Chapter 8

1. Chapter 8 Bonding: General Concepts

2. Chemical Bonds • Bond Energy - energy required to break a bond • Types: ionic and covalent

3. Ionic Bonding • Electrostatic attractions between tightly packed, oppositely charged ions • Between metals (cations) and nonmetals (anions)…can be polyatomic • High melting/boiling points • Ex: NaCl

4. Coulomb’s Law E = (2.31 X 10-19 J * nm)(Q1Q2/r) • Calculates the energy between two ions • E -> Energy in Joules • R -> distance between ion centers in nm • Q1 and Q2 are ions’ charges • When E is negative, the pair of ions has a lower energy than the ions separated (ions are attracted) • When E is positive, the repulsive energy is greater (occurs with two like-charged ions) • In nature, a system wants the lowest possible energy

5. Covalent Bonding • Molecules where electrons are SHARED by nuclei • Equal sharing of electrons occurs between the two atoms • Nonequal sharing = polar covalent

6. Identical Atoms • Diatomic atoms… • Systems of energy favor LOWER energy • If a system can lower energy by forming bonds, it will happen

7. Bond Lengths • Distance at which the system has minimum energy

8. Polar Covalent Bond • Covalent bond where one atom pulls more on an electron than the other • Ex: H-F (fluorine slightly negative) • + and - are used to represent slight charges • Water is another example • Due to electronegativity

9. Electronegativity • (electron affinity)…ability to attract shared electrons • Higher electronegative atom will have the slight negative charge • Higher difference in electronegativities, more likely to be an ionic bond

10. Bond Polarity • Dipolar/dipole moment occurs when a molecule has a slight positive and slight negative charge. • Arrow points towards the slight negative charge with a plus sign on the other end of the arrow • Electrostatic potential diagram also used (red = electron-rich, blue = electron-poor)

11. Polar Bond without Dipole Moment • Polar bonds where slight charges cancel each other out • No lone electrons

12. Ion Electron Configuration and Sizes • Atoms of stable compounds (ionic and covalent), have noble gas electron configurations • Covalent: electrons are shared so that the valence electron configurations of both nonmetals attain noble gas e- configurations • Ionic: Nonmetal achieves e- configuration of the next noble gas atom and the metal’s valance electrons are emptied so both ions achieve noble gas e- configurations

13. States of Ionic Compounds • Usually “ionic compound” means a solid state where ions are close to one another/interacting, minimizing the - - and + + repulsions and maximizing the + - attractions

14. Trend in Ions’ Sizes • Positive ion = loss of e- (cation is smaller than original atom) • Negative ion = gain of e- (anion is larger than neutral atom) • Ion size increases going down a group • Horizontally is more complicated (metals vs. nonmetals)…..

15. Isoelectronic Ions • Ions that have the same electron configuration (b/c have the same number of electrons) • O2-, F-, Na+, Mg2+, Al3+ • Same number of electrons, but different numbers of protons • From O2- to Al3+ attraction to nucleus (increase in protons) increases so smaller ions

16. Bonds • Single Bonds: one pair of electrons shared • Double bond: two pairs of electrons shared • Triple bond: three pairs of electrons shared • More bonds = shorter bond length • More bonds = stronger bonds • Bond energies - amount of energy required to break bond

17. Breaking Vs. Forming Bonds • To break bonds, energy must be added (required, endothermic, energy is positive) • To form bonds, energy must be removed (given off, exothermic, energy is negative) • Change in enthalpy: DH = sum of the energies required to break old bonds (positive signs) plus sum of energies released in formation of new bonds (negative signs)

18. Bond Energies • Energy stored in bonds can be used to determine energy accompanying various chemical reactions. • Bond dissociation energies = energy required to break bonds (positive kJ/mol, endothermic)… reverse for forming bonds. Table 8.4 in book • DH = all bonds broken – all bonds formed • Must use a balanced equation!

19. EXAMPLE • Use data 8.4 in your book to determine DH for the following reaction. CH4(g) + 4F2(g) CF4(g) + 4HF(g) • Answer: -1932 kJ

20. Localized Electron Bonding Model…Highlights • Defined: A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bonded atoms • Localized to one of the atoms (LONE PAIR) • Localized to the space between atoms (BONDING PAIR)

21. LE Model Parts 1. Description of the valence electron arrangement in the molecule using Lewis structures 2. Prediction of the geometry of the molecule using the valence shell electron-pair repulsion (VSEPR) model 3. Description of the type of atomic orbitals used by the atoms to share electrons or hold lone pairs

22. LE Model Part 1: Lewis Structures • All atoms want to have noble gas electron configurations • Only valence electrons are included • Determine total valence electrons • Determine layout • Determine bonds • Place remaining valence electrons

23. Lewis Dot Structure Examples • Ammonia: NH3 • Water: H2O • Acetylene: C2H2 • Carbon Tetrachloride: CCl4 • Dihydrogen Selenide: H2Se

24. Exceptions to Octet Rule • If valence electron total is odd, the octet rule doesn’t work • Some atoms do not require all 8 valence electrons (or have more than 8) • These molecules can exist in stable form • Boron: forms compounds where boron has less than 8 electrons (ex: BF3) • More than 8 only happens with elements in period 3 and beyond ex: SF6 • See pg. 371 purple box for rules if needed

25. Exception Examples • Draw the Lewis dot structure for the following: • ClF3 • XeO3 • RnCl2 • BeCl2 • ICl4-

26. Resonance • Molecules with more than one possible electron dot structure • Do not switch back and forth • Molecules exist as a mixture (hybrid) of the resonance forms • Use double headed arrow to signify • Example:

27. Resonance Example • Draw the Lewis dot (Localized Electron Model) structure(s) of CO32-

28. Formal Charges • Can’t use oxidation numbers because electrons are not shared evenly between atoms (electronegativities) • Atoms can be assigned formal charges using the following: Formal Charge = (# valence e- on atom) – (# valence e- assigned to the atom in the molecule) # valence e- assigned = (# lone pair e-) + (½ # shared e-) • Atoms want to have formal charges close to zero • Negative formal charges should be on the most electronegative atom • ***ESTIMATES of charges (not exact charges)

29. Example • Assign formal charges to each atom in CO2 • Which is more likely? • Answer: two double bonds b/c all formal charges are zero • Draw all resonance structures and show the most stable one for SCN- • Answer: two double bonds b/c N more electronegative than S

30. LE Model Part 2: VSEPR Model • Molecular structure shows 3-D arrangement of molecules • Based on minimizing electron-pair repulsions (bonding and nonbonding electrons will be placed as far apart as possible)

31. No Unshared Pairs • Linear (180°) • BeCl2 • Trigonal Planar (120°) • BCl3

32. No Unshared Pairs • Tetrahedral (109.5°) • CClF3 • Square Planar (90°)

33. With Unshared Pairs • Bent (104.5°) • H2O • Pyramidal (107°) • NH3

34. More to Consider • Tigonal bipyramidal • See-saw • T-shaped • Octahedral  • Deal with atoms that do not obey the octet rule

35. Examples • H2S • CO2 • PCl3 • OCS • H2CO • CH4 • N2H4