1. Key Terms for your Glossary Octet Rule Isoelectronic Lattice

2. Why do atoms form bonds?

4. The Octet Rule Atoms bond in order to achieve an electron configuration (EC) that is the __________________________ _______ When two atoms or ions have the same EC, they are said to be __________ of each other. Eg. Cl- is __________ with Ar

5. Calculating EN differences The first step in defining the polarity of a bond is to calculate electronegativity difference (  EN)  EN = EN large - EN small E.g. for NaCl,  EN = 2.9 - 1.0 = 1.9 Next, estimate from fig 7.12 the % ionic character: about 65% (60 - 70%) Q-Give the % ionic character for MgO, CH, HCl MgO = 3.5 - 1.3 = 2.2 … 80% ionic (75-85) CH = 2.5 - 2.1 = 0.4 … 7% ionic (5 - 10) HCl = 2.9 - 2.1 = 0.8 … 20 % ionic (15-25) Note if % ionic is 20%, then % covalent is 80%

6. Ionic Bonding NaCl = , therefore ________ Na transfers its valence e to Cl = Na + Cl - How does this reflect the ____________ __________________________________ _______________________ Cl gains an e, and becomes anion Cl -, becomes isoelectronic with Ar http://www.youtube.com/watch?v=9pHOiJRBXRc

8. Lewis Structure for Ionic Bonds Tips for Lewis structures dealing with bonds: use two different symbols. For example: o and x’s, or open o’s and filled in o’s

9. Transferring Multiple Electrons What about Mg and O? Electronegativity = 3.4-1.3 = 2.1= ______ Mg = 2 valence e, O = 6 valence e What do they need to do to become isoelectronic with a noble gas? __________________________________ __________ Try drawing a Lewis Structure for Magnesium Oxide.

11. Practice! For each bond, determine  EN, type of bond, and the Lewis Structure: Ca – O K – Cl Li – Br Ba - O

12. Ionic Bonding with MORE than 2 ions Consider Ca and F EN = 4.0-1.0 = 3.0 = ________ __________________________________ So, Ca will form Ca 2+ and F will form F - How do you think this will happen?

13. Calcium Fluoride, CaF 2 Ca bonds with 2 F atoms Ca 2+ + 2F = ____________

14. Practice! For each bond, determine DEN, type of bond, and the Lewis Structure: Mg – Cl Na – O Ca - Br

15. Conductivity of Ionic Compounds What happened in our lab?

17. Conductivity of Ionic Compounds Why are ionic compounds good conductors when dissolved and not in their solid state? What is required for electrical conductivity? What is the structure of ionic compounds in the liquid, solid, and dissolved states?

18. Conductivity of Ionic Compounds An electrical current can flow ______________________________________ _________________________ Is there a mobile charge in solid NaCl? Salt crystals are held together by rigid strong ionic bonds, in a ____________. __________-, not Ladas :O) In a solid state, ions can’t move very much. There is no mobile charge, that’s why solid NaCl _______________________

20. What do you think? What happens when we dissolve NaCl in H20?

22. Aqueous NaCl as a conductor When dissolved, the lattice __________ Ions are free to move around = __________________ ______ Salt water is a __________________ See Figure 3.14 pp.79

23. Stalagmites and Stalactites Crystal columns that form when water containing dissolved lime (CaO) drips down. From limestone (CaCO3) Clear ionic solution over a crystal of the same compound = will continue the arrangement of the crystal

24. Covalent Bonds – Key Terms Diatomic elements Double bond Pure covalent bond Triple Bond

25. Covalent Bonds What happens when the  EN is very small? _________________ (shared e) What happens when the  EN is 0? __________________________________ __________________________________ _______________________________ This happens in ___________________ http://www.youtube.com/watch?v=UR4eG60jjQQ

26. Diatomic elements I Bring Clay For Our New House Have No Fear of Ice Cold Beer ___________________________________ Only 7 elements that exist as pure covalent diatomic elements There are numerous other molecules that diatomic, but not with the same element Noble gases do not form diatomic elements

28. Carbon – The Master Bonder  between C and H = _________________________ ________really small difference, almost __________ Eg. Methane CH 4 Each H shares 1 e with C, C shares each of it’s 4 e with H = _________________________ ___________________ When checking your Lewis Diagrams, count the shared e as If they belong to each of the bonding atoms.

29. Practice! Show how each pair of elements forms a covalent compound, and give the compound formula. H and O Cl and O N and H I and H

30. Multiple Covalent Bonds Some atoms need to share more than one pair of electrons. They may need to share 1, 2, or 3 pairs. Eg. Diatomic molecule of _______________

31. Double and Triple Covalent Bonds These bonds can form between _________________ _____. Eg. Carbon Dioxide (CO 2 ) * Tip* - make sure that each element has 8 e by drawing circles around them in different colours.

32. Double and Triple Covalent Bonds Atoms that share three pairs of e form _________________________ Eg. Diatomic N

33. Practice! CS 2 C-H-N CH 2

34. Polar Covalent Bonds “Monkey in the Middle”

35. Polar Covalent Bond  EN the is _______________ Although higher _____________________________, the  is not high enough to make an ionic bond.  is high enough to make the bonding pair want to be __________________________________ ___________

36. -/+ Partial Charges (      Eg. O – H EN of O = __________ _____________________  = _______ Therefore between 0.5 and 1.7 O has a higher EN so it attracts e more strongly and therefore has a slightly more ______________________ Also known as _________________________ H has a __________________

37. What About Water H 2 0? What do you think the partial charges would be?

39. Water vs. Carbon Dioxide

40. Predicting Molecular Shape Why is H 2 O bent and CO 2 straight? Look at the Lewis Structures e- pairs that are not involved in bonding are called _____________ e- pairs are that involved in bonding are called ______________ e- pairs are arranged around molecules so that they have a ____________________________________

41. Shape of a H 2 0 Molecule

42. Water Water is a polar molecule, as it has two negatively charged parts (the electron pairs that are unbound) and two positive parts (the Hydrogen atoms). These opposite charges are attracted to each other. This holds water molecules together which explains why water is a liquid at room temperature while almost all other similar sized molecules are gases. This also accounts for water's excellent dissolving capacity for other charged substances such as salt, and why uncharged substances (non polar) such as oil do not mix readily with water.

43. Shape of a CO 2 Molecule Each double bond uses 2 bond pairs - which are then thought of as a single unit. Those two double bond units will try to get as far apart as possible, and so the molecule is linear. The structure we've drawn above does in fact represent the shape of the molecule.

44. Other Examples What about ammonia (NH 3 )? Or methane (CH 4 )?

47. LOTS of new stuff tonight!! 2 nd half of lesson #3 on Monday (polarity of bonds) We will be taking up the test Monday We will be doing an activity with molecular building kits as well. Please if you haven’t done the readings…..DO THEM! Make sure that all homework for Lessons 1,2,3 is complete so that you can benefit and understand the activity.