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chapter 12 chemical bonding

chapter 12 chemical bonding. Essentially everything we see around us, natural or man-made, compounds or elements, has atoms bonded with atoms Hard or soft, solid, liquid, or gas; these properties are determined by bonding

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chapter 12 chemical bonding

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  1. chapter 12chemical bonding

  2. Essentially everything we see around us, natural or man-made, compounds or elements, has atoms bonded with atoms • Hard or soft, solid, liquid, or gas; these properties are determined by bonding • Their structures play the major role in chemical rxns all around us, and… • The way they’re bonded determines structure

  3. 12.1 types of chemical bonds • Bond: a force that holds groups of two or more atoms together and makes them function as a unit • They can do this in several ways! ready?

  4. Remember table salt? NaCl • When the water in salt water evaporates away, the Na+ get together with the Cl- to form NaCl • This closely packed collection of oppositely charged ions is an example of Ionic bonding

  5. How strong are ionic bonds judging by this?

  6. These substances are called ionic compounds, and they can be made like this…

  7. Remember: metals like to lose e-s… • and nonmetals gain • So there can be a transferhere of electrons • and the result is two things that love each other… :)

  8. M + X  M+X- • M = metal, X = nonmetal, M+X- is the ionic compound (e.g. NaCl, MgO, KBr…) • but what if the atoms are identical? how do they bond together?…

  9. When these two atoms get close, the e- of one are attracted to the p+ of the other; the e- (or p +) of one are repelled by the e- (or p +) of the other

  10. If they are most comfortable at a distance where their e- shells overlap! - we have a bond!

  11. mini-summary • This is a covalent bondbecause outer shell (valence) electrons are shared by both atoms

  12. Bigger-than-mini summary

  13. Ionic bonding (transferring e-s) and covalent bonding (sharing e-s) are the extremes… • Between these 2 extremes are…

  14. Polar covalent bonds! • Unequal sharing of electrons • Here one atom is not strong enough to strip the electron off the other, but it can hog it • The hog gets a ∂- to showit is almost (not quite) anegative ion; the losergets a ∂+ • but how much do theseatoms like electrons???

  15. Nonpolar Covalent: • Equal Sharing of electrons

  16. 12.2 Electronegativity • When different nonmetals form a covalent bond, they almost never share the electrons equally • = Electronegativity • = the relative ability of an atom in a molecule to attract electrons to itself

  17. Here are electronegativity values • See a general pattern?

  18. Fluorine rules the electonegativity table! The numbers generally decrease as you move away from him

  19. The polarity of a bond depends on the two battling for the electrons • In fact, the difference in electronegativityis the Big Deal here: • great differences in electronegativity = in more ionic sort of bonds • small differences = polar covalent bonds • no difference= (nonpolar) covalent

  20. Where ∆ electronegativity (en) means the difference in electronegativity values: ∆en = 0 covalent ∆en = >0  ~2 polar covalent ∆en = >~2 ionic e.g. here is a polar covalent situation

  21. Nonpolar Covalent Ionic Polar Covalent

  22. example • Classify these as I, PC, or NPC: • KF • O2 • ICl • HBr • NaF • N2  = 3.2 I  = 0 NPC  = 0.5 PC  = 0.7 PC  = 3.1 I  = 0 NPC

  23. example 12.1 • Organize these from least polar to most • H-H • O-H • Cl-H • S-H • F-H  = 0  = 1.4  = 0.9  = 0.4  = 1.9 1 4 3 2 5

  24. 12.3 Bond polarity and dipole moments • If the molecule actually ends up having a partial positive side and a partial negative side we say it has a dipole moment • an arrow tells uswhere thenegative side is

  25. It always happens in diatomic molecules with unequal sharing; it can also happen in polyatomic molecules • like water:

  26. This unequal distribution of charge in a water molecule is unbelievably important to life

  27. it can surround and dissolve ions (more on that later)

  28. they can attach to each other and stick (more later) • which is why it is a liquid and a solid on earth

  29. 12.4 stable electron configurations and charges on ions • a Big Picture so far (see a pattern?)

  30. we can summarize observations like this with a modified PT:

  31. after this and lots more observations we can say this: in almost all stable compounds of the representative elements, all of the atoms have achieved a noble gas configuration • there is just something special about getting a full outer shell!

  32. so when Metals react with Non Metals they’ll exchange e-s until they look NOBLE! • [important note: they don’t become noble gases, they just have an electron config that looks like one] • When Non Metals react with Non Metals they’ll share until they look noble

  33. Predicting formulas of ionic compounds • And now, a Chemistry A review with more depth • In the old days we saw that Ca ion is +2, and O has a -2, to form CaO • Now we know Ca with [Ar]4s2 will lose two electrons (to look noble) • thus Ca2+, and…

  34. O is [He]2s22p4 so it needs two more to look noble • It can take the two electrons from Ca, and now be O2- • so both get a noble appearance • The result is Ca2+ and O2- to form CaO • got it!?

  35. So what happens when Al metal is reacted with oxygen? • Al has its 3 valence electrons, and O only needs 2 • Two Al atoms can provide enough electrons for three O’s (6 total) • so, Al2O3

  36. This is how they all look when they achieve “nobility;” look familiar?

  37. 12.5 Ionic Bonding and Structures of Ionic Compounds • When we learned that sodium chloride is “NaCl” we really were talking empirical (simplest) formula • In reality there are b-zillions of Na’s packed together with b-zillions of Cl’s in a 1:1 ratio

  38. Here is thestructurefor LiF • Again, the “LiF”just tells us theratio of thesetwo is 1:1 • They are closely packed and tough to move around (making it brittle)

  39. it’s easier to see the structure’s geometric shape with the wire frame model

  40. You may havenoticed that thepositive ions arealways dinkierthan the negativeion; why? • Remember the Metal loses its whole outer shell to form an ion; and the Non Metal gains electrons to look noble (getting bigger as a result)…

  41. see an obvious pattern?

  42. Ionic Compounds Containing Polyatomic Ions • [This always causes a bit of confusion so pay attention here] • Polyatomic ions(like NH4+ or NO3-) are true ions with real charges and can get together with other “regular” ions, but! • They are covalently bonded within themselves!

  43. The polyatomic ion acts like a single, welded-together ion • So when ammonium nitrate (NH4NO3) dissolves in water it splits into NH4+and NO3-only; each of the ions stays intact, like a charged family • OK?

  44. 12.6 lewis structures • Remember: bonding involves the valence electrons of atoms!!! • Valence are transferred in ionic bonding and shared in covalent • The Lewis structureis just a simple representation to show how the valence e-s are distributed around a molecule

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