1. 1 BONDING & MolecularGeometry Cocaine Chemistry I – Chapter 8 SAVE PAPER AND INK!!! When you print out the notes on PowerPoint, print "Handouts" instead of "Slides" in the print setup. Also, turn off the backgrounds (Tools>Options>Print>UNcheck "Background Printing")!

2. 2 Bonding & Geometry Problems and questions — How is a molecule or polyatomic ion held together? Why are atoms distributed at strange angles? Why are molecules not flat? Can we predict the structure? How is structure related to chemical and physical properties?

3. 3 Review of Chemical Bonds There are 2 forms of bonding:There are 2 forms of bonding: _________—complete transfer of 1 or more electrons from one atom to another (one loses, the other gains) forming oppositely charged ions that attract one another_________—complete transfer of 1 or more electrons from one atom to another (one loses, the other gains) forming oppositely charged ions that attract one another _________—some valence electrons shared between atoms_________—some valence electrons shared between atoms Most bonds are somewhere in between ionic and covalent.

4. 4 The type of bond can usually be calculated by finding the difference in electronegativity of the two atoms that are going together.

5. 5 Electronegativity Difference If the difference in electronegativities is between: – 1.7 to 4.0: Ionic – 0.3 to 1.7: Polar Covalent – 0.0 to 0.3: Non-Polar Covalent Example: NaCl Na = 0.9, Cl = 3.0 Difference is 2.1, so this is an ionic bond! HCl is polar covalent Cl 2 is n-p covalent

6. 6 Ionic Bonds All those ionic compounds were made from ionic bonds. We’ve been through this in great detail already. Positive cations and the negative anions are attracted to one another (remember the Paula Abdul Principle of Chemistry: Opposites Attract!) Therefore, ionic compounds are usually between metals and nonmetals (opposite ends of the periodic table).

7. 7 Electron Distribution in Molecules Electron distribution is depicted with Lewis (electron dot) structuresElectron distribution is depicted with Lewis (electron dot) structures This is how you decide how many atoms will bond covalently! (In ionic bonds, it was decided with charges)This is how you decide how many atoms will bond covalently! (In ionic bonds, it was decided with charges) G. N. Lewis 1875 - 1946

8. 8 Review of Valence Electrons Remember from the electron chapter that valence electrons are the electrons in the OUTERMOST energy level… that’s why we did all those electron configurations!Remember from the electron chapter that valence electrons are the electrons in the OUTERMOST energy level… that’s why we did all those electron configurations! B is 1s 2 2s 2 2p 1 ; so the outer energy level is 2, and there are 2+1 = 3 electrons in level 2. These are the valence electrons!B is 1s 2 2s 2 2p 1 ; so the outer energy level is 2, and there are 2+1 = 3 electrons in level 2. These are the valence electrons! Br is [Ar] 4s 2 3d 10 4p 5 How many valence electrons are present?Br is [Ar] 4s 2 3d 10 4p 5 How many valence electrons are present?

9. 9 Bond and Lone Pairs Valence electrons are distributed as shared or BOND PAIRS and unshared or LONE PAIRS.Valence electrons are distributed as shared or BOND PAIRS and unshared or LONE PAIRS. HCl lone pair (LP) shared or bond pair This is called a LEWIS structure.

10. 10 Bond Formation A bond can result from an overlap of atomic orbitals on neighboring atoms. Cl HH + Overlap of H (1s) and Cl (2p) Note that each atom has a single, unpaired electron.

11. 11 Steps for Building a Dot Structure Ammonia, NH 3 Decide on the central atom; never H. Why? Decide on the central atom; never H. Why? If there is a choice, the central atom is atom of lowest affinity for electrons. (Most of the time, this is the least electronegative atom…in advanced chemistry we use a thing called formal charge to determine the central atom. But that’s another story!) Therefore, N is central on this one If there is a choice, the central atom is atom of lowest affinity for electrons. (Most of the time, this is the least electronegative atom…in advanced chemistry we use a thing called formal charge to determine the central atom. But that’s another story!) Therefore, N is central on this one Then, Go to the CAR! 1. Count valence e- 2. Arrange octets 3. Recount electrons

12. 12 Count Valence Electrons Add up the number of valence electrons that can be used. NH 3 Add up the number of valence electrons that can be used. NH 3 N = 5 N = 5 H = 1 times 3 H = 1 times 3 Total = 5 + (3 x 1) Total = 5 + (3 x 1) = 8 electrons / 4 pairs

13. 13 Arrange Octets Remember Hydrogen only needs 2! H H H N

14. 14 Recount Electrons 2, 4, 6, 8 electrons total. It works! Check the number of electrons in your drawing with the number of electrons from step 2. If you have more electrons in the drawing than in step 2, you must make double or triple bonds. H H H N

15. 15 In Other Words, Go To The CAR Count valence electrons Arrange octets Recount electrons to make sure they match the valence electrons If you go to the CAR, you will be in the driver’s seat!

16. 16 Carbon Dioxide, CO 2 1. Count valence e- 2. Arrange electrons 3. Recount electrons Place lone pairs on outer atoms. Place lone pairs on outer atoms. This leaves 12 electrons (6 pair). Check to see that all atoms have 8 electrons around it except for H, which can have 2. C 4 e- O 6 e- X 2 O’s = 12 e- Total: 16 valence electrons

17. 17 Carbon Dioxide, CO 2 6. There are too many electrons in our drawing. We must form DOUBLE BONDS between C and O. Instead of sharing only 1 pair, a double bond shares 2 pairs. So one pair is taken away from each atom and replaced with another bond. C 4 e- O 6 e- X 2 O’s = 12 e- Total: 16 valence electrons How many are in the drawing?

18. 18 Double and even triple bonds are commonly observed for C, N, P, O, and S H 2 CO SO 3 C2F4C2F4C2F4C2F4

19. 19 Violations of the Octet Rule (Honors only) Usually occurs with B and elements of higher periods. Common exceptions are: Be, B, P, S, and Xe. BF 3 SF 4 B: 6 ( remember this) Be: 4 P: 8 OR 10 S: 8, 10, OR 12 Xe: 8, 10, OR 12

20. 20 Lewis Dot Structures HOT or not?

21. 21 Definitely HOT!

22. 22 MOLECULAR GEOMETRY

23. 23 Now You Try One! Draw Sulfur Dioxide, SO 2 (hint: go to the CAR!)

24. 24 SO 2 1.Count valence electrons S = 6 O = 6(2) 6 + 12 = 18 electrons 2. Arrange octets O S O → 3. Recount electrons 18 = 18

25. 25 VSEPR VSEPR V alence S hell E lectron P air R epulsion theory.V alence S hell E lectron P air R epulsion theory. Most important factor in determining geometry is relative repulsion between electron pairs.Most important factor in determining geometry is relative repulsion between electron pairs. Molecule adopts the shape that minimizes the electron pair repulsions. MOLECULAR GEOMETRY

26. 26 Some Common Geometries Linear Trigonal Planar Tetrahedral

27. 27 VSEPR charts Use the Lewis structure to determine the geometry of the moleculeUse the Lewis structure to determine the geometry of the molecule Electron arrangement establishes the bond anglesElectron arrangement establishes the bond angles Molecule takes the shape of that portion of the electron arrangementMolecule takes the shape of that portion of the electron arrangement

28. 28

29. 29 Molecular Geometry Recall that atoms share electrons with other atoms, they do so in pairs. A pair of electrons (a bond) in overlapping orbitals holds two atoms together. The bond between two nitrogen atoms creates a diatomic nitrogen molecule. (Linear geometry)

30. 30 Structure Determination by VSEPR Water, H 2 O The electron pair geometry is TETRAHEDRAL The molecular geometry is BENT. 2 bond pairs 2 lone pairs

31. 31 Determine the Shape by Using VSEPR Ammonia, NH 3 The electron pair geometry is tetrahedral. The MOLECULAR GEOMETRY — the positions of the atoms — is TRIGONAL PYRAMID.

32. 32 Boron Trichloride Because boron is stable with 6 valence electrons and the electrons repel each other it makes a trigonal planar geometry.

33. 33 Methane CH 4 in 3 dimensions methane is a tetrahedral geometry

34. 34 Diatomic Elements These elements do not exist as a single atom; they always appear as pairs H 2 –Hydrogen - H 2 N 2 –Nitrogen - N 2 –Oxygen –Fluorine –Chlorine –Bromine –Iodine –Are these molecules polar, nonpolar or ionic? Why?

35. 35 Bond Polarity HCl is POLAR because it has a positive end and a negative end. (difference in electro-negativity of.3 or more) Cl has a greater share in bonding electrons than does H. Cl has slight negative charge (-  ) and H has slight positive charge (+  )

36. 36 Bond Polarity and Molecular Polarity Some molecules have polar bonds but are part of nonpolar molecules. And there are some nonpolar molecules with polar bonds. Molecular polarity is important because it predicts how molecules will behave, since the charges influence how the molecules will interact with each other.

37. 37 Molecular Polarity Bonds are polar when one atom is positive and the other negative. Molecules with many atoms have polarity, with one end positive, the other negatively charged. You can predict the polarity of the molecule by looking at the ends of the molecule to see if it has a positive end and a negative end. Lone pairs of electrons are negative while hydrogen atoms or other low electronegativity atoms tend to be slightly positive. Polar = positive end and negative end Nonpolar = same charge at both ends

38. 38 Molecular Polarity How can you determine if a molecule is polar? Look for lone pairs of electrons at one end (-) And hydrogens (+) at the other end.

39. 39 Molecular Polarity In water, the electron dot formula shows the 6 oxygen valence electrons and the 2 hydrogen electrons bonding to make a bent molecule. In three dimensions, this molecule has length, width and depth. The electrons repel each other (valence shell electron pair repulsion, VSEPR theory) into a bent shape. Because one end is positive and the other end is negative, it is polar.

40. 40 Molecular Polarity of Methane CH 4 The Lewis dot structure shows the carbon atom surrounded by eight electrons, and baby hydrogen with 2 electrons each. The 3-D geometry looks very different with the electron pairs repelling into a tetrahedron. Because there are hydrogens at each end, it is nonpolar. There is no negative- positive dipole.

41. 41 This is why oil and water will not mix! Oil is nonpolar, and water is polar.This is why oil and water will not mix! Oil is nonpolar, and water is polar. The two will repel each other, and so you can not dissolve one in the otherThe two will repel each other, and so you can not dissolve one in the other Molecular Polarity

42. 42 Molecular Polarity Examples “Like Dissolves Like”“Like Dissolves Like” –Polar dissolves Polar –Nonpolar dissolves Nonpolar

43. 43 Hot or Not?

44. 44 Totally Sizzling!

45. 45 The End