Unit 6: Bonding, Shapes & Theory

1. Unit 6: Bonding, Shapes & Theory Chapters 9, 10 & 11

2. Chemist ?s Review • Where do elements get their properties? • How can I predict whether two elements will react? • How do the quantum numbers-n, l, ml and ms differ? n = principal quantum number l = angular momentum (shape) ml = magnetic angular (orientation) ms = magnetic spin

3. 3Types of Bonding • ___________- Electrons donated or gained for element to reach a noble gas electron configuration • __________- Electrons shared to reach noble gas electron configuration (octet rule) • ___________-Electrons free floating in a sea of electrons; tightly packed Use the Periodic Table to predict elements that will create ionic bonds, covalent bonds or metallic bonds.

4. Ionic Bonds Energy • Examine energy exchange: • Li (g)  Li+ (g) + e-1 IE1 = 520 kJ (Ionization Energy) • F (g) + e-1  F- (g) EA = -328 kJ (Electron Affinity) • Combined: • Li (g) + F (g) Li+ (g) + F- (g) IE1 + EA = 192 kJ • Is this endothermic or exothermic? However: • Li+ (g) + F- (g)  LiF (g) ΔH = -755 kJ • Li+ (g) + F- (g)  LiF (s) ΔH = -1050 kJ • Lattice Energy- Enthalpy change that occurs when 1 mole of ionic solid separates into gaseous ions. • High lattice energy influences melting point, hardness, solubility and reactivity.

5. Effect of Nuclear Charge (Z) • Z= Atomic number of an element • Nuclear protons pull on electrons with electrostatic force (opposites attract ) • When there are numerous electrons, they “shield” or “screen” out some of the positive attraction of the nuclear protons. Can you think of a way to illustrate this? Zeff = Effective nuclear charge is the nuclear charge an electron actually experiences as a result of shielding effects due to the presence of other electrons.

6. Shielding & Ionization Energies Shielding affect atomic size as well. As the Zeff increases the electrons fill a greater nuclear proton attraction and the electrons are pulled in tight and the radius decreases. The opposite effect is the n, principal quantum number. As n, increases the electrons are further away and the radius increases. http://www.chemicalelements.com/elements/li.html Going down a Group- n wins and the atom increases in size Going across a Period- Zeff wins and the atom decreases in size

7. Ultimately…The Li Story • Li (s) takes energy to convert to Li (g) • F2 must be separated into F (g) for the reaction • Only ½ of the F2 BE (binding energy) value 159 kJ is used • Pg 346-347 outlines the steps. • The magnitude of the lattice energy overwhelms the multistep process. Even though we saw endothermic reactions previously, the stability dominates the process. Ionic Solids exist only because the lattice energy exceeds the energetically unfavorable electron transfer.

8. Periodic Trends in Lattice Energy • Ionic size increases increasing the charge distance causing the electrostatic energy to decrease. • Ionic charge (like Li+ compared to Mg+2) increases the stability of the ionic bond therefore giving it a bopping lattice energy. In spite of large IE1 coupled with larger IE2, the charge sizes are multiplied yielding a huge Electrostatic energy. Electrostatic energy ≈ Charge A × Charge B ≈ ΔH° lattice cation radius + anion radius

9. Single/Double/Triple Bonds 1- Bond Energy (BE) increases with greater # of bonds 2- Bond Length decreases with greater # of bonds

10. Periodic Trends in Bond Strength & Length • Bond breakage is endothermic- take a part an atomic model; bond energy is positive (ΔH = “+”) • Bond formation is exothermic- put together a model; bond formation is negative (ΔH = “-”) • Just as atomic radii increase going down a Group so does the ionic radii which is directly proportional to the bond length • Atomic radii decreases for cations (cations are smaller than their atoms) and anions (which are larger than their atom) going across a Period so does the ionic radii which means the bond length decreases as well.

11. “Me” vs. “We” complex Water molecule vs. water http://www.google.com/imgres?imgurl=http://www.mbi-berlin.de/en/research/projects/2-04/highlights/water_librational_mode.gif&imgrefurl=http://www.mbi-berlin.de/en/research/projects/2-04/highlights/MolStructDynamics-2005.html&usg=__p_A4RctIYzzcCSUsHQHnK96-WrM=&h=836&w=1032&sz=3196&hl=en&start=0&sig2=KQqJOmqNhobSvS9BhvRldg&zoom=1&tbnid=QYMrej6ojdDYgM:&tbnh=130&tbnw=160&ei=sl0dTZ3jOMG78gbD8-SUAw&prev=/images%3Fq%3Dwater%2Bmolecule%26um%3D1%26hl%3Den%26client%3Dfirefox-a%26rls%3Dorg.mozilla:en-US:official%26channel%3Ds%26biw%3D1920%26bih%3D840%26tbs%3Disch:1&um=1&itbs=1&iact=hc&vpx=1503&vpy=531&dur=4983&hovh=202&hovw=249&tx=127&ty=152&oei=sl0dTZ3jOMG78gbD8-SUAw&esq=1&page=1&ndsp=42&ved=1t:429,r:40,s:0

12. Forces • Strong covalent Forces- Intramolecular • Weak intermolecular Forces or van der Waals Forces- Table 12.2 (p 451) shows these include Ion-dipole, H bonds, dipole-dipole, ion-induced dipole, etc.

13. Bond Energy • Δ°Hrxn= ΣΔH°reactantbondsbroken + ΣΔH°product bonds formed • Exothermic reactions: products have more energy so ΔH = “-” ; Since the products absolute value is greater then the Δ°Hrxn= “-” • Endothermic reactions: products have a smaller absolute value than that for reactant bonds brokesn so the sum Δ°Hrxn = “+” • (Bonds broken requires input of energy so it is positive.)

14. Electronegativity • The relative ability of a bonded atom to attract shared electrons. Make Periodic Table predictions. • This value is for atoms already in a “relationship” (bonded). Electron affinity refers to a single atom in the gas phase gaining an electron to form a gaseous anion Non-polar covalent- bond of evenly shared electrons Polar covalent- bond of UNEVENLY shared electrons

15. Summary of Electronegativity Differences • http://www.google.com/imgres?imgurl=http://www.homework-help-secrets.com/images/bond-types.jpg&imgrefurl=http://www.homework-help-secrets.com/electronegativity.html&usg=__0pmiYyUZTO4_0oFs5L32CrRirvc=&h=649&w=770&sz=39&hl=en&start=84&sig2=6eKguG435sHoG3NzhK4ZrQ&zoom=1&tbnid=FXq_Y_H1a9IxPM:&tbnh=164&tbnw=195&ei=BgYhTe_pKc28nAe_w6XODg&prev=/images%3Fq%3Dnonpolar%2Bbonding%26hl%3Den%26client%3Dfirefox-a%26sa%3DG%26channel%3Ds%26rls%3Dorg.mozilla:en-US:official%26biw%3D1920%26bih%3D840%26gbv%3D2%26tbs%3Disch:10%2C1399&itbs=1&iact=hc&vpx=1636&vpy=170&dur=477&hovh=206&hovw=245&tx=195&ty=157&oei=YgUhTf3cPIyr8AbGjM2ZAw&esq=9&page=3&ndsp=33&ved=1t:429,r:8,s:84&biw=1920&bih=840

16. Working in Groups • Creatively illustrate the following topics with a dramatic production; monologue/dialogue/sketch. • Actors must have names/characteristics consistent with their nature and reactivity; give real examples. • 1) Nuclear shielding affect-explain and how does it influence reactions • 2)Bond Energy- single/double/triple bonds-bond length & strength • 3) Periodic Trends- atomic radius, ionic radius, bond length & bond strength • 4) Intramolecular forces vs. intermolecular forces • 5) Electronegativity in determining ionic, polar covalent or non-polar covalent nature • 6) Electronegativity vs. electron affinity

17. Lewis Dot structures • Single atom/ion- represent each valance electron with a dot. Ions behave to react a noble gas configuration. • Molecular formula- • 1) atom with the lowest Group # or least electronegativity goes in center • 2) add atoms surrounding • 3) sum valence electrons for all atoms • 4) draw single bonds connecting atoms (reducing valance electron sum by two each time) • 5) distribute remaining electrons (8-but only 2 for H)

18. Cl Cl Cl • CCl4 - • C • Add single bonds • Add lone “unpaired” electrons Cl How would you draw the Lewis dot structure for Oxygen gas? Nitrogen gas?

19. Resonance hybrids • Delocalized Electron-Pair Bonding- When the electrons location is split between double/single bonds. • Resonance structures illustrate the multiple locations of the bonding and lone electron pairs. • http://www.google.com/imgres?imgurl=http://www.transtutors.com/Uploadfile/CMS_Images/5346_Canonical%2520structures%2520%26%2520Resonance%2520hybrid.JPG&imgrefurl=http://www.transtutors.com/chemistry-homework-help/S-and-P-block-elements/trioxides-of-oxygen-family.aspx&usg=__Oo9NvX15un_puZSfAMQOf2zAz_A=&h=123&w=544&sz=10&hl=en&start=0&sig2=-QLh2XdK9pCgdwChM8jW2w&zoom=1&tbnid=IFpxBmqXHy0l5M:&tbnh=62&tbnw=275&ei=p0ghTcmvPMG78gbF8-SUAw&prev=/images%3Fq%3Dresonance%2Bhybrid%2Bstructures%26hl%3Den%26client%3Dfirefox-a%26channel%3Ds%26rls%3Dorg.mozilla:en-US:official%26biw%3D1920%26bih%3D840%26gbv%3D2%26tbs%3Disch:1&itbs=1&iact=hc&vpx=375&vpy=393&dur=4960&hovh=98&hovw=435&tx=214&ty=63&oei=p0ghTcmvPMG78gbF8-SUAw&esq=1&page=1&ndsp=37&ved=1t:429,r:15,s:0

20. Formal Charge per Atom • Identify each atom; determine it’s valance electrons • Subtract off unshared electrons • Subtract off half of their shared electrons—after all half belong to the other atom • Sum: That is the total charge • Examine: N03- N 5 e-1 O6 e-1 • N has 0 unshared O1 has 3 unshared or O2 has 2 unshared • N has 4 single bonds: ½ x 4 shared (8) • Each O1 has 1 single bond: ½ x 1 shared (2) or 1 double: 1/2 x 2 shared(4) • N = 5 – 0 – 4 = 1 • O1= 6 – 6 – 1 = -1 O2 = 6 – 4 – 2 = 0

21. Exceptions to the Octet Rule • 1) Less than an Octet of Valance Electrons: • BF3 Drawing: • 2) More than an Octet of Valance Electrons: PCl5 Drawing: • Also-SF4, SF6, AsF6-, ICl4- • 3) Odd Number of Electrons • NO Drawing (double bonded O)

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