I' Covalent Bonds and Reactions

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I. Covalent Bonds and Reactions

• Covalent Bond Model
• A Model is a simple explanation that lets us
  explain complicated things
• Chemical Bonds are a model that help us explain
  why groups of atoms behave as a unit in reactions
• Chemical Bond forms only if that arrangement is
  more stable
• Chemical Bond Energy energy gained by behaving
  as a unit
• Examples
• Methane (CH4)
• a. 4 C-H covalent bonds arranged in a tetrahedral
  shape
• b. The total energy of methane is 1652 kJ/mol
• c. We can estimate any C-H bond as being 413
  kJ/mol
• Chloroform (CH3Cl)
• Total E 1578 kJ/mol
• We can use C-H value
• C-Cl bond E 339 kJ/mol 1578 - 3(413) 339

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• Other bond energies
• Every C-H bond is not exactly 413 kJ/mol, but we
  estimate as such
• Other X-Y bond strengths have been estimated
• Single bond 2 atoms share one pair of e- (C-H)
• Double bond 2 atoms share 2 pairs of e-
  stronger than single bond
• CO 745, C-O 358
• Triple bond 2 atoms share 3 pairs of e- even
  stronger
• C?O 1072
• Bond length shortens as more e- are shared
  between the atoms
• Calculation Enthalpy of a Reaction
• Enthalpy DH sum of energy to break and form
  bonds in a reaction
• a. Bond breaking requires energy (endothermic).
• Bond formation releases energy (exothermic).
• ?H ?D(bonds broken) ? ?D(bonds formed)

energy required
energy released
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• Find DH for 1 H2 1 F2
  2 HF
• 1(432 kJ/mol) 1(154 kJ/mol) 2(565 kJ/mol)
  -544 kJ/mol DH
• 3. Sample Ex. 8.5 CH4 2Cl2 2F2 ----gt
  CF2Cl2 2HF 2HCl
• II. Lewis Structures
• Localized Electron Bonding Model
• A molecule is composed of atoms that are bound
  together by sharing pairs of electrons using the
  atomic orbitals of the bound atoms.
• Pairs of e- are either
• Shared in a bond (bonding pair)
• Localized on a single atom (lone pair)
• Duet and Octet Rules are in effect most stable
  form of any atom
• Lewis Structures
• Uses and to show arrangement of valence
  electrons in a molecule
• Ionic Compounds are straightforward

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• Covalent Compounds
• Shared electrons count for both atoms
• A line stands for one pair of electrons
• Rules for drawing Lewis structures
• Add up total valence e- for all atoms CH4
  4x1 4 8 e-
• Use one pair for each bond
• Arrange leftover e- to satisfy octet/duet
• Examples CCl4, CO2, H2O and Sample Ex. 8.6 N2,
  NH3, CF4, NO

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• 4. Exceptions to the Octet Rule
• a. 2nd row elements C, N, O, F always observe
  the octet rule.
• b. 2nd row elements B and Be (have very small
  size) often have fewer than 8 electrons around
  themselves - they are very reactive.
• c. 3rd row and heavier elements CAN exceed the
  octet rule using empty valence d orbitals.
• d. When writing Lewis structures, satisfy octets
  first, then place electrons around elements
  having available d orbitals.
• e. Sample Ex. 8.7 and 8.8 PCl5, ClF3, RnCl2

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• 5. Formal Charge
• a. The difference between the number of valence
  electrons (VE) on the free atom and the number
  assigned to the atom in the molecule.
• b. We need
• a. VE on free neutral atomfrom its group
  in periodic table
• b. VE belonging to the atom in the
  molecule
• One e- for each bond
• All the lone pairs electrons
• c. Examples
• i. NO
• ii. NO2-
• iii. CO2

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• III. Resonance Forms multiple correct Lewis
  structures for a given molecule
• A. Example NO3-
• B. Rules
• 1. Bracket the set of structures and separate
  with ??
• 2. Only electron pairs move to interconvert
  atoms dont move
• Electron Pushing
• C. Resonance Hybrids The true structure is a
  mixture of all resonance forms
• 1. Every resonance form contributes to the
  overall structure
• 2. None of the resonance forms are true
  pictures of the structure
• The real structure is not interconverting between
  resonance structures, but is a composite all of
  them, all of the time
• Sample Ex. 8.9 NO2-

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• IV. VSEPR Valence Shell Electron-Pair Repulsion
  
• Electron pairs want to be as far apart from one
  another as possible
• This applies to bonding pairs and lone pairs
  alike
• Steps in Applying VSEPR
• Draw the Lewis Structure
• Count atoms and lone pairs and arrange them as
  far apart as possible
• Determine the name of the geometry based only on
  where atoms are

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• B. Complications to VSEPR
• Lone pairs count for arranging electrons, but not
  for naming geometry
• Example NH3 (ammonia)
• Lone pairs are larger than bonding pairs,
  resulting in adjusted geometries
• Bond angles are squeezed to accommodate lone
  pairs

Trigonal pyramidal
Tetrahedral
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• b. Lone pairs must be as far from each other as
  possible
• 4. Double and triple bonds are treated as only 1
  pair of electrons

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5. Names for and examples of complicated VSEPR
Geometries
linear
trigonal planar
bent
trigonal pyramidal
tetrahedral
bent
T-shaped
linear
trigonal bipyramidal.
see-saw
square pyramidal
octahedral
square planar