Molecular Geometry

1
Molecular Geometry
2
Its all about the Electrons

• Electrons decide how many bonds an atom can have
• They also decide the overall shape of the
  molecule
• OPPOSITES ATTRACT!

3
Lewis Structures

• A Lewis structure is basically a diagram of how a
  molecule looks using dots to represent the
  electrons.
• There are 4 rules for making these structures and
  that is where the electrons come into play.

4
Rule number 1

• Count the Number of valence electrons!
• This means of all the atoms present
• With a polyatomic anion, add one for each
  negative charge
• With a polyatomic cation, subtract one for each
  positive charge
• Ex CO2
• C 4 O 6 6 6 4 16

5
Rule Number 2

• Draw a skeleton structure for the molecule
  using all single bonds
• This will most often be one central atom with
  several surrounding ones
• Typically the central atom is written first
• Ex CO2 O-C-O

6
Rule Number 3

• Determine the number of valence electrons still
  available for distribution
• To do this simply deduct two valence electrons
  for each single bond written in step two
• Ex CO2 two single bonds so far so we
  subtract a total of 4
• 16 4 12

7
Fourth Rule

• Determine the number of electrons required to
  fill an octet for each atom
• If this equals the number of electrons left, then
  place them on the atoms as unshared pairs
• If the number of electrons available is less than
  the number needed then you need to make double or
  triple bonds in place of the single bonds

8
CO2 (again)

• O C O
• So far we have used four electrons so 16 4 12
• Carbon still needs 4 more electrons and each
  Oxygen needs 6 more. 6 6 4 16
• But we only have 12 left so lets make some double
  bonds!
• O C O becomes O C O

9

• Now Carbon doesnt need anymore electrons and the
  Oxygens only need 4 more each. Since we used 4
  electrons to make those into double bonds we now
  have exactly 8 electrons left.
• Now we simply distribute them to the Oxygen atoms
  as unshared paired electrons.

10
Practice!
11
Resonance!

• Resonance is invoked whenever a single Lewis
  structure does not adequately reflect the
  properties of a substance
• In other words, resonance comes into play when
  you can make two structures that are the same in
  their placement of atoms but different in the
  bonds
• SO2

12

• Resonance structures are NOT forms where the
  electrons move eternally between them
• Resonance structures are equally plausible or
  they are not a resonance structure
• Resonance forms differ in their distribution of
  electrons, NOT in their arrangement of atoms!
• So just because a formula for a compound is the
  same it does not mean that it is a resonance
  structure

13
VSEPR

• Lewis structures tell us how the atoms are
  connected to each other.
• They dont tell us anything about shape.
• The shape of a molecule can greatly affect its
  properties.
• Valence Shell Electron Pair Repulsion Theory
  allows us to predict geometry

14
VSEPR

• Molecules take a shape that puts electron pairs
  as far away from each other as possible.
• Have to draw the Lewis structure to determine
  electron pairs.
• bonding
• nonbonding lone pair
• Lone pair take more space.
• Multiple bonds count as one pair.

15
Electronegativity

• Electronegativity is a measure of how much an
  element wants to pull electrons towards itself
• This is represented as a unit-less number ranging
  from 0 4.0
• Heres a handy reference sheet with all the
  values. Guard it with your LIFE!

16
So what?

• These numbers can be used mathematically to know
  if a bond is ionic or covalent
• It can also tell you if a covalent bond is more
  polar or less polar (more on polarity in a
  minute)
• So all we have to do is subtract one from the
  other.

17
Example

• Fluorine has an electronegativity of 4.0
• Sodium has an electronegativity of 0.9
• 4.0 0.9 3.1
• So what does that mean?
• It means that it is an ionic bond!
• This makes sense since we know that a bond
  involving one metal and one non-metal is ionic.

18
Example two

• Fluorine has an electronegativity of 4.0
• Carbon has an electronegativity of 2.15
• 4.0 2.5 1.5
• This makes this bond covalent!

19
Sharing is caring, but some elements are greedy!

• This greediness shown by some elements like
  fluorine leads us to the next piece of this
  puzzle
• The more unequal the sharing of electrons is in a
  bond, the more polar it is.
• The smaller that difference in electronegativity,
  the less polar.

20
Polar vs Non-polar

• Polar
• Number greater than 0.4
• Unequal sharing of electrons
• Water is an example
• Non-Polar
• Number less than 0.4
• Equal sharing of electrons
• Methane (CH4) is an example

21
So why is this important?

• Polarity is a major component of organic
  chemistry
• Polarity also explains why certain substances can
  dissolve other substances while others cannot
• Think oil and water