Covalent Bonding Chapter 8

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Covalent BondingChapter 8

• Chemistry 2

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Molecular Compounds 8.1
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Molecules and Molecular Compounds 8.1

• Covalent bond SHARE e-
• Molecule neutral group of atoms joined by
  covalent bond
• Diatomic 2 atoms O2
• Molecular Compound compound composed of
  molecules CO or CO2
• Lower mp bp than ionic compounds
• Normally 2 or more nonmetals

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Molecular Formula 8.1

• Chemical formula
• H20 or CO2 or O2
• 1 is omitted if there is only 1 atom
• No structure or arrangement of atoms
• Use diagrams

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The Nature of Covalent Bonding 8.2
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The Octet Rule in Covalent Bonding 8.2

• Share electrons to attain e- conf. of noble gas
  8 e-
• Combination s of Group 4A, 5A, 6A 7A likely to
  form covalent bonds

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The Electron Probability Distribution for the H2
Molecule
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Single Covalent Bonds 8.2

• 2 atoms held together by single pair of e-
• 2 dots or H-H by each other represent
• Structural Formula H-H
• Represent bonds and arrangement
• Unshared pair valence e- that is not shared

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Double or Triple Covalent Bonds 8.2

• Share 2 pairs or 3 pairs of e-

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Practice Problem

• Write a lewis structure for CCl2F2
• Step 1 Arrange Atoms (Carbon is Central
  Atom because is has the lowest group number and
  lowest electronegativity.
• Step 2 Determine total number of valence
  electrons
• 1 x C(4) 2 x Cl(7) 2 x
  F(7) 32
• Step 3 Draw in valence electrons
• Step 4 Draw single bonds in replace of 2
  electrons between 2 atoms and subtract 2 e- for
  each single bond (4 x 2 8) so 32 8 24
  remaining

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Coordinate Covalent Bonds 8.2

• Covalent bond in which one atom contributes both
  bonding e-
• Molecular Formula CO
• Structural Formula C O
• Polyatomic Ion tightly bound group of atoms
  that has a or charge and behaves as a unit
• H attaches to NH3s unshared e-
• LOOK at page 225 SO3-2

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Bond Dissociation Energies 8.2

• E required to break the bond between 2 covalently
  bonded atoms
• H H ? H2 gives off large amount of heat
• Product more stable than reactants
• Big b.d.e. strong covalent bond normally
  unreactive
• C-C 347 kJ/mol
• C C 657 kJ/mol
• C C 908 kJ/mol

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Resonance

• 2 or more possible e- dot structures
• No back and forth changes actually occur
• Just a way to vision
• Drawing
• Must adhere to octet rule
• Sigma bonds not altered, pi and nonbonding e- are
  altered

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Exceptions to the Octet Rule

• Can occur when odd number of valence e-
• Atom requires less than octet of 8 e-
• BF3-NH3
• Some expand octet to 10 or 12 (esp w/ P and S

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Bonding Theories 8.3
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Molecular Orbitals 8.3

• Orbitals overlap
• REMEMBER atomic orbitals are orbitals in s,p,d,f
• Bonding orbital molecular orbital that can be
  occupied by 2 e- of covalent bond
• SIGMA BONDS s
• 2 atomic orbitals combine
• Directly between 2 nuclei
• Single bonds
• p overlaps end to end
• Pi Bonds p
• 2nd bond of double bond, 2nd and 3rd bond of a
  triple bond (sigma is 1st of double)
• Makes up 2 lobes
• Tend to be weaker than sigma bonds
• Orbital overlapping is less

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VSEPR Theory

• Valence-shell electron pair repulsion theory
  explains 3-D shape

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Hybrid Orbitals

• Provided info about molecular bonding and
  molecular shape
• Atomic orbitals mix to form same total number of
  equivalent hybrid orbitals
• Single Bonds
• CH4 sp3

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http//www.chemguide.co.uk/atoms/bonding/covalent.
html
http//www.mikeblaber.org/oldwine/chm1045/notes/Ge
ometry/Hybrid/Geom05.htm
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Hybrid Orbitals

• Double Bonds

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Molecule of electron pairs Shapes with, and without non-bonding e pair Hybridization of central atom
BeH2 2 linear, linear sp
BF3 3 trigonal planar, trigonal planar sp2
CH4 4 tetrahedral, tetrahedral sp3
NH3 4 tetrahedral, trigonal pyramidal sp3
H2S 4 tetrahedral, bent sp3
PF5 5 trigonal bipyramidal, trigonal bipyramidal dsp3
BrF3 5 trigonal bipyramidal, T-shaped dsp3
TeCl4 5 trigonal bipyramidal, Seesaw dsp3
SF6 6 octahedral, octehedral d2sp3
XeF4 6 octahedral, square planar d2sp3
XeF2 5 trigonal bipyramidal, linear dsp3
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Polar Bonds and Molecules 8.4
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Bond Polarity 8.4

• Nonpolar covalent bond equally share electrons
• Polar Covalent bond unequal sharing
• The more electronegative, the more strongly pulls
  on e-
• Less electronegative atom slightly d charge
• More electronegative atom slightly d- charge

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• Use table 6.2 in Chapter 6 for electronegativity
  of elements
• HCl
• H 2.1
• Cl 3
• Electronegativity .9
• Conceptual Problem Page 239 30-31

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Polar Molecules 8.4

• Often in a polar bond ?One end of molecule is
  slightly and other end slightly
• Call DIPOLE
• Ex HCl

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Attractions Between Molecules 8.4

• Intermolecular attractions weaker than ionic or
  covalent bonds but they are important!!! HOW?
• Determine if solids, liquids, and gases
• Surface tension
• Van der Waals ForceS
• Dipole Interactions polar molecules attracted
  to one another
• Similar to ionic but weaker
• Dispersion Forces caused by motion of e-
• Temporarily attractive force that results when
  the e- in 2 adjacent atoms occupy positions that
  make them temporarily dipole
• Weakest of all interactions
• Occurs even in non-polar

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Hydrogen Bonds 8.4

• Attractive forces in which a hydrogen covalently
  bonded to a very electronegative atom is weakly
  bonded to unshared e- pair

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Intermolecular Attractions and Molecular
Properties 8.4

• Physical properties depends on type of bonding
• Melting/boiling point lower for covalent
  compared to ionic
• Few covalent bonds have high mp
• Most network solids (crystals) solids in which
  all the atom s are covalently bonded
• Ex diamond
• Each C is attracted to 4 other Cs