Covalent Bonding

1
Covalent Bonding

• Covalent bonds versus ionic bonds
• Electronegativity and polarity
• Drawing Lewis (electron dot) structures
• Resonance
• Formal charge
• VSEPR theory - Molecular geometry
• Hybridization
• Sigma and Pi bonds

2
Ionic vs Covalent Bonding
3
The Covalent Bond

• Covalent bonds are formed by sharing at least one
  pair of electrons.
• Every covalent bond has a characteristic length
  that leads to maximum stability.
• This is the bond length.

4
The Covalent Bond

• Energy required to break a covalent bond in an
  isolated gaseous molecule is called the bond
  dissociation energy.

5
The Covalent Bond
6
Ionic vs Covalent Bonding
7
Polar Covalent Bonds

• Bond polarity is due to electronegativity
  differences between atoms.
• Pauling Electronegativity is expressed on a
  scale where F 4.0

8
Polar Covalent Bonds

• As a general rule for two atoms in a bond, we can
  calculate an electronegativity difference (?EN )
  ?EN EN(Y) EN(X) for XY bond.
• If ?EN lt 0.5 the bond is covalent.
• If 0.5 lt ?EN lt 2.0 the bond is polar covalent.
• If ?EN gt 2.0 the bond is ionic.

9
Polar Covalent Bonds

• Using electronegativity values, predict whether
  the following bonds are nonpolar covalent, polar
  covalent, or ionic
• SiCl4 CsBr FeBr3 CH4
• HCl CCl4 NH3 H2O

10
Electron-Dot Structures (Lewis)

• The electron-dot structures provide a simple, but
  useful, way of representing chemical reactions.
• Ionic
• Covalent

11
Drawing Lewis-Dot Structures

• Determine the total valence electrons in the
  molecule
• Determine bonds needed by
• Bonds in molecule (total electrons needed to
  fill all valence shells actual valence
  electrons)/2
• Draw the structure by placing the least
  electronegative atom (other than H) in the center
  and attach other atoms using correct of bonds.
• Complete the valence shells of the terminal atoms
  -
• Put remaining pairs on central atom.
• CHECK YOUR ANSWER by counting electrons and
  valence shells

12
Drawing Lewis-Dot Structures

• NH2F Amino Fluoride In this molecule,
  nitrogen is the central atom.
• Number of valence electrons 5 (2 x 1) 7
  14 e-
• Bonds needed (2288)-(1157)/2 6/2 3
  bonds

Rule 3 Rule 4/5 Rule 5
13
Drawing Lewis-Dot Structures

• Draw electron-dot structures for
• CO C2H4 H2O2 Cl2CO
• O3

14
Resonance Structures

• When multiple structures can be drawn, the actual
  structure is an average of all possibilities.
• The average is called a resonance hybrid. A
  straight double-headed arrow indicates resonance.

15
Resonance Structures

• The nitrate ion, NO3, has three equivalent
  oxygen atoms, and its electronic structure is a
  resonance hybrid of three electron-dot
  structures. Draw them.
• Draw as many electron-dot resonance structures as
  possible for SO2, CO32, HCO2 , SO42, PO43.

16
Drawing Lewis-Dot Structures

• SF6
• Some molecules with central atoms below the
  second row sometimes expand their valence
  shell
• In this compound there are 10 valence electrons
  on bromine this is called an expanded octet.
  The extra pairs go into unfilled d orbitals.
  (can t predict of bonds with these atoms!!!!)

17
Formal Charge

• Formal Charge Helps determine the best resonance
  structure.
• FC ( Valence e- in free atoms) (valence e-
  in molecular atom)
• (valence e- in molecular atom) ( unbonded e-
  bonds)
• Draw two different Lewis structures for SO2 and
  do formal charges on each atom
• Determine which is more stable

18
Molecular Shapes VSEPR

• The approximate shape of molecules is given by
  Valence-Shell Electron-Pair Repulsion (VSEPR).

19
Molecular Shapes VSEPR

• Step 01 Count the total electron groups
• Step 02 Arrange electron groups to maximize
  separation.
• Electron groups are collections of bond pairs
  between two atoms or a lone pair. (a single bond
  is a group, a lone pair is a group, a double bond
  is 1 group).
• Groups do not compete equally for spaceLone
  Pair gt Triple Bond gt Double Bond gt Single
  Bond

20
Molecular Shapes VSEPR

• Two Electron Groups Electron groups point in
  opposite directions.

21
Molecular Shapes VSEPR

• Three Electron Groups Electron groups lie in the
  same plane and point to the corners of an
  equilateral triangle.

22
Molecular Shapes VSEPR

• Four Electron Groups
• Electron groups point to the corners of a
  tetrahedron.

23
Molecular Shapes VSEPR
Five Electron Groups Electron groups point to
the corners of a trigonal bipyramid.
24
Molecular Shapes VSEPR

• Six Electron Groups Electron groups point to the
  corners of a octahedron.

25
Molecular Shapes VSEPR Summary
26
Molecular Shapes VSEPR

• Draw the Lewis electron-dot structure and predict
  the shapes of the following molecules or ions
• H2O H3O CF4
• PF6 SiCl4
• XeF2 ICl4

27
Valence Bond Theory

• 1. Covalent bonds are formed by overlapping of
  atomic orbitals, each of which contains one
  electron of opposite spin.
• 2. The electron pair in the overlapping orbitals
  is shared by both atoms.
• 3. The greater the amount of orbital overlap, the
  stronger the bond.

28
Valence Bond Theory

• Think about CH4
• 4 equal bond lengths,
• 4 equal bond angles,
• 4 equal bond strengths
• Requires 4 equal types of bonds equal orbitals,
  but C has an s and 3 p orbitals???????????
• Linus Pauling - Wave functions from s orbitals
  p orbitals could be combined to form hybrid
  atomic orbitals.

29
Valence Bond Theory

• sp3 hybrid

30
Valence Bond Theory

• sp hybrid
• sp2 hybrid

31
Valence Bond Theory

• sp3d hybrid
• sp3d2 hybrid

32
Valence Bond Theory

• s bonds
• p bonds

33
Molecular Orbital Theory

• The molecular orbital (MO) model provides a
  better explanation of chemical and physical
  properties than the valence bond (VB) model.
• Atomic Orbital Probability of finding the
  electron within a given region of space in an
  atom.
• Molecular Orbital Probability of finding the
  electron within a given region of space in a
  molecule.

34
Molecular Orbital Theory

• Additive combination of orbitals (s) is lower in
  energy than two isolated 1s orbitals and is
  called a bonding molecular orbital.

• Subtractive combination of orbitals (s) is
  higher in energy than two isolated 1s orbitals
  and is called an antibonding molecular orbital.

35
Molecular Orbital Theory

• Molecular Orbital Diagram for H2

36
Molecular Orbital Theory

• Molecular Orbital Diagrams for H2 and He2

37
Molecular Orbital Theory

• Bond Order is the number of electron pairs shared
  between atoms.
• Bond Order is obtained by subtracting the number
  of antibonding electrons from the number of
  bonding electrons and dividing by 2.