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Chapter Six Representing Molecules

Chapter Six Representing Molecules. Section 6.1 The Octet Rule. The Octet Rule. Recall: atoms want noble gas configurations Octet Rule: atoms will gain, lose, or share electrons to achieve a noble gas configuration Typically 8 valence electrons

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Chapter Six Representing Molecules

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  1. Chapter SixRepresenting Molecules

  2. Section 6.1The Octet Rule

  3. The Octet Rule • Recall: atoms want noble gas configurations • Octet Rule: atoms will gain, lose, or share electrons to achieve a noble gas configuration • Typically 8 valence electrons • Atoms will bond with each other to achieve a full octet

  4. Lewis Structures • A pair of shared electrons can be represented by either with 2 dots or with a dash • Unshared electrons are called lone pairs F F F F

  5. Lewis Structures • Types of bonds: • Single bonds: bond containing only 2 electrons • Multiple bonds: bond containing more than 2 electrons • Double bond: 4 electrons (or 2 pairs of electrons) • Triple bond: 6 electrons (or 3 pairs of electrons)

  6. Bond Strength • In a particular pair of elements • Triple bonds are the shortest • Double bonds are in the middle • Single bonds are the longest Bond energy is the energy required to BREAK bonds between atoms

  7. Section 6.2Electronegativity and Polarity

  8. Electronegativity • Electronegativity is a periodic trend • Ability of an atom to attract electrons to itself when bonded to another atom • Quantified by the Pauling Scale

  9. Electronegativity

  10. Electronegativity

  11. Electronegativity

  12. Categories of bonds • Let’s consider three molecules: • H2, HF, NaF

  13. Ionic, Polar Covalent, Nonpolar • Take the difference in electronegativities of two atoms bonded together • If the difference is 0.5 or lower, the bond is nonpolar covalent • If the difference is between 0.5 and 2.0, the bond is polar covalent • If the difference is greater than 2.0, the bond is ionic

  14. Determine if the bond is ionic, polar covalent, or nonpolar covalent • The bond in ClF (chlorine and fluorine) • The bond in CsBr • The carbon-carbon double bond in C2H4 • In which of the following molecules are the bonds most polar: H2O, BCl3, PCl5

  15. Section 6.3Drawing Lewis Structures

  16. Drawing Lewis Dot Diagrams • 1) Determine the central atom and place terminal (“outside”) atoms around central atom • Central atom typically is the least electronegative element in compound, the element with only 1 atom, and/or the element written first in compound • 2) Count total # of v.e. • 3) Bond all terminal atoms to central using single bond • Each bond is 2 electrons; subtract from total # of v.e.

  17. Drawing Lewis Dot Diagrams • 4) Complete the octets of terminal atoms w/ remaining v.e. • 5) If any electrons left over, put on central atom • 6) Use multiple bonds to complete octet of any elements where necessary

  18. Examples of Lewis Dot Structures • CH4 • H2O • O2 • CO2 • CN-

  19. Group Quiz #1 • Draw the Lewis Dot Structures for the following: • CS2 • NF3 • ClO3-

  20. Section 6.4Lewis Structures & Formal Charge

  21. Formal Charge • Another way of keeping track of electrons in a molecule • Formal Charge = (# of v.e.) – (# of associated electrons) • Ex: Ozone (O3) • Now you try: NO3-

  22. Using Formal Charge • Formal Charge can help us determine the best Lewis Structure when there are options • Consider the following two skeletal structures for CH2O. Which one is preferred?

  23. Formal Charge Rules • Lewis structures where all formal charges are zero is preferred • Small formal charges (0 and +/-1) are preferred to big formal charges (+/-2, +/-3, etc.) • The best arrangements are where the more electronegative atoms have the more negative formal charge

  24. Group Quiz #2 • Draw the Lewis Structures for the following compounds and determine the formal charge on EACH atom • SO32- • CO32-

  25. Section 6.5Resonance

  26. Resonance Structures • Consider the molecule NO3- and its Lewis Structure

  27. Section 6.6Exceptions to the Octet Rule

  28. Exceptions to the Octet Rule • Central atom has fewer than 8 v.e. due to electron shortage • Ex: Boron (happy w/ 6); Beryllium (happy w/ 2) • Central atom has fewer than 8 v.e. due to odd # of electrons (known as radicals) • Ex: Nitrogen (NO2) • Central atom has more than 8 v.e. • Ex: Sulfur (SF6) and Xenon (XeF4) • See pp. 201—204 for examples/explanations

  29. Formal Charges, Resonance Structures, AND Exceptions! • Consider the polyatomic ion: SO42-. What would be the BEST Lewis dot structure? • What about PO43-?

  30. Group Quiz #3 • Draw the Lewis Structure for antimony pentafluoride (SbF5) • Draw the Lewis Structure for Borane (BH3) • Draw the Lewis Structure for Nitrogen Disulfide (NS2)

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