This slideshow provides a brief overview of molecular polarity and dipole-dipole forces in chemistry. With easy-to-understand examples and visual representations, your understanding of polarity and intermolecular forces should improve drastically.
Introduction
History
Definition
Types of H bond
Hydrogen bond in water
Bifurcated and over - Coordinated hydrogen bond in water
Hydrogen bonds in DNA and proteins
Hydrogen bonds in polymers
Systematic hydrogen bond
Importance of hydrogen bond
Conclusion
References
Introduction
History
Definition
Types of H bond
Hydrogen bond in water
Bifurcated and over - Coordinated hydrogen bond in water
Hydrogen bonds in DNA and proteins
Hydrogen bonds in polymers
Systematic hydrogen bond
Importance of hydrogen bond
Conclusion
References
A non-covalent interaction differs from a covalent bond in that it does not involve the sharing of electrons, but rather involves more dispersed variations of electromagnetic interactions between molecules or within a molecule.
Intermolecular attractions are attractions betwee.pdfannaelctronics
Intermolecular attractions are attractions between one molecule and a neighbouring
molecule. The forces of attraction which hold an individual molecule together (for example, the
covalent bonds) are known as intramolecular attractions. These two words are so confusingly
similar that it is safer to abandon one of them and never use it. The term \"intramolecular\" won\'t
be used again on this site. All molecules experience intermolecular attractions, although in some
cases those attractions are very weak. Even in a gas like hydrogen, H2, if you slow the molecules
down by cooling the gas, the attractions are large enough for the molecules to stick together
eventually to form a liquid and then a solid. In hydrogen\'s case the attractions are so weak that
the molecules have to be cooled to 21 K (-252°C) before the attractions are enough to condense
the hydrogen as a liquid. Helium\'s intermolecular attractions are even weaker - the molecules
won\'t stick together to form a liquid until the temperature drops to 4 K (-269°C). van der Waals
forces: dispersion forces Dispersion forces (one of the two types of van der Waals force we are
dealing with on this page) are also known as \"London forces\" (named after Fritz London who
first suggested how they might arise). The origin of van der Waals dispersion forces Temporary
fluctuating dipoles Attractions are electrical in nature. In a symmetrical molecule like hydrogen,
however, there doesn\'t seem to be any electrical distortion to produce positive or negative parts.
But that\'s only true on average. The lozenge-shaped diagram represents a small symmetrical
molecule - H2, perhaps, or Br2. The even shading shows that on average there is no electrical
distortion. But the electrons are mobile, and at any one instant they might find themselves
towards one end of the molecule, making that end -. The other end will be temporarily short of
electrons and so becomes +. Note: (read as \"delta\") means \"slightly\" - so + means \"slightly
positive\". An instant later the electrons may well have moved up to the other end, reversing
the polarity of the molecule. This constant \"sloshing around\" of the electrons in the molecule
causes rapidly fluctuating dipoles even in the most symmetrical molecule. It even happens in
monatomic molecules - molecules of noble gases, like helium, which consist of a single atom. If
both the helium electrons happen to be on one side of the atom at the same time, the nucleus is
no longer properly covered by electrons for that instant. How temporary dipoles give rise to
intermolecular attractions I\'m going to use the same lozenge-shaped diagram now to represent
any molecule which could, in fact, be a much more complicated shape. Shape does matter (see
below), but keeping the shape simple makes it a lot easier to both draw the diagrams and
understand what is going on. Imagine a molecule which has a temporary polarity being
approached by one which happens to be entirely non-polar just at .
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2. Intermolecular Forces
• the strongest intermolecular forces exist between polar
molecules
– polar molecules act as dipoles
– these dipoles are short range
– they have equal but opposite charges
– positive to negative poles (the arrow’s head points towards
the more electronegative of the two atoms in the bond)
EX. H Cl
3. Dipole-Dipole Forces
• dipole-dipole – attraction between polar molecules
• polar molecules have greater boiling points than
nonpolar molecules
– this is because nonpolar molecules are held together by
weaker forces, called van der Waals forces
– since the dipole-dipole forces that hold polar molecules
together are stronger than van der Waals forces, it is harder
for their bonds to be broken by a boiling liquid
4. Dipole-Dipole Forces Cont.
• polar molecules have a partially charged positive end
and a partially charged negative end
• polar molecules can also induce a dipole in nonpolar
molecules because polar molecules can temporarily act
as electrons
– however, these dipoles are weaker than dipole-dipole forces
5. What Do Dipole-Dipole Forces
Look Like?
the idea of “opposites attract” may not be applicable to
real life relationships, but in dipole-dipole forces, this is the
basic principle
!
!
in this figure, the dipole-dipole forces that exist between
atoms in the chemical compound iodine monochloride are
set up in a way where the negative and positive poles are
connected, not negative-negative or positive-positive
6. But Why do Opposites Attract?
• in dipole-dipole interactions, both ends of the pole (“di”
means two, so “dipole” means two ends of one pole)
have partial charges
– the two “complete” each other, resulting in a neutral charge