Tuesday, November 9, 2010

Lewis Structures – November 9th 2010

      When we first walked into class today, we picked up pages 9-15. Then, Mrs. M stamped our page #6 (the homework assigned yesterday). Once everyone was settled again, we went over all the answers to page 6. These answers can be found on Moodle!
                  WARNING!: Tomorrow is the beginning of a streak of 6 quizzes! The quiz tomorrow will be seeing if you can tell the difference between and Ionic and Covalent compound, and if you can name each of them.                  
Makes compounds
Transfers electrons
Made with metals and nonmetals
Positive and negative charges
Weak bonds!!
Ex: NaCl
Makes molecules
Shares electrons to be noble
Made of ONLY nonmetals
Neutral charges
Strong bonds!!!
Ex: H₂O

                  Next, we did pages 7, 8, 10, & 11. You may read page 9 at your own leisure, as it only describes how to create a Lewis structure. Mrs. M knows 2 ways to figure out how to create Lewis Structures. She began to teach us the first way, (described on page 9) and the taught us “her way”. On our periodic tables, we added the amount of valance electrons. Under the first family we put 1, indicating there is 1 valance electron on all the above atoms. 2nd family we put 2, 13th we put 3, 14th we put 4, 15th we put 5, 16th we put 6, 17th we put 7, and 18th we put 8, indicating that these atoms are stable.


                  -Consult the molecular formula and sum up all the valence electrons from the separate atoms.
                  - Choose central atom
                  - The first element in the compound becomes the central atom (excluding hydrogen).
                  -Insert pairs of electrons between all pairs of atoms that are to be bonded together
                  -Place any remaining electrons on peripheral atoms as unshared pairs, starting with the most                   electronegative such atom.

Mrs. M began teaching us “her way” of drawing Lewis Structures. This method includes NHS (Needed; Have; and Shared electrons.)
                  First you take the amount of atoms in the compound (excluding Hydrogen) and multiply                   it by 8 electrons which results in the amount of electrons NEEDED to have a stable compound.
                  Then, you fine the amount of valance electrons you have. Look at the charge of the atom that we listed at the bottom of the periodic table and add them all up to equal the amount of electrons you have in that particular compound.
                  Lastly, you subtract the amount of electrons you have from the amount needed. This resulting number is the amount of electrons that need to be shared between the atoms.

FOR EXAMPLE: (done in class)
                  Take the compound NF₃. The amount of atoms in this compound is 3. 1 Nitrogen and 3 Fluorine’s. You multiply 4 by 8 to get the amount of electrons needed to create a stable compound.

4 x 8 = 32

                  Next you find the amount of electrons you have. There are 5 valance electrons in Nitrogen, and 7 in Fluorine. Remember, there are 3 Fluorine atoms!!! Add all the valance electrons together to get the amount of electrons there are.

5 + 7 + 7 + 7 = 26

                  Finally, you would subtract 26 from 32 to find the shared electrons.

32 – 26 = 6 electrons

                  Since N is the first atom used, it would become the central atom. You write N in the middle of the space provided. There are 3 fluorine atoms so you draw 3 F’s on 3 sides of the N. Because there are 6 electrons being shared, you draw a line from the N to each F. Each atom needs 8 total electrons to become stable. To show the addition of the last 2 atoms to N, you would put 2 dots on the last side unoccupied by the lines. Finally, you need to make each F atom stable. You add 2 dots on each side left on the F, to create 8 electrons on each atom.

                  HOMEWORK! Finish the 2nd column of page 11.

Annika S

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