The Intermolecular Forces (forces between molecules)

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The Intermolecular Forces
(forces between molecules)
In general, the weaker the
intermolecular forces, the less
energy which is required for the
substance to undergo a physical
change ( change in state).
– Thus the physical properties of melting point
and boiling point gives a rough measure of
the intermolecular forces which bind the
molecules to each other.
• The higher the melting and boiling points,
the higher the strength of the
intermolecular forces which bind the
molecules together.
• The intermolecular forces:
1. Forces between covalent molecules
– London dispersion forces (Van der Waals
forces)
– Dipole-dipole attraction
– Hydrogen bonding
2. Ionic bonding
3. Metallic bonding
4. Network covalent bonding
1. Covalent molecule forces
a) London Dispersion Forces: are the
dominant forces between covalently
bonded, non-polar molecules.
Ex. N2, O2, CH4
• formation of “ instantaneous dipoles”
• weak and act over a short distance.
• Force increases with the increase in the
number of electrons.
London Dispersion Forces
• Small molecules generally have a low
melting and boiling points ( molar mass of
less than 50, generally will be a gas at
room temp)
• -covalent compounds are usually gases,
liquids or low melting solids.
Force varies with the shape of the molecule.
• A long molecules has many electrons exposed
and therefore a relatively high boiling point. The
structure is flexible and does not stack well, so it
has a low melting point.
• A more compact molecule has fewer exposed
electrons and will have a low boiling point but
stacks better, so its melting point will be higher.
Ex Normal pentane (C5H12)
mp -130°C bp 36°C
Neopentane (C5H12)
(2,2 dimethylpropane)
mp -20°C
bp 9°C
b) Dipole-Dipole Attraction: is a force which
acts between polar molecules (ex: H2S ).
It results from the attraction of opposite
poles of permanent dipoles.
- polar molecules will generally have
higher melting and boiling points than nonpolar molecules.
c) Hydrogen Bonding: occurs when a lone
pair on a O, N or F atom forms a “bond”
to the partial positive charge of a H atom
(part of a polar covalent bond).
- molecules will have high boiling and
melting points.
Figure 4.4
The hydrated proton
Relative magnitudes of forces
The types of bonding forces vary in their
strength as measured by average bond
energy.
Strongest
Covalent bonds (400 kcal)
Hydrogen bonding (12-16 kcal )
Dipole-dipole interactions (2-0.5 kcal)
Weakest
London forces (less than 1 kcal)
2. Ionic bonding:
generally occurs in compounds of metals and nonmetals. Ex. NaCl, K2SO4, NH4Cl.
It is the result of the attraction of the opposite
charges.
-formation of a crystal lattice.
-ionic solids are called crystals and are very
brittle and hard
-ionic solids have very high melting and boiling
points.
3. Metallic bonding:
Most metals have only one or two valence
electrons and low ionization energies.
The valence electrons do not seem to belong to
any individual atom but move easily from one
atom to another.
Metals can be thought of as positive ions
immersed in a “sea” of mobile electrons.
The attractive forces that bind metals together are
called metallic bonds.
Ex: Hg, Cu, Au, Fe, alloys.
-high electrical and thermal conductivity
-shiny surfaces are due to the de-excitation of
the electrons (the result of valence electrons
absorb and re-emit light energy)
-strength varies with the nuclear charge of the
metal atoms and the number of electrons in the
metal’s sea of electrons
-metals can be flattened out (malleable) or
stretched out into a wire (ductile) because the
electrons can move into other positions without
breaking up the essential structure.
-Very low to very high melting point
-Very low to very high boiling point
4. Network Covalent Bonding:
Occurs between two non-metals
-network is infinite
-includes diamond, graphite, quartz and most
rocks.
-covalent is stronger than other bonds due to a
hard network.
-solids are brittle
-network solids have high melting (1200 –
3500°C) and boiling points.

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