A Molecular Comparison of
Liquids and Solids
• Physical properties of substances understood in terms
of kinetic molecular theory:
– Gases are highly compressible, assumes shape and volume
• Gas molecules are far apart and do not interact much with each
– Liquids are almost incompressible, assume the shape but
not the volume of container:
• Liquids molecules are held closer together than gas molecules,
but not so rigidly that the molecules cannot slide past each
– Solids are incompressible and have a definite shape and
• Solid molecules are packed closely together. The molecules are
so rigidly packed that they cannot easily slide past each other.
A Molecular Comparison
of Liquids and Solids
• Converting a gas into a liquid or solid requires the
molecules to get closer to each other:
– cool or compress.
• Converting a solid into a liquid or gas requires the
molecules to move further apart:
– heat or reduce pressure.
• The forces holding solids and liquids together are
called intermolecular forces.
• The covalent bond holding a molecule together
is an intramolecular forces.
• The attraction between molecules is an
• Intermolecular forces are much weaker than
intramolecular forces (e.g. 16 kJ/mol vs. 431
kJ/mol for HCl).
• When a substance melts or boils the
intermolecular forces are broken (not the
• Dipole-dipole forces exist
• Polar molecules need to be
• Weaker than ion-dipole forces.
• There is a mix of attractive and
repulsive dipole-dipole forces
as the molecules tumble.
• If two molecules have about
the same mass and size, then
dipole-dipole forces increase
with increasing polarity.
London Dispersion Forces
• Weakest of all intermolecular forces.
• It is possible for two adjacent neutral molecules to
affect each other.
• The nucleus of one molecule (or atom) attracts the
electrons of the adjacent molecule (or atom).
• For an instant, the electron clouds become
• In that instant a dipole is formed (called an
London Dispersion Forces
• Polarizability is the ease with which an electron cloud can
• The larger the molecule (the greater the number of
electrons) the more polarizable.
• London dispersion forces increase as molecular weight
• London dispersion forces exist between all molecules.
• London dispersion forces depend on the shape of the
• The greater the surface area available for contact, the
greater the dispersion forces.
London Dispersion Forces
• Special case of dipole-dipole forces.
• By experiments: boiling points of compounds with
H-F, H-O, and H-N bonds are abnormally high.
• Intermolecular forces are abnormally strong.
• H-bonding requires H bonded to an
electronegative element (most important for
compounds of F, O, and N).
– Electrons in the H-X (X = electronegative element) lie
much closer to X than H.
– H has only one electron, so in the H-X bond, the + H
presents an almost bare proton to the - X.
– Therefore, H-bonds are strong.
• Hydrogen bonds are responsible for:
– Ice Floating
Solids are usually more closely packed than liquids;
Therefore, solids are more dense than liquids.
Ice is ordered with an open structure to optimize H-bonding.
Therefore, ice is less dense than water.
In water the H-O bond length is 1.0 Å.
The O…H hydrogen bond length is 1.8 Å.
Ice has waters arranged in an open, regular hexagon.
Each + H points towards a lone pair on O.
The Octet Rule
most compounds, the representative elements in the
compounds achieve noble gas configurations”
• This statement is usually called the Octet Rule, because
the noble gas configurations have 8 e in their outermost
shells (except for He, which has 2 e).
• Many Lewis formulas are based on this idea
How to write Lewis structures for
molecules & polyatomic ions?
• Step 1 : Count the total number of valence
electrons in the molecule
Doesn‟t matter which element it comes from, just the total of all
• Step 2 : Set up a skeleton structure
least electronegative element, or carbon in center.
H is always at the outskirts.
Most of the time (but not always) halogen will also be at the
• Step 3 : Place the valence electrons until every
atom has a stable octet by:
i) drawing bond (either „ ‟ , „ ‟ or „ ‟)
ii) drawing lone pair (non-bonding pair of electrons) (:)
**Make sure every atom has a stable octet, which means,
every atom needs to have 8 valence electrons around it.
Except for H of course, which satisfy its stability with 2
**Don‟t forget to show all lone pair electrons.
Now, let’s practise writing Lewis
Draw Lewis structures for:
Exceptions to the Octet Rule
1) Electron deficient structures
Gaseous molecules containing either beryllium or boron as central atom ~ they have
fewer than eight electrons around Be or B atom.
Lewis structure for BeCl2 and BF3
Multiple bonds to the central atoms give unlikely structures:
Halogens are much more electronegative than Be or B.
• A molecule or polyatomic ion for which two or more Lewis
formulas with the same arrangements of atoms can be
drawn to describe the bonding is said to exhibit resonance.
• The three structures below are resonance structures of the
The relationship among them is indicated by the doubleheaded arrows.The double-headed arrow indicates
that the structures shown are resonance structures.
• Formal charge is the hypothetical charge on an
atom in a molecule or polyatomic ion.
• The concept of formal charges helps us to write
correct Lewis formulas in most cases.
• The most energetically favorable formula for a
molecule is usually one in which the formal charge
on each atom is zero or as near zero as possible.