Chemical Bonding: Electronegativity, Types of Bonds, and Molecular Geometry

1. Chapter 9 and 10 Chemical Bonding

2. TOPICS in Chapter 9 • Electronegativity • Types of Bonds • Bond Polarity • Lewis Dot Structure

3. TOPICS in Chapter 10 • RESONANCE • FORMAL CHARGE • MOLECULAR GEOMETRY – drawing shapes of molecules

4. Why are things the way they are? • Why are some things soft? • Why are some things sturdy and hard? • What determines the properties of compounds?

5. Why do bonds form? • Chemical bonding is the universe’s answer to creating order amid disorder. It results from the tendency of a system to seek its lowest energy.

6. Chemical bonds • The bond is a human concept developed to provide a method for dividing up the energy evolved when a stable molecule is formed from its component atoms.

7. One Thing to Note! • Since bonding is a human invention, it is just an oversimplification of what is actually involved when atoms or ions combine to form a molecule.

8. Changes in Exam Dates • Nov. 10/12 LAB EXAM • Nov. 17/19 EXAM 2

9. Types of Chemical Bonds • Ionic Bonds - involves transfer of electrons • Covalent bonds - involves sharing of electrons • Polar Covalent bonds • Coordinate Covalent bonds – involves lone pairs

10. Bond Energy • energy required to break a bond • The higher the bond energy, the stronger the bond. It means that it will take more E to break the bond.

11. Bond Formation • A bond will form if the system can lower its total energy by doing so. • If more energy is used up in the process, no bond will be formed.

12. Bond Length • Is the distance between 2 ion centers • Is the distance at which the system has minimum energy • The system will act to minimize the sum of the repulsive energy and attractive energy.

13. Localized Electron ModelakaLewis Structures for Molecules and polyatomic ions • Predicts the bonding in covalently bonded molecules and ions • (indicates which atoms are bonded together)

14. Localized Electron Model • Sum up the valence e- (add or subtract to account for ion charge) • Pick central atom (usually one of a kind, seldom oxygen, never hydrogen) • Connect all atoms to central atom with one bond (which is 1 pair of electrons)

15. Localized Electron Model Rules • Arrange remaining electrons to satisfy: • - DUET for Hydrogen • -Octet for all other elements • (Note: B and Be are exceptions) • -If too few e-’s, use double or triple bonds as necessary • If too many e-’s, assign them to elements in the 3rd period or higher, starting with the central atom

16. Drawing the Localized Electron Model (aka Lewis Structure for Molecules) of Compounds • Count the # of valence electrons in the compound. • [NOTE: the # of valence electrons MUST always be EVEN.] 2. Determine the central atom • [NOTE: the central atom is usually the 1st Element, unless it is H. H only forms 1 bond and can never be the central atom.]

17. 3. Distribute the other elements around the central atom and connect them to the central atom with Single Bonds. 4. If the outer elements follow the Octet Rule, give them an octet. If they don’t, leave them be.

18. 5. To have the correct drawing, count the number of electrons in the drawing. Case I: If # of electrons = # of valence electrons AND central atom has an OCTET (if it forms an octet), then you are done. Case II: If # of electrons in drawing < # of valence electrons AND central atom CAN HAVE MORE THAN an OCTET (because it belongs to period 3 and higher), then add LONE PAIR(s) of electrons to the central atom. (You essentially already do this WHEN you follow MY METHOD.)

19. Case III: If # of electrons = # of valence electrons AND central atom DOES NOT HAVE an OCTET (but it DOES follow the OCTET Rule), then move one or more LONE PAIR(s) of electrons from the outer atom to the central atom. Moving the lone pair of electrons does not change the # of electrons in the compound. STOP moving lone pairs around as soon as central atom has an octet! Do not overdo things.

20. Localized Electron Model Rules • TWO Special Cases • When more than 1 valid and equivalent structure can be drawn for a molecule/ion, use resonance • When more than 1 valid but non-equivalent structures can be drawn for a molecule/ion, use formal charge

21. Formal Charge Method • Draw all valid but non-equivalent structures • Determine the formal charge on each atom in the structure • The structure with the smallest formal charges on each atom and with the negative formal charges on the most electronegative elements is the correct structure

22. Formal Charge Method • Formal charge is defined as the # of valence electrons minus the # of electrons directly on the atom in the molecule • # of electrons assigned to atom: • Lone pairs belong to the atom • ½ of shared electrons belong to atom

23. Valence Shell Electron Pair Repulsion ModelakaVSEPR • Expands the localized electron model by predicting the shape of the covalently bonded molecules and ions

24. VSEPR • Draw the localized electron model • Distribute the central atom’s effective electron pairs as far apart as possible. • Effective Electron pairs are: • Lone pair of electrons • Single bond electron pairs • Treat double or triple bonds as an effective single electron pair • Add atoms, keeping remaining lone e- pairs as far apart as possible

25. Molecular Shape • Shape of molecule is determined by position of atoms about the central atom, not by the location of electron pairs

26. VSEPR Model Key Points • Electron pairs are placed as far apart as possible to minimize electron repulsion • Molecule shape is determined by location of atoms while keeping remaining lone electron pairs as far apart as possible

27. VSEPR Model • Valence Shell Electron Pair Repulsion • useful in predicting geometries of compounds formed from non-metals

28. VSEPR Postulate • Main Postulate: The structure around an atom is determined by minimizing electron-pair repulsions. • In other words, the bonding and non-bonding pairs around an atom will be as far apart from each other as possible.

29. VSEPR Model • Lone pair electrons require more room than bonding pairs and tend to compress the angles between the bonding pairs.

30. Geometry vs # of lone pairs • Linear - 0 • Bent - 2 • Trigonal planar - 0 • Trigonal pyramidal - 1 • Tetrahedral - 0

31. Geometry vs # of electron pairs • Linear - 2 • Bent - 2 • Trigonal planar - 3 • Tetrahedral - 4 • Trigonal bipyramidal - 5 • Octahedral - 6

32. VSEPR dictates Geometry! • Examples: • BeCl2 - linear • CH4 - tetrahedral • H2O - bent

33. Types of Geometry • Linear - CO2 • Bent - H2S • Trigonal planar - BF3 • Trigonal pyramidal - NH3 • Tetrahedral - CCl4 • Octahedral

34. Dipole Moments • Polar bonds have dipole moments! • Example: If NaCl is place in an electric field, the molecule will have a preferential orientation. • Na- Cl d+ d-

35. Dipole Moment • Representation for dipole moment: • d+ d- • Polar bonds have dipole moments. • Polyatomic molecules can also exhibit dipole moments.

36. Always Remember! • Water is a polar molecule and thus has a dipole moment!

37. Also remember…... • The presence of a polar bond does not necessarily yield a polar molecule! • If polar bonds cancel out, then molecule is non polar.

38. When do polar bonds cancel out?

39. Polar Bonds Cancel out in…. • Linear molecules with 2 identical bonds • example: CO2, CS2 • Planar molecules with 3 identical bonds that are 120 degrees apart • example: BF3, AlCl3 • Tetrahedral molecules with 4 identical bonds that are 109.5 degrees apart • example: CCl4, CBr4

40. Hybridization • To determine the hybridization of the atom: • A. Include ALLlone pair of electrons and ALL atoms attached to element in question. • Therefore, consider double bonds and triple bonds as 1 bond since only 1 atom is attached to it.

41. Valence Bond Theory • Is the theory that deals with hybridization • 1. states that valence electrons reside in atomic orbitals that can be hybridized • 2. A chemical bond results in the orbital of 2 half-filled orbitals occupied by electrons of opposite spins • 3. The geometry of overlapping orbitals determines the shape of the molecule.