lewis structure, satisfy the octet first. If electrons are left over, put them on elements that have d orbitals

24 What should you do... If more than one element could accept the extra electrons, assume the central element wants them

25 Even More Practice: 1. Boron trifluoride 2. Phosphorus pentachloride 3. Chlorine trifluoride 4. Xenon trioxide 5. Beryllium dichloride 6. Iodine tetrachloride