1. Chapters 8 & 9 Review By Robert Liu and Julie Baldassano

2. Chapter 8: Bonding-General Concepts Ionic bonds: bonds between two ions, usually with a very large electronegativity difference. Example: CaCl2 Covalent bonds: bonds in which electrons are shared between the two nuclei. Example: H2 Polar Covalent bonds: bonds in which electrons are shared, but unequally. Example: H2O

3. Chapter 8: Electronegativity The ability of an atom in a molecule to attract shared electrons to itself. Generally increases across a period and decreases going down a group

4. Chapter 8: Bond Polarity Molecules with a center of negative charge and a center of positive charge are known as dipolar. Any molecule with a polar bond will have dipole moments, making the entire molecule polar. However, symmetrical arrangements will cancel out dipole moments.

5. Chapter 8: Ions Ionic compounds are formed from a nonmetal ion and a metallic ion. (MgCl2, Cacl2) Cations are always smaller in radius than their parent atoms. Anions, always bigger. With isoelectronic ions(those with the same amount of valence electrons), ionic radius decreases as nuclear charge increases.

6. Chapter 8: Lattice Energy As charge increases, the lattice energy increases. Mg2+>Na+. As distance between the two nuclei in an ion decrease, lattice energy increases.

7. Chapter 8: Ionic Character Electronegativity and % ionic character have a positive relationship.

8. Chapter 8: Bond Energy Simple Arithmetic. Take H2 +F2=2HF for example. Bonds broken-bonds formed=(triangle)H

9. Chapter 8: The LEBM 1. Description of valence electron arrangement using lewis dot structures. 2. Predicting geometry using VSEPR 3. Description of the type of atomic orbitals used to share electrons

10. Chapter 8: Lewis Structures Use valence electrons only. Hydrogen obeys duet rule, everything else typically obeys octet rule. One dot= one electron, one line=a pair of electrons.

11. Chapter 8: Exceptions Boron tends to form compounds in which it is electron deficient. Third period elements can exceed the octet rule using their open d orbitals (citation needed) Odd numbers of electrons (NO), oxygen has 6, nitrogen has 5. 6+5=11 o_O? Results in 5 around nitrogen, 8 around oxygen.

12. Chapter 8: Resonance Occurs when more than one lewis structure can be drawn for a molecule.

13. Chapter 8: VSEPR Minimize repulsions by spacing atoms as far away as possible from each other. Bond angles: Linear=180 Trigonal Planar=120 Tetrahedral=109.5 Trigonal Pyramidal=<120 ~107 Bent=<120 ~104.5

14. Chapter 8 - Quiz 1) The geometry of the SO 3 molecule is best described as (A) trigonal planar (B) trigonal pyramidal (C) square pyramidal (D) bent (E) tetrahedral 2) For which of the following molecules are resonance structures necessary to describe the bonding satisfactorily? (A) H 2 S (B) SO 2 (C) CO 2 (D) OF 2 (E) PF 3 3) The electron-dot structure (Lewis structure) for which of the following molecules would have two unshared pairs of electrons on the central atom? (A) H 2 S (B) NH 3 (C) CH 4 (D) HCN (E) CO 2 1) A 2) B 3) A

15. Chapter 9 - Hybridization Hybridization: the mixing of native atomic orbitals to form special orbitals for bonding Molecules using localized electron model: 1) Draw Lewis structure 2) Place electron pairs using VSEPR model 3) Match effective electron pairs to hybrid orbitals

16. Chapter 9 - sp sp: linear shape 2 effective pairs 180 degree angles

17. Chapter 9 - sp 2 sp 2 : trigonal planar shape 3 effective pairs 120 degree angles

18. Chapter 9 - sp 3 sp 3 : tetrahedral shape 4 effective pairs 109.5 degree angles

19. Chapter 9 - dsp 3 dsp 3 : trigonal bipyramidal shape 5 effective pairs 90 and 120 degree angles

20. Chapter 9 - d 2 sp 3 d 2 sp 3 : octahedral shape 6 effective pairs 90 degree angles

21. Chapter 9 - σ and π bonds σ: electrons shared in area on a line between the atoms, localized bonding π: a shared electron pair occupies the space above and below the line between the atoms, delocalized bonding

22. Chapter 9 - MO Model Molecular orbital model: -Pros: no resonance, effective for molecules with unpaired electrons, bond energy information Antibonding (MO2) are higher energy than bonding (MO1) orbitals Bond order = # of bonding electrons - # of antibonding electrons 2

23. Chapter 9 - MO Model - Start at lowest energy level - Each orbital can hold two electrons (each with different spins) - Use number of valence electrons to fill - Degenerate = B2, C2, N2 so pi2p is lower energy than sigma2p, flipping the lines

24. Chapter 9 - Magnetism & Trend Paramagnetic: attracted to magnetic field - unpaired electrons Diamagnetism: repelled from magnetic field -paired electrons - High bond order = high bond energy = short bond length

25. Chapter 9 Quiz 5) Using molecular orbital theory, determine the magnetism of O 2 and O 2 −. a- O 2 is paramagnetic; O 2 − is diamagnetic. b- O 2 is diamagnetic; O 2 − is paramagnetic. c- Both O 2 and O 2 − are diamagnetic. d- Both O 2 and O 2 − are paramagnetic. 2) HCN sp sp 2 sp 3 sp 3 d sp 3 d 2 4) SF 4 sp sp 2 sp 3 sp 3 d sp 3 d 2 1) SH 2 sp sp 2 sp 3 sp 3 d sp 3 d 2 3) NH 2 1- sp sp 2 sp 3 sp 3 d sp 3 d 2

26. Chapter 9 Quiz Answers 1) SH 2 is sp 3 2) HCN is sp 3) NH 2 1- is sp 3 4) SF 4 is sp 3 d 5) d- Both O 2 and O 2 − are paramagnetic.