Chemical Bonding:

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Chapter 5

• Chemical Bonding
• Covalent Bond Model

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Sec 5.1 Covalent Bond Model

• Covalent bonding is different from ionic bonding
  in several ways
• Covalent bonds are usually between two nonmetals,
  even the same element
• Covalent bonds involve sharing of electrons
  rather than transfer of electrons
• Covalent compounds are discrete molecules as
  units (ie not crystals)
• Covalent compounds can be solid, liquid, or gas
  at room temperature

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Sec 5.1 Covalent Bond Model

• Covalent bond a chemical bond resulting from
  elements sharing electrons
• The orbitals overlap so that both elements in the
  bond satisfy the octet rule
• Example H2, Cl2

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Sec 5.2 Lewis Structures

• Drawing Lewis structures for covalent bonds
• Dots stand for unshared electrons
• When two electrons are shared between elements,
  we indicate a the bond with a line
• Differentiate between bonding pairs and
  nonbonding pairs
• Example H2, Cl2, CH4

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Sec 5.2 Lewis Structures
Figure 5.1 page 100
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Sec 5.2 Lewis Structures

• Exceptions to the octet rule
• Boron has 3 valence electrons, it can be stable
  with 6 electrons (3 bonds)
• Sulfur can be stable with 8, 10, or 12 valence
  electrons due to its size
• Hydrogen needs only 2 electrons
• BF3

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Sec 5.3 Single, Double, Triple

• More than 2 electrons can be shared between two
  elements
• 2 electrons is called a single bond
• 4 electrons is called a double bond
• 6 electrons is called a triple bond
• There is no quadruple bond
• Examples O2, CO2, N2, HCN

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Sec 5.4 Determining Bonds Formed

• In general the number of covalent bonds an
  element can form from the valence
• Carbon with 4 electrons can form 4 total bonds
• Nitrogen with 5 forms 3 total bonds
• Oxygen with 6 forms 2 total bonds
• Halogens (F, Cl, Br, I) and Hydrogen form 1 bond.
  These elements will never be involved in double
  or triple bonds

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Sec 5.5 Coordinate Covalent Bonds

• In general skip this concept

Figure 5.3 Page 105
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Sec 5.6 Drawing Lewis Structures

• Tactics for completing Lewis Structures
• Find the total valence electrons for the
  molecular (add valences for each atom)
• Determine arrangement of atoms or work from given
  backbone
• Connect atoms with single bonds and subtract
  those e- from the total valence
• Leave lone pairs untouched, and fill in other
  valance e- to fit octet rule (the hard part)

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Sec 5.6 Drawing Lewis Structures

• Practice makes it much easier
• Rules to remember
• Oxygen will never have more then 2 bonds (1
  double or 2 single)
• Hydrogen and Halogens will never be in the middle
  of a chain, they always cap an end
• Dont break up pairs on elements such as Oxygen,
  Phosphorus, Sulfur, Halogens

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Sec 5.6 Drawing Lewis Structures
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Sec 5.7 Bonding with Ions

• Slightly different from the book
• We examine the bonding within polyatomic ions
• If negative, remember to add an extra electron
  (for example NO2-)
• If positive, remember to remove an electron (for
  example NH4)
• Then the polyatomic associates with the opposite
  ion (book example of K2SO4)

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Sec 5.8 Molecular Geometry

• VSEPR Valence Shell Electron Pair Repulsion
• Method of predicting molecular shape based on the
  number of bonding and non-bonding electron pairs
• Bond angles arise due to electron repulsion

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Sec 5.8 Molecular Geometry

• Three Primary Classes for Shape
• Tetrahedral 4 regions of electron density
• Trigonal Planar 3 regions of e- density
• Linear 2 regions of e- density
• Double and triple bonds are one density
• Subclasses are based on the number of lone pairs
• Example CF4 vs. PF3

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Sec 5.8 Molecular Geometry
Chemistry at a Glance Page 112
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Sec 5.9 Electronegativity

• How can we tell if the bond is covalent or ionic
  ?
• Electronegativity a measure of an atoms
  attraction for electrons
• Metals have lower EN, want to give up electrons
• Nonmetals have higher EN, wanting to gain
  electrons

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Sec 5.9 Electronegativity
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Sec 5.9 Electronegativity
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Sec 5.10 Bond Polarity

• In general, if the elements have a difference in
  electronegativity of 1.9 or greater, the bond is
  ionic.
• If the difference in electronegativity is less
  than 1.9, the bond is covalent
• Check metals and nonmetals to confirm

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Sec 5.10 Bond Polarity

• Even with covalent bonds, there can be unequal
  sharing of electrons
• If the difference in the electronegativity is
  between 0.5 and 1.9 the bond is polar
• Polar covalent bond means that the bond is not
  even, there is a partial positive pole and a
  partial negative pole
• These poles are indicated with a and sign and
  the lowercase greek delta (d)

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Sec 5.11 Molecular Polarity

• In the same way that a bond can have uneven
  sharing making it polar, the overall molecule can
  also be the same
• This concept of polar molecules applies to
  covalent molecules (ionic is already known)
• There are two factors to consider
• Are the bonds in the molecule polar
• What is the geometry and does it have an effect

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Sec 5.11 Molecular Polarity

• In some cases, multiple polar bonds cancel out
  the opposite direction, rendering the molecule
  overall nonpolar
• Example CO2, CCl4 (Compared with CH3Cl)
• In other cases, a molecule is polar due to the
  lone pairs
• Example H2O, NH3
• In general, figure out the lewis structure and
  then determine polarity

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Sec 5.11 Molecular Polarity

• Figure 5.13
• Page 118
• The difference of
• CH4 and CH3Cl

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Sec 5.12 Naming Molecular Compounds

• Remember Binary Ionic Compounds?
• Binary Covalent Molecular compounds are slightly
  different in naming
• Remember binary means only 2 different elements
• When naming, use a prefix to announce how many
  atoms of each type are present in the compound

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Sec 5.12 Naming Molecular Compounds

• Prefixes for numbers
• Table 5.1 Page 119
• Note that these prefixes are common

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Sec 5.12 Naming Molecular Compounds

• Generally, give the prefix and name of the first
  element, then the prefix and stem of the second
  element with the suffix ide
• N2O3 would be dinitrogen trioxide
• CCl4 would be carbon tetrachloride
• Some compounds such as H2O, CH4, and NH3 go by
  common names but can be named by the same method
  as above

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Sec 5.12 Naming Molecular Compounds
Table 5.2 Page 119
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Problems

• Assigned problems pages 121 - 124
• 5.16, 5.17, 5.21, 5.22
• 5.24, 5.27, 5.29, 5.30
• 5.36, 5.39, 5.43, 5.51, 5.52, 5.53
• Practice Test page 124