Molecular Compounds Covalent Bonding

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Molecular Compounds(Covalent Bonding)

• CPC Ch. 8
• Science with Mr. A.

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Molecules and Molecular Cpds

• The sharing of electrons is covalent bonding.
• When the atoms bond they form a molecule.
• Each molecule is an electrically neutral group of
  atoms joined together by covalent bonds.
• Combined molecules that form compounds are
  molecular compounds.
• Usually this is 2 or more non-metals that form a
  compound with a low MP/BP.
• These molecular compounds are written out by a
  molecular formula.

3
Single Covalent Bonds

• SCB have one pair of shared electrons.
• All Halogens form SCBs like H2.
• The electron dot diagram for this is HH.
• The two electrons in the bond can be replaced
  with a line (H-H).
• The diagram with lines for bonds such as
• H-H is called a structural formula.

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Single Covalent Bonds

• Any valence electrons that are not paired are
  called unshared pair.
• Such is the case for NH3

• Atoms will form double and triple covalent bonds
  if it allows them to attain the Nobel Gas
  configuration.
• Ex N2

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N
H
H
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H
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Coordinate Covalent Bonds
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• Carbon Monoxide bonds differently than water,
  ammonia, and carbon dioxide.
• Carbon needs four electrons, and Oxygen needs
  two.
• C 1s2 2s2 2p2
• O 1s2 2s2 2p6

• Here, in the CCB, one atom contribute BOTH the
  electrons needed for the bond.

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O
C
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C
O
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Coordinate Covalent Bonds

• Most polyatomic ions contain both Covalent and
  Coordinate Covalent Bonds.
• Ex SO32-

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S
O
O
O


2 e-
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2-
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S
O
O
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O
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Bond Dissociation Energy

• When H2 forms, heat is given off.
• This suggests that the product is more stable.
• Breaking the Cov. Bond of 1 mole of H2 requires
  435KJ of energy.
• The energy required to break a bond is the Bond
  Dissociation Energy.

• The large the energy the stronger the Covalent
  Bond.
• Ex
• C-C 347 KJ
• CC 657 KJ
• CC 908 KJ

8
Resonance

• When a molecule has two or more possible
  structures, it shows resonance.
• Here the electrons can move around.
• Ex
• O O O -----gt O O O

9
Bonding Theories

• Molecular Orbitals
• When two atoms combine, their orbitals overlap
  forming a Molecular Orbital.
• Atomic Orbitals belong to the atom.
• Molecular Orbitals belong to the molecule.
• s s sigma
• Horizontal p Horizontal p sigma
• Vertical p Vertical p pi

10
VSEPR Theory

• Valence Shell Electron Pair Repulsion Theory.
• Remember that a 2-D drawing for CH4 is really a
  3-D molecule.
• The like charges of Hydrogen atoms on CH4 will
  repel each other.

• The angles between the Hydrogen atoms will vary.
• Ex
• CH4 109.5o
• H2O 105o
• NH3 107o

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Hybrid Orbitals

• VSEPR Theory deals with the shape of the molecule
  only.
• Hybrid Orbital deals with shape and type of bond.
• Hybrid Single Bond
• CH4 C has an e- config of 1s2 2s2 2p2
• It changes to 1s2 2s 2p3
• Sp3 has an angle of 109.5o
• Sp2 has an angle of 120o
• Sp has an angle of 180o
• Sp bonds always form a tetrahedron.

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Polar Bonds

• Cov. Bonds share, and by sharing the atoms all
  pull on each other.
• When they pull evenly it is a Non-Polar Cov.
  Bond.
• When they pull unevenly it is a Polar Cov. Bond.
• The more electronegative atom attracts electrons
  and therefore has a negative charge.

• H Cl-
• H electroneg. 2.1
• Cl electroneg. 3.0
• HCl is polar
• A polar bond makes a polar molecule.
• A molecule with two different charged poles is
  called a Dipole.

13
Molecular Attractions

• Intermolecular attractions are weaker than Ionic
  or Cov. Bonds.
• Van der Waals are the two weakest molecular
  attractions
• Dipole Interactions polar charge of one
  molecule pulls on another molecule.
• Dispersions Forces normal movement of the
  electron can momentarily polarize a non-polar
  molecule.

14
Misc.

• Hydrogen Bonds
• Hydrogen is Cov. Bonded to a very electronegative
  atom already bonded to another electronegative
  atom.

• Intermolecular Attraction
• Network solids are very stable solids where all
  atoms have Cov. Bonds.

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The End