Chem. 1-2 Chapter 8
Molecular Shape
Formulas
• Molecular Formula:
– Specifies how many atoms are in a single
molecule of a compound:
– Ex. Glucose is C6H12O6
• Empirical Formula:
– Specifies the ratios in which molecules are
present.
– Ex. Glucose empirical formula is C3H6O3
– You simply divide until a subscript cannot be
divided by a whole number
Formulas…..
• Structural Formula:
– Shows the arrangement of atoms in a
molecule… The Lewis structures you have
been doing.
• However, structural formulas don’t give
information on the shape of the molecule.
• This information is often shown with ball
and stick models
Shapes
• Shapes are most often symmetrical.
• Why?
• Valence electrons will arrange themselves
as far away from each other as possible.
This is called the Valence-Shell Electron
Pair Repulsion Theory (VSEPR)
– Exceptions to this are molecules involving
transition metals.
• There are five main shapes of molecules
Linear
• A straight line.
• All molecules with only 2 atoms.
• But CO2 is too…..
– This is because the valence electrons
try to get as far away from each other as
possible.
– No unshared central electrons.
• Bond angle is 180°
Trigonal Planar
• “Triangular laid out flat”
• Example BCl3
• Generally when a central atom has 3
surrounding atoms AND the central atom
has no unshared pairs of electrons.
• Bond angle is 120°
Tetrahedral
• “4 surfaced shape”
• Ex. CH4
• It is like a pyramid with a flagpole on the
top.
• It is not a flat shape because the electrons
can be further apart (at an angle of 109.5°)
using a 3D spread instead of 90° as a flat
shape.
Pyramidal
• Specifically a 3-sided base (not like the
Egyptian pyramids)
• When we have a central atom with
unshared electron surrounded by 3 other
atoms.
• Ex. NH3
• Bond angle is 107°
• Unshared pairs repel stronger than bond
pairs
Bent
• Example : H2O
• Again, the unshared pairs exert a greater
repulsion force. So the 2 H atoms are
separated by a slightly smaller angle than
one would expect.
• Bond angle is 105°
Home work
• Write the Structural Formula (Lewis
Structure) and predict/draw the shape of
the following molecules:
• CCl4, HCN, BF3, BeCl2, H2S, NF3, PCl3,
OF2, SiO2, CF4
• Due Tomorrow
Bond Length
• The distance between nucleii of atoms in a
molecule.
• Two trends….
– Triple bonds are shorter than double bonds
which are shorter than single bonds.
– As you move down a group, atoms are
_________, therefore bond length is _______
Polarity
• Polar molecules remember from ch 7?
• Polar molecules form Dipoles, meaning
two poles – the molecule has a negative
and a positive end.
• If a molecule has only non-polar (less than
0.4 E.N. difference) bonds it is a non-polar
molecule.
• Strangely, if a molecule contains polar
bonds, it is not necessarily polar. If
symmetry cancels out the polarity, the
molecule is non-polar.
Polarity….
• The shape of a molecule AND the polarity
of its bonds determine molecule polarity.
• We must look again at electronegativity…
the ability of an atom to attract electrons in
bonds.
• Remember, electronegativity increases
going left to right in a period.
• Electronegativity increases going up a
group.
• See table on page 241
Why is polarity important?
• Polar molecules only dissolve polar
molecules
• Non-polar molecules only dissolve nonpolar molecules.
How do we figure out if a
molecule is polar?
• Draw ball and stick model.
• Look at the electronegativity difference
between each atom involved in a bond.
• Draw arrows in the direction of highest
E.N. on each bond.
• The length of each arrow is proportional to
the E.N. difference across each bond.
• Imagine the arrows show an electron tug
of war….
• If the tug-of-war is a draw, the molecule is
non-polar. There is symmetry of
electronegativity.
• If there is not good symmetry, the
molecule is polar.
Polarity ….
• Lets take the examples of H2CO, CO2 and
H2O.