Download structures the octet rule lewis dot structure time to try a few

16/09/2015
THE OCTET RULE
All elements want to get a full outer shell
Will lose or gain electrons to get to 8 electrons with the exception of
Hydrogen and Helium
atoms bond in order to achieve the same number of electrons as the
nearest noble gas in the periodic table.
bonded atoms are said to be isoelectronic to the nearest noble gas.
(All noble gases have 8 valence electrons except He.)
STRUCTURES
1.
LEWIS DOT STRUCTURE
Uses the octet rule
Represent the valence electrons of
an element
You can follow these rules…
2. Count all the valence electrons. – Determine the
total number of valence electrons in the compound.
3. Place two electrons (a sigma covalent bond)
between the central atom and each of the outer
atoms
4. Place "lone pairs" of electrons about each terminal
(outer) atom [except H atoms] to satisfy the octet rule.
Count how many electrons you have used to this point.
5. Place the extra electrons on the central atom(s) in
pairs.
6. If the central atom does not have an “octet”, try
forming double or even triple bonds until it does.
Decide how to arrange the various atoms with respect to one
another.
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Although there are no definite "rules" you can follow, you usually
do the following:

Go for maximum symmetry. For instance, if you have one of
Atom A and four of Atom B, put Atom A in the middle and
arrange the four B atoms symmetrically around it.

Carbon is always a central atom in a Lewis structure. In
compounds with more than one carbon atom, the carbon atoms
are joined in a chain.

Hydrogen atoms always go on the outside. Not too hard to
figure out why!!

Halogen atoms form only single bonds when oxygen is not
present, therefore they are usually placed on the “outside”.
When halogen atoms are on the “outside”, they are always
joined by single bonds and always obey the octet rule.

Oxy-acids are acids containing oxygen (how surprising!). Some
examples are HNO3, H2SO4, H3PO4 and HClO3. When writing
the Lewis structure for an oxy-acid, always bond the hydrogen(s)
to an oxygen.
TIME TO TRY A FEW
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16/09/2015
VSEPR THEORY
What it stands for: Valence Shell Electron Pair Repulsion
Theory
Predicts what the molecules will look like in 3D
Developed in 1957 by Canadian chemist, Ronald Gillespie
LONE PAIR
Based on the idea that the geometry of a molecule is determined primarily by repulsion among the pairs of electrons associated with the central atom. The pairs can be bonded or un‐bonded (lone pairs)
BONDED PAIR
Only valence electrons of the central atom influence the molecular shape
1) Pairs of electrons in the valence shell of a central atom repel each other.
2) These pairs of electrons tend to occupy positions in space that minimize
repulsions and maximize the distance of separation between them.
GRAB A MOLYMOD KIT!
-There are three kinds of repulsions between valence
electrons of a molecule
-LP LP repulsions are stronger than LP BP repulsions which
are stronger than BP BP repulsions
SHAPE – Tetrahedral
LONE PAIR – LONE PAIR (LP, LP)
ANGLE – 109.5º
LONE PAIR – BONDED PAIR (LP, BP)
- Used to show the 3D shape
of CH4
BONDED PAIR – BONDED PAIR (BP, BP)
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* Shapes of Covalent Molecules with NO Lone Pairs of
Electrons Around the Central Atom*
Bond Angle - 180º
# OF Electron Pairs – 3
Examples – BeCl2, CO2, HgCl2
Bond Angle – 109.5º
# of Electron Pairs – 4
Examples – CH4, ClO4, PO4, SO4
Bond Angle - 90º
# of Electron Pairs – 6
Examples – SF6
Bond Angle - 120º
# of Electron Pairs – 3
Examples – BCl3, BF3, SO3
Bond Angle – 90º + 120º
# of Electron Pairs – 5
Examples – PCl5
Planar triangular
Tetrahedral
Trigonal bipyramidal
Octahedral
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