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Ionic and Covalent Bonding

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Title: 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
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