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8.7 Covalent Bonding

8.7 Covalent Bonding. Molecules arise from localized attractive forces between atoms, which we call covalent bonds Atoms in the molecule are connected strongly, but molecules are not strongly attracted to each other.

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8.7 Covalent Bonding

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  1. 8.7 Covalent Bonding • Molecules arise from localized attractive forces between atoms, which we call covalent bonds • Atoms in the molecule are connected strongly, but molecules are not strongly attracted to each other. • Molecular compounds are usually gases or liquids unless the molecules are very large

  2. Single Covalent Bonds • Sharing of 1 pair of electrons • Each atom has one half-filled valence orbital that overlap one another • H. + .H  H:H • Single bond represented asH:H or H–HCalled a Lewis formula or electron-dot formula

  3. 8.10 Lewis Symbols • Lewis Symbols:The number of valence electrons available for bonding are indicated by unpaired dots.

  4. Lewis Symbols • We generally place the electrons on four sides of a square around the element symbol.

  5. Single bonds between like atoms • Halogens

  6. Single bonds between unlike atoms • HF Figure 7.10

  7. Some atoms can form bonds with more than one atom • CCl4Figure 7.10How many valence electrons are supplied by each atom?

  8. Lewis Model of Covalent Bonding (Localized electron model) • A covalent bond is formed by the sharing of at least two valence electrons between two atoms. • At normal temperatures and pressures, for non-metals, the maximum number of valence electrons per atom is eight (with the exception of two for H and He). • Each shared electron is counted twice (one per atom). • A stable molecule is formed when sharing of valence electrons lowers the chemical energy.

  9. Procedure for Drawing Lewis Structures • Arrange the atoms (symbols) around the central atom. • Count up total valence electrons • Draw a line to represent a single bond (2 shared electrons) between the central atom and each surrounding atom. This is a 2-electron bond between the cores. • Place remaining electrons on the outside atoms, then the central atom • Shift electrons, as necessary, to make multiple bonds and satisfy the octet rule

  10. Whiteboard Work #1: • Draw the Lewis Structures for the following molecules: • BF3 • NH3 • NO2 • CO2

  11. 8.11 Exceptions to the Octet “Rule” • Incomplete Octets • Usually only seen for Be, B, Al • Leads to stable, yet reactive molecules! • Coordinate Covalent Bonding • Molecules with too few electron pairs can bond with molecules with unshared electron pairs to form a new shared-electron-pair bond • BF3 + NH3 F3BNH3

  12. Exceptions to the Octet Rule • Why is BH4- more stable than BH3? • Why is BF4- more stable than BF3? • Why does aluminum chloride exist in the gaseous state as Cl2AlCl2AlCl2 (that is, Al2Cl6) instead of AlCl3?

  13. 8.11 Exceptions to the Octet Rule • Odd-Electron Molecules • Also stable, yet reactive! • Examine the Lewis formula for NO2 • Why does NO2 combine with itself to form N2O4?

  14. CO2

  15. Multiple Bonds • Can share more than one pair of electrons to form double or triple bonds 2 electron pairs 3 electron pairs

  16. Whiteboard Work #2: Draw the Lewis structure for SO2

  17. 8.12 Resonance Structures • Lewis formulas don’t always accurately represent bonds. Sometimes it takes two formulas to adequately represent the bonds.

  18. Resonance Structures In the sulfur dioxide example, we could move the electron pair from either oxygen to form the double bond. The two structures we drew were equivalent. These are examples of equivalent resonance structures. The Lewis-dot representation suggests that the molecule exists as either one structure or the other, and that it sort of “flips” back and forth. This is NOT the case.

  19. Rather the electron pair of the double bond is sort of “smeared out” over both bond positions. This “smearing out” is referred to as delocalization, and can be represented using a dashed line. Each S–O bond is neither a single or a double bond. The three electron pairs distribute themselves among the 2 bonding positions and is basically a 3/2 = 1.5 bond.

  20. How many different valid Lewis formulas can you write for the following molecules or ions? How do they differ? SO2 SO3 H2SO4 CO32- NO3- HNO3 NCS- The different resonance forms represent delocalized bonding. Resonance

  21. SO2 SO3 H2SO4 CO32- NO3- HNO3 NCS- Write Lewis formulas for the following molecules or ions:

  22. Formal Charge • Can be used to decide between alternate Lewis structures • There is considerable controversy as to whether the concept of formal charge dictating electron distribution is in fact correct. • You can read more about it in your text. • Is helpful for evaluating Lewis structures

  23. Whiteboard Work #3 Draw the Lewis structure for I31– Draw the Lewis structure for SbF5

  24. Exception to the Octet Rule • Some bonding situations result in the central atom having > 4 electron pairs ( > 8 electrons) around it: • S can form SH2 with S having 8 valence electrons • S can form SCl4 with S having 10 valence electrons • S can form SCl6 with S having 12 valence electrons • How is this possible??? (Think QM…)

  25. Expanded Octets Non-Metal atoms in Period 3 or higher can have more than 8 electrons around them (i.e. in their valence shell or top shelf) Why is this? Recall period 3 atoms (and period 4, 5, 6… atoms) have d-orbitals that can hold those extra electrons!

  26. Exceptions to the Octet Rule • Expanded Valence Shells • What do you do if there are too many electrons to be accommodated by octets? • Write Lewis formulas for the following:SF4 SF6 IF4+ XeF4 XeF2 PF5 BrF3 BrF5 • When do we find expanded octets? What • is the origin of the octet rule?

  27. Modern View of Lewis Bonding Model • A stable molecule is formed when sharing of electrons lowers the chemical energy. • For non-metals, the number of valence electrons equals the column number. The maximum number of valence electrons per core is 8 (2 for H and He). • Usually a covalent bond is formed by sharing 2 valence electrons between 2 cores. • At low temperatures, a non-reactive substance usually has the maximum number of electrons per core (octet rule). There are some exceptions.

  28. Website: http://www.kentchemistry.com/links/bonding/lewisdotstruct.htm

  29. NH3 NH4+ CCl2F2 SOCl2 SO2 CO2 CO32- SO32- H2SO4 HCN CN- NCS- Write Lewis formulas for the following molecules or ions:

  30. Lewis Structures

  31. Write Lewis formulas for the following molecules or ions: PCl3 CO H2O2 (HOOH) SO42- Ch. 8 Group Quiz Put your names and class date in the heading!!!

  32. Origins of the “Octet Rule” • Octet rule: we know that s2p6 is a noble gas configuration. We assume that an atom is stable when surrounded by 8 electrons (4 electron pairs). (there are many exceptions)

  33. Octet Rule • What ion or compound is formed from the following to approximate a noble gas electronic configuration? What is the configuration? • Na1s22s22p63s1 • Na+1s22s22p6 • H 1s1 • H+ or H-1s0 or 1s2

  34. Octet Rule • Cl 1s22s22p63s23p5 • Cl-1s22s22p63s23p6 • O 1s22s22p4 • O2-1s22s22p6 • H + O 1s1 + 1s22s22p4 • H2O 1s22s22p6 for O, 1s2 for H • Na + O 1s22s22p63s1 + 1s22s22p4 • Na2O 1s22s22p6 for Na and O

  35. Octet Rule • C + H 1s22s22p2 + 1s1 • CH4 1s22s22p6 for C, 1s2 for H • C + Cl 1s22s22p2 + 1s22s22p63s23p5 • CCl4 1s22s22p6 for C, 1s22s22p63s23p6 for Cl • C + O 1s22s22p2 + 1s22s22p4 • CO2 1s22s22p6 for C and O

  36. 1. NO 1. NO2 2. N2O4 2. BH3 3. BH4- 3. AlCl3 4. SF4 4. SF6 5. IF4+ 5. XeF4 6. XeF2 6. PF5 BrF5 Write Lewis formulas for the following molecules or ions:

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