Lewis Symbols of Atoms

1
Lewis Symbols of Atoms

• use symbol of element to represent nucleus and
  inner electrons
• use dots around the symbol to represent valence
  electrons
• put one electron on each side first, then pair
• elements in the same group have the same number
  of valence electrons

2
Lewis Bonding Theory

• atoms bond because it results in a more stable
  electron configuration
• atoms bond together by either transferring
  electrons(losing or gaining) or sharing
  electrons so that all atoms obtain an outer shell
  with 8 electrons like the noble gases
• Octet Rule

3
Lewis Symbols of Ions

• Cations have Lewis symbols without valence
  electrons
• (electrons are lost)
• Anions have Lewis symbols with 8 valence
  electrons
• (electrons gained)

4
Ionic Bonds

• metal to nonmetal, in ionic compounds
• metal loses electrons to form cation
• nonmetal gains electrons to form anion
• ionic bond results from to - attraction
• larger charge stronger attraction
• smaller ion stronger attraction
• Lewis Theory allow us to predict the correct
  formulas of ionic compounds

5
Covalent Bonds

• Form between two or more nonmetals, in molecular
  or covalent compounds
• Sharing of electrons
• Equally shared---nonpolar covalent bond
• Unequally shared----polar covalent bond
• Bond order single, double and triple

6
Single Covalent Bonds

• Nonpolar Polar

H


H
O


H
H

O


7
Double Covalent Bonds

• two atoms sharing two pairs of electrons
• 4 electrons
• shorter and stronger than single bond

8
Triple Covalent Bonds

• two atoms sharing 3 pairs of electrons
• 6 electrons
• shorter and stronger than single or double bond

9
Bonding Lone Pair Electrons

• Electrons that are shared by atoms are called
  bonding pairs
• Electrons that are not shared by atoms but belong
  to a particular atom are called lone pairs
• also known as nonbonding pairs

O S O


Bonding Pairs
Lone Pairs




10
Lewis Structures

• some common bonding patterns
• C 4 bonds 0 lone pairs
• 4 bonds 4 single, or 2 double, or single
  triple, or 2 single double
• N 3 bonds 1 lone pair,
• O 2 bonds 2 lone pairs,
• H and halogen 1 bond,

11
VSEPRValence Shell Electron Pair Repulsion

• Helps us predict the electronic and molecular
  shapes (geometries)
• The shape around the central atom can be
  predicted by assuming that the areas of electrons
  on the central atom will try to get as far from
  each other as possible in order to minimize the
  repulsions between electrons

12
Areas of Electrons

• Each Bond counts as 1 area of electrons
• single, double or triple all count as 1 area
• Each Lone Pair counts as 1 area of electrons
• Lone pairs take up slightly more space than
  bonding pairs

13
Some Geometric Figures

• Linear
• 2 atoms on opposite sides of central atom
• 180 bond angles
• Trigonal Planar
• 3 atoms form a triangle around the central atom
• Planar
• 120 bond angles
• Tetrahedral
• 4 surrounding atoms form a tetrahedron around the
  central atom
• 109.5 bond angles

14
Linear geometryTrigonal planar bent geometries
15
Tetrahedral, trigonal pyramidal and bent
geometries
16
Resonance Structures

• More than one Lewis structure can be drawn for a
  molecule
• CO2 OCO ? O-CO ? OC-O
• SO2 OS-O ? O-SO
• O3 OO-O ? O-OO
• SCN- S-CN ? SCN-
• CO32- NO2- NO3-

17
Exceptions to the Octet Rule

• Electron poor (deficient 4 or 6 electrons)
  elements Be, B.
• BF3, BeCl2
• Electron rich(expanded octet of 10 or 12
  electrons) elements period 3 and beyond period 3
  have vacant d orbitals S, P, Cl, Br, etc.
• SF4, SF6, SeF4, PF5,ClF5, IF5
• Radicals (odd number of electrons) NO, NO2

18
Shortcomings of Lewis theory

• Could not explain
• why metals conduct electricity
• how a semiconductor works
• why liquid oxygen is attracted into the
  magnetic field of a large magnet (oxgen is
  paramagnetic---contains unpaired electrons) NOT
  diamagnetic.

19
Electronegativity

• measure of the pull an atom has on bonding
  electrons
• increases across period (left to right)
• decreases down group (top to bottom)
• larger difference in electronegativity, more
  polar the bond
• negative end toward more electronegative atom

20
Bond Polarity
3.0-3.0 0.0

• .

4.0-2.1 1.9
3.0-0.9 2.1
covalent
ionic
non polar
polar
0
0.4
2.0
4.0
Electronegativity Difference
21

• .

polar bonds, but nonpolar molecule because pulls
cancel
polar bonds, and unsymmetrical shape causes
molecule to be polar
22
Adding Dipole Moments