Ionic and Covalent Bonding

1
Ionic and Covalent Bonding

• Chapter 9
• Ebbing Textbook

2
Ion Sizes

• Cations Shrink
  
• Anions Increase

Ba
Ba2
Ba
Cl
Cl-
3
Periodicity and Electron Configurations

• Which of the following is the smallest ion? and
  WHY?
• S-, Fe2 Ca2, Be

4
Ions Want to Form a Closed Shell Too!

• S- Ne3s23p4 1 electron Ne3s23p5
• This is equivalent to Cl
• Fe2 Ar4s23d6 - 2 electrons
  Ar4s23d4 Ar4s13d5
• This is equivalent to Cr
• Ca2 Ar4s2 - 2 electrons Ar
• Be He2s2 - 1 electron He2s1

5
Lewis Structures A Crude Predictive Model

• By concentrating on the valence electrons we
  can predict something about the reactions of
  elements.
• Using electron-dot structures we can envision how
  electrons are shared in poly-atomic molecules.
• Every element wants to have a nobel gas
  configuration. -- OCTET RULE

6
Lets Go Back to Our Examples

• H2
• 2 H. HH

7
OR

• What about NaCl?

Cl
Na
8
OR

• What about NaCl?

Cl
Na
9
Polarity in a Molecular Bond

• An unequal sharing of electrons
• IONIC bonding

Cl
Na
10
Polarity in a Molecular Bond
Cl-
Na
11
Another Example - Cl2

• A non-polar bond
• COVALENT - equal sharing of electrons

Cl-
Cl-
12
Electronegativity

• Linus Pauling - Two Nobel Prizes
• The ability of an atom in a molecule to attract
  shared electrons to itself

EN is Electron Affinity D is the difference in
Electronegativity Fluorine is Defined as 4.0
13
Electronegativity and Bond Type
14
Concept Check

• Rank the following diatomics from most covalent
  to most ionic.
• SH, HCl, HF, HH, OH
• H-HgtS-HgtCl-HgtO-HgtF-H 0.0 0.4 0.9
  1.4 1.9

15
What About O2?

• The number of electrons doesnt add up.
• HOW CAN WE FORM TWO OCTETS?

2
12 electrons / 2 6 pairs
O
16
What About O2? The Solution

• Double Bonds

O
O
17
Writing Lewis Structures

• Determine the total number of valence electrons
  in the species by adding together the numbers of
  valence electrons of each atom.
• Remember Column number gives the number of
  valence electrons for the atom.

18
Writing Lewis Structures

• If the species is an anion (has a negative
  charge) add the charge number to the overall
  number of electrons.
• If the species is an cation (has a positive
  charge) subtract the charge number from the
  overall number of electrons.

19
Writing Lewis Structures

• Place the atoms in their relative positions with
  element around the central atom as far apart as
  possible.
• Draw a line representing a single bond,
  containing two electrons, between the atoms.

20
Writing Lewis Structures

• Distribute the remaining electrons evenly in
  pairs on the outer atoms so that each atom has
  eight electrons. (Remember that hydrogen is an
  exception and only has two electrons)
• Subtract the number of electrons used from the
  total and place the remaining electrons on the
  central atom. Pair them if possible.

21
Writing Lewis Structures

• If the central atom has fewer than eight
  electrons, move a sufficient number of non-bonded
  pairs of electrons from the outer atoms,other
  than hydrogen and halogens, to between the atoms
  such that the central atom has at least eight
  electrons.

22
Lewis Formulas

• Write the Lewis formula for P2.

23
Lewis Formulas

• Write the Lewis formula for P2.

Total valence electrons 2 x 5 10.
The skeleton is P-P. Distribute the remaining
electrons symmetrically
24
Lewis Formulas

• Write the Lewis formula for P2.

Total valence electrons 2 x 5 10.
The skeleton is P-P. Distribute the remaining
electrons symmetrically
..
..
P____ P
25
Neither P atom has an octet. Make one lone pair
from each P a bonding pair. P
P
26
PCl5
27
PCl5
P 5 Cl 5 x 7 40 valence electrons
Cl
Cl
P
Cl
Cl
Cl
28
Beryllium Chloride

• BeCl2

29
Beryllium Chloride

• BeCl2

Be
Cl
Cl
30
Basic Exceptions

• B and Be often have less than the octet.
• a complete 2s shell is stable
• 3rd row and higher can exceed the octet.
• d-orbitals are available
• If you must exceed the octet, place extra
  electrons on the central atom with a charge.

31
Oxidation States

• When refering to a molecule, we can refer to the
  oxidation state of each atom.

The oxidation state is equal to the charge on an
ion if a noble configuration is assumed for each
atom.
32
Got it?

• NaCl
• Na oxidation state 1
• Cl oxidation state -1
• NH3
• N oxidation state -3
• H oxidation state 1

33
What about O3?

• 3 x 6 valence electrons 18

34
What about O3?

• 3 x 6 valence electrons 18

O
O
O
35
Resonance

• Resonance is a representation of the fact that
  electrons are not stationary in a molecule.

O
O
O
O
O
O
36
Resonance

• Resonance is a representation of the fact that
  electrons are not stationary in a molecule.

O
O
O
O
O
O
Or
O
O
O
37
What Do We Predict About the Bond Structure in O3?

• Based on our resonance model
• Bond lengths should be the same
• Molecule should be bent
• Only one kind of bond energy

38
Formal Charges

• If more than one Lewis structure is possible and
  it is not certain which should be the central
  atom (sometimes there is not a central atom),
  formal charge is used to derive the most likely
  molecular structure.
• Formal charge is a theoretical assessment and
  does not represent an actual charge on the
  species.

39
Formal Charges

• Subtract the number of formal valence electrons
  assigned to an atom in a molecule from the number
  of electrons normally around the atom.
• Formal Charge
• Valence electrons - local molecular electrons

40
Formal Charges

• The number of valence electrons is the same as
  the group number of the element.
• Local molecular electrons 1/2 the number of
  shared electrons the number of unshared
  electrons.

41
Formal Charges

• Calculate the formal charge for each element in
  all possible structures. The structure closest
  to a formal charge of zero is the most likely
  structure.
• Also, within a structure, the most
  electronegative element usually has the most
  negative formal charge.

42
Example of Formal Charge
2-

• SO42-

O
S
O
O
O
Formal Charge on S 6 - (1/2 x 12)0 Formal
Charge on O 6 - (6 1/2x2) -1
OR 6
- (41/2x4) 0
43
If More Than One Lewis Structure Exists

• Atoms in molecules try to achieve formal charge
  0 first.
• Negative formal charges are expected to reside on
  the most electronegative atoms.

44
A Problem to Consider

• Consider the Lewis structures that may exist for
  XeO3. Use the concept of formal charges to
  predict the structure.

45
Formal Charge and XeO3

• Assign the most effective Lewis dot structure
  based on formal charges.

(3)
Xe
O
O
O
(-1)
(-1)
(-1)
46
Formal Charge and XeO3

• Assign the most effective Lewis dot structure
  based on formal charges.

(0)
Xe
O
O
O
(0)
(0)
(0)
47
Bond Length, Order and Energy

• Use bond energies to estimate ?H for the
  following gas-phase reaction. H H
  H H
  H H H H

Br
H
C
-
C
H
Br
C
C

This is called an addition reaction, because a
compound is added across the double bond.
48
An Addition Reaction
H
H
H
Br
C
C

H
H
H
H
H
C
C
Br
H
H
...
49
Enthalpy of ReactionFrom Bond Energies (BE)

• In the reaction, a CC double bond is converted
  to a C-C single bond. An H-Br bond is broken,
  and one C-H bond and one C-Br bond are
  formed. ?H BE(CC) BE(H-Br) -
  BE(C-C) - BE(C-H) -BE(C-Br)
• (602 362 - 346 - 411 - 285) kJ
• -78 kJ