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Understanding Molecular Structures: Octet Rule, Electronegativity, Lewis Diagrams

Explore the Octet Rule, Electronegativity, and Lewis Structures in chemistry. Understand how atoms achieve noble gas configurations, determine bond types, and grasp concepts of formal charges and resonance structures.

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Understanding Molecular Structures: Octet Rule, Electronegativity, Lewis Diagrams

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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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