Download Lewis Structures

05/10/99
Ionic bond formation
2 peripheral Lewis’ bonded to a central Lewis
Ionic Compounds
• Arrows can be used to show electron transfer
• Electrons are transferred and accepted by atoms to
achieve the stable electron configuration of (become
isoelectronic with) noble gases.
Molecular Elements and
Compounds
• Ionic compounds are typically
formed between a metal and a
non-metal atom
• Electrons are transferred to the
atom with higher
electronegativity (usually from
the metal to the non-metal)
• The transfer of electrons forms
two ions of opposite charge
which are attracted to each other
by electrostatic forces
The ions arrange themselves in a 3
dimensional lattice so that attractions between
opposite charges are maximized and
repulsions between like charges are
minimized
Ethane Molecule
• Covalent bonds
typically form
between atoms of nonmetal elements
• Atoms share the
electrons in their outer
shells to achieve full
octets (or duets)
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Lewis Theory of Bonding
• Atoms and ions are stable if they have a full valence
shell of electrons
• Electrons are most stable when they are paired
• Atoms form chemical bonds to achieve a full valence
shell of electrons
• A full valence shell of electrons may be achieved by
an exchange of electrons between metal and a nonmetal atoms
• A full valence shell of electrons may be achieved by
the sharing of electrons between non-metal atoms
• The sharing of electrons results in a covalent bond
Duets, Octets and Exceptions to
the Rules
• Duet rule: Hydrogen forms stable molecular
configurations when it shares 2 electrons
Exceptions to the rule:
• Under-filled octets: molecules whose central
atoms are surrounded by fewer than 8 electrons
Lewis Structures
• Shows the arrangement of electrons and
bonds in a molecule or polyatomic ion,
Octet rule:
• non-metals will share electrons in order to
form an outer shell of 8 electrons
– Carbon, nitrogen, oxygen, fluorine and other
halogens always obey the octet rule
Exceptions to the rule:
• Overfilled octets: molecules whose central atoms
are surrounded by more than 8 electrons
• Consequently, the central atom is electron deficient and
reacts readily with electron rich molecules
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Non-Octet Compounds
The four possibilities for non-octet compounds are:
1. Where more than 4 atoms are bonded to the
central atom such as PCl5, SF6.
2. A noble gas is participating in bonding such
as XeF4.
3. Where the central atom has less than 8
valence electrons such as BH3.
4. Where molecules contain an odd number of
nonbonding electrons such as NO.
Drawing Lewis Structures
• Because there are exceptions to the octet
rule, we need a set of rules to determine
how many electrons surround atoms in a
molecule
N O
Steps for Drawing Lewis Structures
Skeleton Structure
1. Identify the central atom (skeleton structure)
– Usually the element with the highest bonding
capacity (largest number of unpaired electrons),
or the highest electronegativity, or the element
that there is only one of in the chemical formula
– Write the symbol for the central atom, then
arrange the symbols of the atoms for the rest of
the elements around it
Example: Lewis Structure
• Draw the Lewis Structure for CH2O
• Total valence electrons in atoms: C=4, H=1,
O=6
Steps for Drawing Lewis Structures
Count Electrons
2. Add up the total number of valence
electrons in the atoms of each element in
the molecule
– This number represents the total number of
electrons (dots) you will draw in the
structure
– Add electrons if you are drawing a
polyatomic anion
– Subtract electrons if you are drawing a
polyatomic cation
Example: Polyatomic Cation
• Polyatomic ions, NH4+
– Total electrons N = 5, H=1, subtract one for 1+
charge
(1 x 5) + (4 x 1) -1 = 8 electrons, therefore there are 8
electrons in the structure of the NH4+ ion
(1 x 4) + (2 x 1) + (1 x 6) = 12 electrons,
therefore there are 12 electrons in the structure
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Example: Polyatomic Anion
• Polyatomic anion, PO43– Total electrons, P = 5, O = 6, 3 electrons for 3charge
• (1 x 5) + (4 x 6) + 3 = 32 electrons,
therefore there are 32 electrons in the PO43structure
Steps for Drawing Lewis Structures
Electron Pairs in Bonds
3. Place one pair of electrons between the
central atom and each of the surrounding
atoms (ligands)
– Each pair of electrons represents a single
covalent bond
Steps for Drawing Lewis Structures
Steps for Drawing Lewis Structures
Octet for peripheral atoms
Electrons available
4. Place pairs of the remaining valence
electrons as lone pairs on the surrounding
atoms (not the central atom)
– Follow duet rule for H, and octet rule for
other atoms
5. Determine how many electrons are still
available by subtracting the number of
electrons used in the structure so far, from
the total number of valence electrons.
Total
Electrons
Electrons
- placed in
bonds
number of electrons
= needing placement as lone pairs
on ligands and/or the central atom
Steps for Drawing Lewis Structures
Steps for Drawing Lewis Structures
Remaining electrons on central atom
Create multiple bonds if necessary
6. Place the remaining electrons on the central
atom in pairs.
7. If the central atom does not have an octet,
move lone pairs from the surrounding atoms
into a bonding position with the central atom
until all octets are complete.
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Lewis Structure
vs.
Structural Formula
Steps for Drawing Lewis Structures
Add excess electrons to central atom
8. If the surrounding atoms have complete
octets and there are electrons remaining,
add these electrons as lone pairs to the
central atom.
Steps for Drawing Lewis Structures
Check final structure
Check the final structure:
 All surrounding atoms (ligands) except H should
have a complete octet counting shared (bonded)
and lone pair electrons
The central atom is the only atom in the structure
that can be under-filled or over-filled
Polyatomic ions should be surrounded by large
square brackets and the charge indicated on the
top right
Placing electrons around atoms
Compound
1) Skeleton
Structure
2) Count
electrons
3) Electron
pairs in
bonds
4) Octet for
peripheral
atoms
5) Remaining e–s on
center atom
6) Create
multiple
bonds?
Final
structure
• Lewis Structures have 2 dots representing a
bond
• Structural formulas have a line representing
a bond
• Structural formulas are easier to see and
use!
Draw Lewis Structures for the
following:
• ClO2• CO
Placing electrons around atoms
Compound
1) Skeletal
Structure
2) Count
electrons
3) Electron
pairs in
bonds
4) Octet for
peripheral
atoms
5) Remaining e–s on
center atom
6) Create
multiple
bonds?
Final
structure
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Resonance structures
Co-ordinate Covalent Bonds
• Lewis structures for certain atoms do not
match experimental observations
• For example, the bond lengths of CHO2–
predicted by the Lewis structure are incorrect
–
O
H
C
O
• The double CO bond should be shorter, and
possess a greater bond energy (due to the
higher concentration of e–s in a double bond)
• Yet, experimentally, both bonds are the same
• The reason is due to “resonance”
• Bonds that occur when both electrons in a
bond are donated by a single atom
[
• Investigate how the coordinate covalent
bond is formed in NH4+
]
Resonance structures
• A resonance structure can be drawn for any molecule
in which a double bond can be formed from two or
more identical choices
• Resonance structures can be drawn 2 ways…
C
O
[
] [
O
H
C
O
H
–
H
]
C
O
2
[
O
O
1
]
Examples of Resonance
Structures
–
–
• Resonance structures have a structure intermediate to
what we can depict with a Lewis Structure, therefore we
need to show all possibilities in the structure
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