3. Wednesday, Feb. 19 th : “A” Day Thursday, Feb. 20 th : “B” Day Agenda  Collect homework: pg. 198: 1-14  Collect lab: “Nonmetal Reaction”  Sec. 6.1 quiz: “Covalent Bonds”  Begin Section 6.2: “Drawing/Naming Molecules” Valence electron, Lewis structure, unshared pair, single bond  In-Class: Practice pg. 202: #1,2 Practice pg. 203: #1,2  Homework: Concept Review

4. Homework Pg. 198: #1-14 Lab: “Non-metal Reaction”  Questions, problems?  Turn in

5. Section 6.1: Quiz “Covalent Bonds”  You may use your notes, your book, and your 10:00 lab partner on the quiz….  Question #2 IS NOT in section 6.1. The answer is “B”“B”  When you’re done, pick up the notes for today…

6. Sec. 6.2: “Drawing and Naming Molecules”  Both ionic and covalent bonds involve valence electrons, the electrons found in the outermost energy level of an atom that determine the atom’s chemical properties.

7. Determining Number of Valence Electrons Review  How many valence electrons are in the following elements? Ca P I C Li

8. Lewis Electron-Dot Structures  In 1920, G.N. Lewis, an American chemist, came up with a system to represent the valence electrons of an atom.  This system, known as electron-dot diagrams or Lewis structures, uses dots to represent valence electrons.

9. Lewis Structures Model Covalently Bonded Molecules  Lewis structures: a structural formula in which electrons are represented by dots. Dot pairs or dashes between two atomic symbols represent shared electrons in covalent bonds.  A Lewis structure shows only the valence electrons, NOT all of the electrons in an atom. Na Ca

10. Lewis Structures Show Valence Electrons  As you move from left to right across a period, each element has an additional valence electron.  These valence electrons are represented by dots on each side of the element’s symbol.

11. Lewis Structures Show Valence Electrons  You do not begin to pair dots (valence electrons) until all four sides of the element’s symbol have a dot.

12. Octet Rule  An element with an octet of valence electrons, such as a noble gas, has a stable configuration.  The tendency of bonded atoms to have octets of valence electrons is called the octet rule.

13. Lewis Structures Show Valence Electrons  When two chlorine atoms form a covalent bond, each atom contributes one electron to form a shared pair.  With this shared pair, both atoms can have a stable octet.  Unshared pair: a non-bonding pair of electrons in the valence shell of an atom; also called a lone pair.

14. Lewis Structures Show Valence Electrons  Single bond: a covalent bond in which two atoms share one pair of electrons.  The shared pair of electrons are replaced with a dash, indicating a single bond. (A dash counts for 2 electrons)

15. Rules for Drawing Lewis Structures Pg. 201 There are 5 steps to follow to draw a Lewis structure: 1.Determine the number of valence electrons in each atom. Add them up to get the total number of valence electrons in the molecule.

16. 2. Arrange the atoms Hydrogen atoms can only form 1 bond! Halogen atoms often bond to only one other atom and usually at the end of the molecule. Carbon wants to be the center of attention and is often placed in the center of the molecule. With the exception of carbon, the atom with the lowest electronegativity is often the central atom. Molecules also like to by symmetrical. Rules for Drawing Lewis Structures Pg. 201

17. Rules for Drawing Lewis Structures Pg 201. 3.Distribute the dots (valence electrons) Arrange the dots so that each atom (except for hydrogen, beryllium, and boron) satisfies the octet rule (8 dots). 4.Draw the bonds Change each pair of dots that represents a shared pair of electrons to a long dash. 5.Verify the structure Count the number of electrons around each atom. Make sure the octet rule is satisfied and that you have the correct amount of dots that you determined in Step 1.

18. Sample Problem A, Pg. 202  Draw a Lewis structure for CH 3 I. 1.Determine the total number of valence electrons in each atom and in the molecule. 4 + 1 + 1 + 1 + 7 = total number of dots 14 2. Arrange the atoms. Put carbon in the center with the other atoms around it.

19. Sample Problem A, Pg. 202 3.Distribute the dots. Put one shared pair of electrons between each of the bonded atoms. Distribute the remaining electrons, in pairs, around the remaining atoms to form an octet for each atom.

20. Sample Problem A, Pg. 202 4.Change each pair of dots that represents a shared pair of electrons to a long dash. H H C I H 5.Verify the structure! Count the electrons and make sure you have the correct amount. (Each dash counts for 2 electrons.)

21. Additional Examples Draw the Lewis structures for the following molecules: 1. SCl 2 2. AsF 3 3. SiH 4 4. CHF 3

22. Lewis Structures for Polyatomic Ions  Lewis structures can be drawn for polyatomic ions as well.  If the charge on the polyatomic ion is negative, you must add electrons.  If the charge on the polyatomic ion is positive, you must subtract electrons.  The entire Lewis structure is put in [brackets] to show that the charge of the ion is distributed over all of the atoms. [ ]

23. Sample Problem B, Pg. 203  Draw a Lewis structure for the sulfate ion, SO 4 2- 1. Determine the total number of valence electrons in each atom and in the molecule. Add two additional electrons to account for the 2 − charge on the ion. 6 + 6 + 6 + 6 + 6 = 30 valence electrons Plus 2 extra (for the 2 - charge) = 32 total

24. Sample Problem B, Pg. 203 2. and 3. Arrange the atoms and distribute the 32 dots so that there are 8 dots around each atom. (Sulfur has the lower electronegativity, so it’s placed in the center, surrounded by the 4 oxygen atoms. Plus, this is symmetrical.)

25. Sample Problem B, Pg. 203 4.Draw the bonds. Change each shared pair of electrons to a long dash. Place brackets around the whole structure and a 2  charge outside the bracket to show that the charge is spread out over the entire ion.

26. Additional Examples Draw the Lewis structures for the following polyatomic ions: 1.OH - 2.NH 4 + 3.H 3 O + 4. ClO -

27. In-Class  Practice, pg. 202: #1,2  Practice, pg. 203: #1,2