#  Ionic bonds › Strongest (melting point?) › Charged ions  Metallic bonds › Strong, but weaker than ionic › “Sea of electrons” – all shared among all.

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 Ionic bonds › Strongest (melting point?) › Charged ions  Metallic bonds › Strong, but weaker than ionic › “Sea of electrons” – all shared among all of them  Covalent bonds › Weakest (melting point?) › Probably the most important substances of all are molecules

 When you need to draw a Lewis structure, always list the number of electrons you have first. This is especially important for when you have polyatomic ions (charged particles).  SO 4 2- is a polyatomic ion – what would its Lewis structure be?

 First, how many electrons? › S  6 › O  6 x 4 = 24 › Charge  2 extra e - › TOTAL  6 + 24 + 2 = 32  Place S in the center, surrounded by the 4 O’s. Make a single bond first on each one. Then place electrons until you have a model that obeys the octet rule.

 Let’s try one more for ammonium. › N  5 › H  1 x 4 = 4 › Charge  -1 e - (a positive charge means we’re short electrons) › Total  5 + 4 – 1 = 8

 Remember we can use electronegativity to calculate the polarity of a bond?  We find the electronegativity values for each atom in the bond on the periodic table (p. 174)  Then we subtract the smaller number from the bigger one – the difference tells us how polar the bond is.  0.0 ------- 0.5 ----------- 1.7 ------------ 3.3  nonpolar polar mostly ionic

 Consider the compound HCl. Is this bond nonpolar covalent, polar covalent, or mostly ionic? › Hydrogen’s electronegativity is 2.2 › Chlorine’s electronegativity is 3.19 › What is the difference?  What about CCl?

 Has anyone seen the movie Casino Royale ? Know the name of the main girl?  Vesper.  Now I have to ask: what are the odds that a scriptwriter would just happen to put a guy named Bond with a girl named Vesper? Clearly that writer knew their chemistry.

 Say that 5 times fast!  What does that name tell you about the theory?  It has to do with valence shell electrons, it seems they come in pairs, and they repel each other.  Why would electrons repel?

 VSEPR theory only applies to molecules (substances with only covalent bonds).  Its difficult to see VSEPR shapes on paper, so I’ll be showing you some shapes with a modeling kit.

 In a molecule, there are two types of electron pairs: bonding pairs (those used to form bonds), and lone pairs (those not used to form bonds)  For counting purposes, double bonds and triple bonds count as a single electron pair.  Each electron pair wants to be as far away from all the others as possible. As you’ll see, this results in some very common shapes.

 Electronegativity can tell us more than just the polarity of a bond. It can tell us if a molecule will be polar overall.  First, consider CH 4. What is the geometry? What is the polarity of each bond in the molecule? Is the molecule polar?  Look at CH 3 Cl. Calculate the polarity of each bond in the molecule. Is the molecule polar?  Last one: CCl 4. Are the bonds polar? Is the molecule polar?

 No! If all the polar bonds in a molecule cancel out, the molecule will be nonpolar.

 Determine the geometry of each of the following molecules, then determine if the molecules are polar or nonpolar overall. (Remember, a compound can have polar bonds and still be nonpolar!) › H 2 O XeF 2 › SO 2 NH 3 › BF 3 › (the ones in bold don’t obey the octet rule for the central atom)

 P. 181 (4, 6, 8, 10, 14, 20)  P. 225 (1, 5, 7, 10, 12, 14b,c,e, 19)  This assignment will be due the day before the unit test – you should have at least a week to finish it.

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