Download Classifying Intermolecular Forces

Classifying Intermolecular Forces
In general, intermolecular forces can be divided into several categories. The four prominent types are listed
below and are, in general, listed from strongest to weakest. However, you must realize some forces from the
“weakest” category may actually be stronger than those listed in the stronger categories!
Covalent and ionic attraction/bonding are classified as intramolecular forces rather than intermolecular forces.
However, they are mentioned here because some solid substances are formed due to covalent and ionic bonding
between each of the particles. Examples: a covalently bonded crystal of diamond or sand (SiO2) and ionically
bonded crystals of NaCl or other salts. These solids are composed of atoms attached via covalent or ionic bonds
to all other atoms in the solid. Thus, the intermolecular force is actually an intramolecular force or vice versa!
1. Covalent bonding
For example, in diamond, silicon, quartz etc., the all atoms in the entire crystal are linked together by
covalent bonding. These solids are hard, brittle, and have high melting points. Covalent bonding holds
these particular particles tighter than ionic attraction. Thus, these crystals will have higher melting and
boiling points than ionic solids. (Example: all of the carbon atoms in diamond are covalently bonded to
the other carbon atoms in the crystal…thus diamond has one of the highest melting points of all known
substances.)
Ionic Bonding (strong ionic attraction)
Similar to covalent bonding above, ionic solids are composed of a large crystal of positive and negative
ions bonding to other ions in the crystal and thus also have high melting and boiling points! Greater
differences in electronegativity between metals and nonmetals do relate to properties of the solid. The
greater the difference in electronegativity, the more ionic the bond, the higher the attractive forces.
Examine the following list of energy (in kilojules/mol) needed to melt the ionic compound and see if
you can explain the observed values by way of ionic attraction: LiF 1036; LiI 737; KF 821; MgF2 2957
[Metallic bonding]
[Forces between atoms in metallic solids belong to another category. Valence electrons in metals are
delocalized. They are not restricted to certain atoms or bonds, rather they run freely in the entire solid,
providing good conductivity for heat and electric energy. The behavior of the electrons in metals give
special properties such as malleability, ductility and mechanical strength. We will not be analyzing
metals at this time!]
2. Dipole-Dipole Forces:
Strong dipole-dipole forces = Hydrogen ‘bonding’
Certain substances which have a hydrogen atom directly bonded to either a N, O, or F atom, (such as
H2O, HF, NH3), form a strong intermolecular force called hydrogen bonds, which affects the properties
(melting point, boiling point, solubility) of substances. Other compounds, containing OH and NH2
groups, also form hydrogen bonds. Molecules of many organic compounds such as alcohols, acids,
amines, and amino acids contain these groups, and thus hydrogen bonding plays an important role in
biological science.
Intermediate dipole-dipole forces
Substances whose molecules have a permanent dipole moment have higher melting point or boiling
points than those of similar molecular mass but whose molecules have no dipole moment.
3. London Dispersion Forces (LDF’s) or van der Waal's force- These forces always operate in any
substance containing electrons. The force arises from a temporary dipole that occurs when the electrons
surrounding the particle are unequally distributed. Typically, the interaction is weaker than the dipoledipole interaction. In general, the greater the number of electrons in the molecule, the stronger the van
der Waal's force of interaction. For example, the boiling points of inert gases increase as their atomic
masses increases due to stronger London dispersion interactions and the phases of the halogen are an
excellent example of the strength of these forces with increased number of electrons. (At room
temperature both F2 & Cl2are gases while Br2 is a liquid and I2 is a solid.)
The division into types of forces is for convenience and discussion. Of course multiple types of IMF’s can be
present simultaneously on a particle for many substances. Usually, intermolecular forces are best collated with
the phase of matter for the involved substance…if you know the phase then you should also have a good idea of
the strength of the IMF! Solids have strong IMF’s as do liquids while gases have weak IMF’s!
*All particles containing electrons will automatically have London Dispersion Forces available…but this
may not be the important IMF available so you must analyze the structure of the particle and determine if it
has other forces available to it as well!
Additionally, you must also consider ion dipole and induced dipole forces and how they affect the properties
and interactions of matter.
1. Ion-dipole interaction occurs when the oppositely charged portion of a polar molecule is attracted
strongly to a charged ion. These are strong attractive forces!
2. Ion-induced dipole – This occurs when an ion induces a stronger dipole on a nonpolar molecule. For
example, an iodide ion in solution can induce a strong dipole on a nonpolar iodine molecule
(I2)…stronger than the instantaneous dipole induced dipole force (LDF).
3. Dipole-induced dipole – A polar molecule is able to induce a weaker dipole on a nonpolar molecule as
well. For example, nonpolar O2 dissolved in polar water.
A summary of some of the interactions are illustrated in the following diagram:
In the above diagram, the ion-induced dipole will be very strong (clearly water strongly attracts the sodium ion) while the
rest of the interactions follow in order from L to R from strong to weaker interactions.
See if you can answer the following questions.
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What are dipoles?
What are dipole moments?
How do dipoles interact?
Why do molecules attract one another?
How do London dispersion forces come about?
What parameters cause an increase of the London dispersion forces?
What is a hydrogen bond?
What type of hydrogen bonds are strong?
What chemical groups are associated with hydrogen bonds?
Use your textbook to define the following terms and then use the concepts of intermolecular forces to analyze why certain
substances have differing values for these!
Melting pointBoiling pointVapor pressureSurface tensionViscosity“Confidence Building” Questions
1. Which of the following molecules have a permanent dipole (which are classified as polar)?
a. H2O b. CO2 c. CH4 d. N2 e. SO2 f. NH3
2. Order the above set of molecules from expected strongest to weakest IMF and justify your answer (note, you
should always think about the phase of these substances at room temperature for your initial order).
3. Order the above from lowest to highest melting point.
4. Will stronger intermolecular forces generate higher or lower melting and boiling points? Explain.
5. Which has the higher boiling point, Br2 or ICl? Explain.
6. An atom or molecule can be temporarily polarized by a nearby species. Polarization creates centers of charge
giving which of the following?
a. permanent dipole b. dipole c. hydrogen bonding d. induced dipole e. induced ions
7. How would you characterize the degree of polarization of a particle and what would it depend upon?
8. If only London dispersion forces are present on the following molecules, which should have a lower boiling
point? H2O or H2S Explain. (Note: at 25C water is a liquid while hydrogen sulfide is a gas.)
9. Which has a higher boiling point, I2 or Br2? Explain.
10. Ethanol (C2H5OH, molar mass 46) boils at 351 K, but water (H2O, molar mass 18) boils at higher temperature,
373 K. Diagram these molecules and explain the facts addressing IMF’s.
11. Ethanol C2H5OH and methyl ether CH3OCH3 have the same molar mass. Which has a higher boiling point?
Explain.