Lewis Dot Diagrams _LDDs_ by dffhrtcv3

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									Lewis Dot Formulas (LDFs)

             Chapter 7

Take Home Test # 4 due next Tuesday
            10/24/06
         HW Due 10/17/06
• Read Chapter 7

• P 300, # 2, 4, 6
• Also, give me LDFs for all elements in
  groups 1, 2 and 13 -18. If you use more
  than one color, Extra credit!
All atoms have electrons organized
          in energy levels
• 7 possible energy levels
  – Named by numbers: 1 through 7
  – Electrons at EL 1 are closest to nucleus, at EL
    7 they are farthest away

• Each energy level has 1 or more sublevels
  – Named by letters: s, p, d, f
  – Each holds different # electrons:
     • s (2), p (6), d (10), f (14)
          Valence electrons
• The electrons in the energy level farthest
  from the nucleus (s and p sublevels)

• These are the electrons that participate in
  chemical reactions (they are shared,
  gained or lost)

• We will only look at elements from groups
  1, 2, 13 – 18
Lewis Dot Formulas of atoms
• Show one element’s valence electrons
• First 4 are drawn alone in a side of an
  imaginary square around element’s
  symbol:


• Second 4 are filled in as pairs of the first 4
  electrons:
          Lewis Dot Formulas
            of compounds
• Show all valence electrons in a compound
• Covalent
  – All elements want to be like noble gases, who
    have full octet (8) of valence electrons
  – When elements get into compounds, it is to
    share/lose or gain electrons to feel like noble
    gases.
     • But! Why will they never be exactly like nobles?
  – Use this information to draw LDD’s for
    compounds
LDFs of Covalent Compounds
When there are just two atoms:
• Count total # of available valence
  electrons from all elements = A
• Count total # of valence electrons needed
  from all elements to achieve stable octet =
  N
• S = # electrons shared  S = N – A
LDFs of Covalent Compounds
When there are just two atoms:
• Place S electrons as dashes
  – Each dash = a shared pair of 2 electrons
• Place any leftover electrons as unshared
  pairs (as dots) so that each element (that
  can) has full octet

• Examples: CO, Br2, Cl2
           HW For Monday
• P 301: 14, 34, 36, 40, 42, 52
LDFs of Covalent Compounds
Arranging atoms when more than 2 atoms:
• Center atom usually least electronegative
  – Electronegativity = how much an element wants
    electrons. Ranges from 0.8 – 4.
     • Page 250 = table of electronegativity values
• Oxygen usually not bonded to itself
  – Exceptions: O2, O3, H2O2
• In ternary acids, hydrogen usually does not bond
  to central atom (bonds to oxygen)
  – Exceptions: phosphoric and phosphorous acid
LDFs of Covalent Compounds
When more than 2 atoms:
• Calculate S by subtracting Ntotal – Atotal
• Place S electrons as dashes
  – Each dash = a shared pair of 2 electrons
• Place any leftover electrons as unshared
  pairs (as dots) so that each element (that
  can) has full octet
• Total # of electrons (dashes plus dots)
  should = A
                Hydrogen!
• Only needs 2 electrons to fill its valence
  energy level. (1s)
LDFs of Covalent Compounds
When more than 2 atoms:
• Examples:
  – NH3         – C3H8
  – CO2         – H2SO4
  – CH4         - C2H4
  – CS2
  – CHCl3
  – CH2O
          HW For Tuesday
• P 301 # 14, 34 and 36 d, 40 and 52
              Exceptions
• See section in your book for section on
  exceptions to octet rule.
      Lewis Dot Diagrams
        of compounds
• Ionic:
  – Between a metal and a non-metal
  – Put brackets and write charge around each
    ionic species and do each separately
  – Give/take away electrons from each species
    according to their charges
  – Don’t show electrons of positive cation
• Examples:
  – NaCl       – SrO
  – LiCl       – MgCl2
        Lewis dot formulas of
          polyatomic ions
• Put brackets & write charge on outside
  right hand corner
• If negative charge, add value to A
• If positive charge, subtract value from A
• Examples
  – CO32-
              Resonance
• When there is more than one possible
  Lewis dot formula, molecule or ion will be
  an “average” of all possible formulas.
• Resonance = lowers energy = more stable
  compound
Examples:
• NO3-         – CO32-
• SO2
• O3
       Checking your Lewis Dot
              Formula
Formal Charge!
    FC = # valence – [(# bonds) + (# unshared)]
          electrons                 electrons

• Σ of FC in a covalent molecule = zero
• Σ of FC in an ion = charge of the ion

• Negative FC on more electronegative elements
• FC represented by + or -
                 Radicals
• Compounds or elements where A (total
  valence electrons) is an odd #
• Energetically unstable!
• Radical elements  reactive with other
  elements to form compounds where A is
  an even #
  – Groups 1, 13, 15, 17
• Radical compounds: odd electron will be
  unshared & unpaired on the radical
  element (and Σ FC ≠ 0)
  – NO & NO2

								
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