The Octet Rule

1
The Octet Rule

• Except for hydrogen and helium, atoms are
• most energetically stable if they have a
• completely filled valence shell.
• A completely filled valence shell is called an
• octet because eight electrons are involved.
• Atoms will form ionic compounds by gaining
• or losing valence electrons or covalent
• compounds by sharing electrons.

2
Lewis Structures

• A simple way to show the valence electrons
• present in an atom.
• Valence electrons are those electrons found
• in the highest numbered principal energy
• level (PEL).
• Valence electrons are found only in the s and
• p sublevels and in most cases are the
• electrons responsible for bonding.

3
Lewis Structures

• The chemical symbol represents the kernel of
• the atom.
• The kernel of an atom consists of the nucleus
• and the core electrons.
• Example
• X

Start with s and proceed cw using Hunds Rule!
s
pz
px
py
4
Lewis Structures Second Period
Li Be B C N O F Ne
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Drawing Lewis Structures

• Count the number of valence electrons.
• Remember, the number of valence
• electrons can be determined by the
• group (family) numbers.
• Draw a skeleton structure for the covalent
• structure assuming single (sigma) bonds.
• Terminal atoms will usually be hydrogen,
• oxygen, and halogens.

6

• Determine the number of valence electrons
• that are left.
• Distribute the remaining electrons to the
• atoms surrounding the central atom to satisfy
• the octet rule.
• Distribute the remaining electrons as pairs
• around the central atom.

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• When all of the valence electrons have been
• used, ensure that the central atom has an
• octet.
• If the central atom does not have an
• octet, form one or more double or triple
• (pi) bonds.
• To form multiple bonds, move one or
• more pairs of electrons from a
• surrounding atom to the bond
• connecting the central atom.

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Lewis Structure Examples

• NH3 ammonia
• 1 N --- 5 val e-
• 3 H --- 3

H
N
H
H
H
8
H
N
- 6
H
2
shared pair
unshared pair
9

• OF2 oxygen difluoride
• 1 O --- 6 val e-
• 2 F --- 14

O
F
F
F
20
- 4
16
O
F
F
- 12
4
F
F
O
10

• HCN hydrogen cyanide
• 1 H --- 1 val e-
• 1 C --- 4
• 1 N --- 5

H C N
H C N
10
- 4
H C N
6
H C N
A triple bond is needed for both C and N to have
a complete octet.
11

• NH4 ammonium

1 N --- 5 val e- 4 H --- 4 -1

H
N
H
H
H
8
-8
Square brackets are used for anions and cations.
0
A 1 is subtracted because a positively charged
polyatomic ion has 1 less electron.
12
O
O

• O2 oxygen

2 O --- 12 val e-
O
O
-2
10
O
O
13

• C2H5OH ethyl alcohol (ethanol)

2 C --- 8 val e- 6 H --- 6 1 O --- 6
H
H
H
H
C
C
O
H
H
20
-16
H
H
4
H
H
C
C
O
H
H
14

• N2H4 hydrazine

2 N --- 10 val e-
N
N
H
H
4 H --- 4
H
H
14
-10
N
N
H
H
4
H
H
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Resonance

• When there is more than one equivalent
• Lewis structure, all the structures are given
• to represent the molecule or the polyatomic
• ion.
• Resonance can only occur when all the
• structures
• satisfy the octet rule
• have the same type and number of bonds

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Resonance
O
O
N

• NO2- nitrite
• 1 N --- 5 val e-
• 2 O --- 12
• 1
• 18
• - 4
• 14
• -14
• 0

N
O
N
O
O
O
N
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Resonance

• - -
• O N O
• Note that the placement of the atoms in these
• structures is the same but the electron
• arrangement is different.

O
N
O
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• Note that polyatomic ions are placed inside
• square brackets and the charge is placed
• after the brackets as a superscript.
• It is necessary to include all the structures
• separated by a double-headed arrow.

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Exceptions to the Octet Rule

• Three major exceptions to the octet rule
• Molecules or ions with more than eight
• electrons around the central atom.
• Species with fewer than eight electrons
• around the central atom.
• Species with an odd number of valence
• electrons.

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Expanded Octets

• Starting with period three, atoms have the
• capability to accommodate d electrons (3d).
• AX4E molecules such as SF4 are able to
• accommodate 4 bonding pairs of electrons
• and one nonbonding pair of electrons. This
• results in S being surrounded by 5 electron
• pairs.

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• The favored bonding scenario includes large
• central atoms (starting in the third period)
• and small terminal atoms such as fluorine,
• chlorine, and oxygen.
• As shown below, S also has the ability to
  accommodate six pairs of valence electrons as
  found in SF6.

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Less Than an Octet

• Molecules having either boron or beryllium as
  their central atom result in the central atom
  having only 2 or 3 valence pairs of electrons.
• These molecules are very reactive with a molecule
  having an unshared pair of
• electrons.

BF3
BeCl2
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Odd Number of Valence Electrons

• Most molecules have an even number of
• valence electrons.
• In rare cases, molecules such as NO and
• NO2, there is one unpaired electron which is
• very reactive.

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Odd Number of Valence Electrons

• nitrogen(II) oxide nitrogen(IV) oxide

N
O
O
N
O
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Electronegativity

• Electronegativity is a measure of the
• attraction of an element for a shared pair of
• electrons.
• H Cl H Cl
• Comparing the electronegativity values of
• hydrogen and chlorine, chlorine has a value
• of 3.2 and that of hydrogen is 2.2.

d-
d
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Range of Ionic Character

• Ionic 100 50 5 0
• Character
• ?EN
• 0.0 0.3 Nonpolar Covalent
• 0.4 1.7 Polar Covalent
• 1.8 4.0 Ionic

4.0
0.3
0.0
1.7
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Properties of Ionic Compounds

• The building blocks of ionic compounds are
• cations (positive ions) and anions (negative
• ions).
• The force holding an ionic compound
• together is the ionic bond.
• Properties
• have only empirical formulas

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• high melting and boiling points
• solids at room temperature (25C, 298K)
• conductors in the liquid phase (l) or in an
• aqueous solution (aq)
• some are water soluble, some are not
• check solubility chart

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Nonpolar Covalent Compounds

• The building blocks of nonpolar covalent
• (molecular) compounds are molecules.
• The force holding a nonpolar molecule
• together is the covalent bond.
• Properties
• have both empirical and molecular
• formulas

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• low melting and boiling points
• often a gas or a liquid at room
• temperature (25C, 298K)
• nonconductors of heat and electricity
• insoluble in water
• exists as discrete molecular units

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Polar Covalent Compounds

• The building blocks of polar covalent
• (molecular) compounds are molecules.
• The force holding a polar molecule together
• is the covalent bond.
• Properties
• have both empirical and molecular
• formulas

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• usually higher melting and boiling points
• often a gas or a liquid at room
• temperature (25C, 298K)
• some are conductors of heat and
• electricity
• more likely to be water soluble