 Worksheet from side counter FIND YOUR NEW SEAT

 Pass back work  Extra credit slips: pass out  Presentation grades posted 3 rd period: finish presentations  Tests will be graded by the end of Tuesday Go over Wednesday  Work ethic

 Demo intro to reactions  Elements vs. compounds  Vocab cards  Practice  Kahoot check  Valence electrons: revisit Objective: how do elements make up compounds?

 Volunteer to keep track of elements and compounds demonstrated in the demo

 Work on part 1 of your worksheet ElementCompound

 15 minutes to finish practice problems with your neighbor

 Go to kahoot.it on your iPad/laptop

 How do you find the number of valence electrons an element has by looking at the periodic table?

 How many valence electrons in Carbon?

 How many valence electrons in Nitrogen?

 How many valence electrons in Beryllium?

 How many valence electrons in Sodium?

 How many valence electrons in Argon?

 Back of worksheet  Finish by Wednesday (more HW will be assigned Wed, collect on Th/F)

 ½ sheet of paper for bellwork quiz  Homework (compound vs. element)  Write name and date 10/29/2014

1. Is Xe an element, compound or diatomic element? 2. How many elements make up the compound Mg(SO 4 ) 2 ? 3. How many Sulfur atoms are there in the compound Mg(SO 4 ) 2 ?

1. Is Xe an element, compound or diatomic element? 2. How many elements make up the compound Mg(SO 4 ) 2 ? 3. How many Sulfur atoms are there in the compound Mg(SO 4 ) 2 ? 4. How many Carbon atoms are there in the compound CH 3 COOCH 2 C 6 H 5 ?

 How do we go from elements on the periodic table to building compounds?  We must first visit Lewis Dot Structures of single elements to determine how they will fit together.

 Look at the two pictures below. One is a Lewis Dot Structure of Sulfur and one is the Bohr model of Sulfur. How are they similar and how are they different?

 Each atom has 4 sides (top, bottom, left, right). Only 2 electrons can fit on each side.  Place electrons 1 at a time clockwise or counterclockwise on each side before doubling up  No more than 2 electrons can sit on one side of the atom Exceptions: Hydrogen and Helium only have 1 orbital  These rules make it clear later on how atoms bond together

 Create Lewis Structures using beans  With your partner: Create the Lewis structure for the elements on your note-taker using knowledge of valence electrons and the proper location of Lewis Dot electrons Once you have created it, draw it on your note- taker in the box indicated

 ½ sheet of paper for bellwork quiz  Write name and date

 Draw the Lewis dot structure of Iodine (I)  How many electrons does Iodine need to have a full outer shell?  How many electrons does Hydrogen need to have a full outer shell?

 What are the rules for how do Lewis Dot Structures bond together to make compounds?

 IN COMPOUNDS ELECTRONS COME IN PAIRS!  Only in rare circumstances do you see a lonely electron on an atom in a compound

Electronegativity – the atoms desire for electrons or attraction to electrons Compare the following elements in terms of electronegativity: H vs. O C vs. O C vs. F B vs. F

1. LEAST electronegative atoms usually go in the middle 2. More electronegative atoms like to have electron pairs un-bonded (floating around them 3. Hydrogen’s are always external (not central atoms) 4. Each atom must have a satisfied octet (8 valence electrons, 2 for hydrogen)

 Octet rule: atoms of main group elements tend to bond in such a way that each atom has 8 electrons in its valence shell.  Bonds: each bond is a shared pair of electrons and each atom associated with that bond contains both electrons in its valence shell.

1. Formal Charge = The charge on an atom based on its bonding configuration in a compound 2. Formal Charge = V – B 2 V = valence electrons of element in question B = bonded electrons of element in question Ex. C 2 H 2 Element: Carbon V = 4B = 8 = 4 2 FC = 4 – 4 = 0

Goal:  The structure of the compound must have the lowest possible formal charges on all atoms to be stable  If formal charges are high, the compound likely does not exist in that form.

 Practice – with beans  Create the following compounds with your beans and then draw them on your note-taker in the box provided  H 2 OCO 2

 What would be a good definition for a bond after creating these structures?  How many electrons would be required to make a double bond? A triple bond?

 Task: Step 1: add up all valence electrons for each atom in the compound Step 2: draw the Lewis Structure of the compound with electrons ONLY Step 3: draw the Lewis Structure of the compound by substituting bonds for electrons between atoms (2 electron = 1 bond, 4 electrons = 2 bonds, 6 electrons = 3 bonds)

 Task: Step 1: add up all valence electrons for each atom in the compound Step 2: draw the Lewis Structure of the compound with electrons ONLY Step 3: draw the Lewis Structure of the compound by substituting bonds for electrons between atoms (2 electron = 1 bond, 4 electrons = 2 bonds, 6 electrons = 3 bonds)

 Use the beans each time you create a compound  If you need help with one, I cannot help you unless you have started with the beans so we can manipulate them

 Covalent bond: Shared pair of electrons Can be equal or unequal sharing  polarity  Ionic bond: Loss of gain of electrons in creating a bond Metal + non-metal (positive ion + negative ion)

 More Lewis Dot Structure  Honors: formal charge with each atom in each compound!