Na (g)  Na + (g) + e - [Ne] 3s 1 [Ne] Cl (g) + e -  Cl – (g)

2. Fomation of ionic compounds Na(g) Na+(g)+e- [Ne]3s1 [Ne] Cl(g)+e- Cl–(g) [Ne]3s23p5 [Ne]3s23p6 Sum: Na(g)+Cl(g)+e- Na+(g)+e-+Cl–(g) Net: Na(g)+Cl(g) Na+(g)+ Cl–(g)

3. Ionization energy Electron affinity Step 1 Step 2a Step 2b Step 1: ionization energy + electron affinity (+147 kJ/mol) Step 2a: formation of a single ion pair(-493 kJ/mol) Step 2b: formation of crystal lattice (-293 kJ/mol) from many ion pairs TOTAL energy change: -639 kJ/mol

4. Bond formation releases energy • when a covalent or non-covalent (i.e. ionic) bond is formed between atoms (or ions), energy is released • to break bonding interactions, energy is required (“absorbed”)

5. “Ball and stick” Cl H S Cl Cl “spacefilling”

6. Covalent Bonding = BDE Figure 9.11 p 342

7. The Electron Probability Distribution for the H2 Molecule

8. = BDE

9. CH4 Cl2 CH3Cl HCl Is this reaction energetically favorable? Is it exothermic or endothermic? You can answer this question by calculating bond energies for the reactants and products ... this is presented in more detail in Chem 122

10. . . . . . . . . . . . Lewis electron-dot symbols the electrons in the valence shell of an atom or ion are represented by dots placed around the symbol for the element. IA IIA IIIA IVA VIIIA Period ns1 ns2 ns2np1 ns2np2……. ns2np6 second Li Be B C ……... Ne third Na Mg Al Si ……... Ar . .. . . . . .. .. . . .. .. . .. .. . ..

11. Lewis diagrams • two electrons per “orbital”; four “orbitals” per atom (except H, He) • in a Lewis structure the four “orbitals” (represented by four sides of the atomic symbol) are considered equivalent • electrons 1-4 are added singly, 5-8 are added to form pairs • singly occupied “orbitals” indicate the potential to form bonds • a bond between atoms involves overlapping orbitals then filling the overlap with two electrons • non-bonding electrons are called “lone pairs”, they are not shared between atoms.

12. Outer shell “octet” • “noble gas core” = 8 electrons (an “octet”) in the valence shell. The exception is He which has only 2 e- in valence shell. • this is a very stable configuration of electrons • every atom wants an octet (except elements 1-5: H, He, Li, Be, B) Li+, Be2+, B3+all have [He] core

13. The octet rule • the number of electrons in lone pairs plus the number in covalent bonds should add up to 8 (except for H where the number of electrons should be 2) • lone pair electrons count as 2; bonding electrons count as 2 (even though they are shared!)

14. Lewis diagrams (continued) • electrons 1-4 are added singly, 5-8 are added to form pairs • singly occupied “orbitals” indicate potential to form covalent bonds. The atom with the greatest potential to form bonds is frequently a good choice as the “central atom” (e.g. C, N, S, P, O; NOT: H, F) • Total # of e- shown in Lewis diagram = total # of valence e- contributed by all atoms (adjusted for charge on molecular ion when applicable)

15. Drawing Lewis diagrams* • Select “central atom” based on bond-forming potential or electronegativity (in general, less electronegative atoms are “central atoms”) • Draw a potential skeleton for the molecule and distribute valence e- • Redistribute e- to satisfy octet rule (multiple bonds or coordinate covalent bonds may be required) • Make sure the total # of e- shown in the Lewis diagram matches # of e- predicted from valence e- of all atoms (and charge on ion) *see p 348-350

16. Draw Lewis structures for the following molecules [H3O] 1+ C2Cl2 COCl2

17. Electronegativity • electronegativity = the tendency of an atom to attract electrons in a bond (it describes the tendency of an electron to be found near the nucleus of an atom in a chemical bond). • NOT the same as “electron affinity” (electron affinity is measured in energy/mol)

18. “Electronegativity” Trends within Periodic Table (p346) arrow direction indicates increasing electronegativity most electronegative non-metals metals least electronegative

19. Electronegativity: measure of the tendency for an atom to attract bonding electrons Trends: increases left to right within period and from bottom to top within group.

21. covalent bonds revisited • in a covalent bond between atoms with same electronegativity, there is “equal” probability of finding electrons around either bonded nucleus. • in a polar covalent bond the more electronegative atom gains partial (-) charge, and the less electronegative atom gains partial (+) charge. This partial charge separation gives rise to “dipoles”.

23. “Resonance” • When e- density is delocalized over several atomic nuclei, equivalent Lewis structures can be drawn. These structures are called “resonance” structures and are meant to represent a close approximation of the actual (delocalized) electronic structure of the molecule. N2O4

24. Bond order and bond length • Bond order = number of bonds between atoms • Bond length decreases (bonds get shorter) as bond order increases.

25. “polar” vs. “non-polar” molecules • A molecule is considered “polar” if it has a net “dipole moment”. A molecule is “non-polar” if it has no net dipole moment (or if the dipole moment is negligibly small) • Characteristics of a polar molecule: • Polar covalent bond(s) • Molecular geometry that gives rise to a net dipole moment

26. Lewis structures: Formal Charge Formal charge (FC) = a hypothetical charge on an atom that is calculated as follows: FC = (# of Ve-) – (# of LPe- + ½# of Be-) Where: Ve- = valence electrons for the unbonded atom LPe- = lone pair electrons around the atom Be- = bonding electrons around the atom