Chemistry SM-1131 Week 15 Lesson 1

1. Chemistry SM-1131Week 15 Lesson 1 Dr. Jesse Reich Assistant Professor of Chemistry Massachusetts Maritime Academy Fall 2008

2. Class Today • Poem • Review Chapter 10 • Test Wednesday • Class on Friday- student evals and test returned • I’ll try to make some review sheets, but you should make your own in case I don’t • Final is on Thursday. Check the schedule to find out where you are supposed to meet and be.

3. Poem • “Follow your true north.”

4. Bonding theories • How bonding works. What hold molecules together. • Ionic bonding- metal and a non-metal- one element loses an electron, gains a positive charge, and the other gains an electron and a negative charge. They then become “electrostatically attracted.” • Covalent- two non-metals- share electrons to obey the octet rule. Form a “chemical bond.”

5. Lewis Theory • Valence electrons are dots • Every main group element wants 0 or 8 valence electrons around it to be more like a noble gas. (except H which wants 0 or 2 and He which wants 2).

6. Figuring out Valence • Count up all the S and P electrons in the highest shell. • Or, look at group number IA=1, IIA=2, IIIA=3, IVA=4, VA=5, VIA=6, VIIA=7, VIIIA=8

7. Lewis Dot Theory

8. Adding electron dots • Normally, I want you to fill out the pattern like we did above going around in a clockwise or counterclockwise fashion. • When it’s time to pair electron pair them away from other atoms. • If you have a central atom and only two side atoms. On the sides pair one up on top and the other on bottom.

9. Ions • So Lithium has 1 valence Electron. If it loses that electron it has a +1 charge and 0 valence electrons. We’d write it like this: Li+ and the Lewis Dot would be this: Li

10. Ions • Oxygen has 6 valence electrons. It wants eight. So, typically it gains two electrons. We’ll write it like O2-, and it’s dot structure would be the following. O

11. Ionic Compounds Li +1 O -2 Li2O Li O Li What happens if the two Li each lose 1 electron? How many VE will each Li have? The O? Is the Octet Rule satisfied?

12. Li2O Li O Li

13. Ionic Compounds • MgO • Mg+2O-2 • Mg2O2 • MgO Mg O Mg has 0 VE and O has 8 VE. Octet Rule! Don’t draw the arrows!

14. Ionic Bonding • Notice in all those cases it was between a metal and a non-metal. The metals lost electrons. The non-metals gained them. Charges were formed. Oppositely charged particle were attracted to each other. That’s how ionic bonding works!

15. Covalent Bonding • Now, these example are two non-metals. They will share electrons to fulfill the octet rule. When we draw the line by connecting the dots it’s called a covalent bond. It just means two elements are sharing electrons.

16. H2 • H has 1 valence electron. It wants either 0 or 2. So far, we’ve seen it give or take an electron to get to 0 or 2 valence electrons. This time it’s going to do something different. It’s going to share. H H we draw it this like: H H

17. I2 • I has 7 valence electrons. It wants either 0 or 8. It’s going to share. I I we draw it this like: I – I I - I

18. Bonds have 2 electrons • Because they share electrons when we count how many electrons each one has we say both get the benefit of the two electrons in the bond. • H-H each one has 2 electrons • I-I each one has 8 electrons, each I has 3 lone pairs for 6 electrons, and then we say each one has 2 electrons from the shared bond

19. Multiple Bonds • O2 has a double bond. • N2 has a triple bond. • I’ll show you why on the board.

20. Dot Placing Exceptions • Every now and again you hit an exception. • SO2, which was on your packet was an exception. I had forgotten. • If you fill out dots like normal you end up with an unpaired electron on both oxygen atoms. You’d think there should be a way to make a bond somehow. • If you think something should work, but it didn’t, try pairing dots on the central atom facing one of the adjacent elements.

21. Lewis Dot exceptions • PF5 (P has 10 Ve), SF6 (S has 12 Ve), BF3 (B has Ve), SO4-2 (S has 12 Ve). PO4-3(P has 10 ve), NO (N has 7 Ve) • You don’t have to memorize these, but if you see them or anything like them you have to know why it’s an exception to the octet rule.

22. Resonance • Sometimes molecules have 2 or more ways to write the exact same thing, but now a double bond is on the left rather than the right and so on. • These structures are called resonance structures. • Sometimes a molecule exists as one, then the other, and sometimes it’s in between. • Electron sloshing.

23. VSEPR • Linear- straight line • Trigonal Planer- connect outside atoms to make a triangle • Tetrahedral- 4 corners of a cube- 2 in the plane and 2 out of the plane • Trigonal pyramidal- looks like a pyramid • Bent- Lewis looks straight, but lone pair(s) make these bent.

24. VSEPR • Linear (central atom with a triple bond or 2 double bonds) • Trigonal planer (three terminal atoms, no lone pairs, this is the only one where the atom isn’t trying to look tetrahedral) • Tetrahedral (four terminal atoms) • Trigonalpryamidal (three terminal atoms and one lone pair) • Bent (1 double bond, a single bond, 1 terminal atom or 2 single bonds and two lone pairs)

25. 3D • When you draw atoms in the plane of the paper just use normal lines for bonds. • If you want to draw something coming towards you draw a wedge as a bond • If you want to draw something going away from you draw a hashed line as a bond.

26. Electronegativity • Many of you boneheads didn’t graph this • You won’t know the trends as well • I’ll probably ask for this on the test • It relates back to shielding, atomic radius, and ionization energy (all of which you should have graphed.

27. Shielding • As you go across the period the charge increase, but the shielding does not. • The nucleus is getting more and more positive, and therefore electrons are getting drawn closer and closer to it. • That is until the next period is reached, and then the next electron is pushed much further out by a whole new set of elements shielding it from the nucleus.

28. Atomic Radius • Radius shrinks as you go from left to right because the effective nuclear charge (how much the electron actually feels a pull towards the center) increases • As you go down the periodic table each new period gets a new set of shielding and therefore the electrons are further away from the nucleus and the radius expands.

29. Ionization energy • To ionize means to make an ion. We typically mean ripping out an electron. • The electron that is the mostly weakly bound is the one on the far left. If feels the weakest force holding it in place. That atom is easy to let go. • Across the period the effective nuclear charge is increasing and its pulling electrons closer into it. As you move right it’s harder to dislodge an electron. • As you go down there is more shielding in the way making it easier to rip out an electron.

30. Electronegativity • In the top right (F) you have a very strong nucleus with very little shielding. You have a very strong pull towards F. So strong it can pull neighboring atom’s electrons towards it. • In the bottom left you have an atom with an enormous amount of shielding and very little effective nuclear charge making it to the valence electron. There is very little pull towards the nucleus.

31. Tug of War • Electronegativity is like two atoms playing tug of war in the playground. In general the big kid wins and the little kid goes home crying. • The big kids are in the top right with the very positive nucleus and little shielding. • Everyone to the left is a littler kid then the ones on the right. • For covalent bonds, it’s still sharing, but it’s not an equal sharing.

32. Dipole moment • When the more electronegative element starts pulling the electron closer to it we say it kinda gets a negative charge. When the less electronegative element holds onto its electron less we say it kinda gets a positive charge. These aren’t full charges, just partial. • The electronegative elements get a partial negative, the electropositive gets a partial positive. We call the + and – a dipole.

33. BF3 • F is the most electronegative element. B is not. F steals electron density from B. Each F has a partial negative charge. The B has a partial positive.

34. CCl4 • Cl is one of the most electronegative elements. C is not. Watch what happens.

35. Water

36. Calculating Bond Type • You can calculate bond types by subtracting a less electronegative element from the more electronegative element. • If EN-EP= 0-.4 then it’s a non-polar bond where everything is shared equally (Cl-Cl) • If EN-EP=.4-2.0 then it’s polar and one is the big kid and one is the little kid (HF, CO2) • If EN-EP = 2.0+ then it’s ionic. It’s not shared, the EP gave its electron(s) to EN.

37. A difference • One of those three is not like the other. • BF3, it looks like all the arrows are pulling in symetrically opposite directions. • CCl4, it looks like all the arrows are pulling in symetricallyoppsotivedriections. • H2O, it looks like the H electrons are pulled up • All three have polar bonds, but because H electrons are only going up and aren’t balanced out the molecule is uneven.

38. Polarity • Molecules that have polar bonds and are overall will be asymetrical polar molecules.

39. Polar vs. Non-Polar • Polar molecules will dissolve in water • Non-Polar will not • Polar Molecules will dissolve salts • Non-Polar will not • Polar Molecules will be more dense and harder to boil then similar looking non-polar molecules

40. THE END • Study for test on Wednesday • Get test, posted grades, student eval on Friday • Final on Thursday (I think 2:45pm) • It’s been fun and I’m proud of those of you who worked your butts off.