Intermolecular Versus Intramolecular Force Comparison
Date: Winter 2011-2012
Why are intramolecular forces of attraction stronger than intermolecular ones?
Wow! There are so many ways to think about and answer this question. The simplest answer is that if the intermolecular forces were stronger it would form a bond creating a compound.
All I can think is
that is mostly by definition.
Or else it is too big to answer.
There are several discernibly different attraction mechanisms between atoms and/or molecules.
Whatever ones are strongest, they pull atoms together firmly into molecules and we call them intramolecular.
The much weaker ones remaining will weakly try to pull molecules together and we call them intermolecular.
There are sometimes forces with strength between those two ranges;
they actually blur our definition of a molecule sometimes.
Adduct formation comes to mind, Hydrogen-bonding is not quite strong enough to be in that range. (In my view.)
I guess the fact that covalent bonds usually have well-defined saturation numbers
(I.E., hydrogen makes one and only one covalent bond, with one and only one neighbor...)
tends to make molecules clearly defined.
But they are not always.
The list of distinguishable chemical attractions I tend to think of:
binary covalent bonding [ intra ]
ternary covalent bonding (boron hydrides) [ intra ]
higher-number de-localized covalent bonding (benzene ring) [ intra ]
metallic sea-of-electrons bonding (Is a lump of metal a single molecule?)
(How do you choose to relate it to more-localized covalent bonding?)
[ intra, I think, at least by strength ]
ionic bonding [ intra ? ]
coordination bonding [ intermediate ]
hydrogen bonding (might be modeled as a disadvantaged form of ternary covalent bonding)
[ inter-molecular, or intermediate ]
polar-polar attraction [ inter ]
polar-induced attraction [ inter ]
quantum-uncertainty-induced polar attraction (better known as Van der Wall forces) [ inter ]
( I suppose I might have missed one or two.)
I cannot begin to explain all their distinctions in a short way.
I do not know how to explain why covalent bonds exist and are stronger than Van der Waal forces.
It seems obvious but I just have not thought it through. Perhaps someone else could.
Or why electron_affinity of some species is stronger than
the ionization_potential of other species,
enabling an electron to be stolen, and causing ionic bonding to happen.
Maybe you could think about those as you learn.
One might distinguish 3 levels of bonding instead of two:
2) equilibrium-stable compounds, (this would include hydrated salt crystals and adduct gasses)
3) bulk phases and their adhesion,
At the root is the simplest context:
if the prevailing thermal energy is equivalent to the bond energy or greater,
rearrangements occur, at whatever level that is.
If not, the bonding is firm, of whatever type.
It is difficult to predict and systematize the quantum-mechanical results of electron orbitals.
So we have somewhat messy facts of life shaping our chemical world.
We organize them as best we can.
The words "intermolecular" and "intramolecular" are part of that effort.
This is not ALWAYS the case, but as you have noted, it is the most
common condition. The reason is that most all intra-molecular -- bonds
between atoms and/or molecular fragments -- are the result of sharing (or
pairing) of electrons. These are the common types of chemical bonds, and the
type of bond referred to when you say a chemical bond is formed.
Intermolecular bonds are formed between molecular fragments. These might be
hydrogen bonds, like those formed between molecules of water, pairing of
electric dipoles, and some types of bonding such as between ions. For
example, the heat of sublimation of LiF is 147 kJ/mol. This is the energy
required to sublime solid LiF into Li(+1) + F(-1). However, the dissociation
energy of LiF, the energy required to dissociate diatomic LiF into Li(0) +
F(0) -- atomic gaseous lithium and a fluorine atom is 577 kJ/mol. These are
different chemical processes. Diatomic gaseous LiF is a polar diatomic
molecule, but the atoms are bound together by the formation of an electron
pair. This is a stronger bond than the dissociation product of LiF into a
pair of ions.
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Update: June 2012