INTERMOLECULAR FORCES: HIGHLIGHTS, NOTES AND OBSERVATIONS
INTERMOLECULAR FORCES:
Surface Tension – The energy or work required to increase the surface of a
liquid. Liquids thus have a tendency to minimize their surface area creating an "elasticlike" skin on the surface. An increase in temperature decreases the surface tension.
You should be able to explain why. A wetting agent (soap or detergent) decreases the
surface tension of a liquid and allows it to spread more easily. The wetting agent (or
surfactant) is a molecule with a hydrophilic "head" and a hydrophobic "tail." The
hydrophilic ends interact with water and the hydrophobic ends only weakly interact with
water – and physically prevent them from hydrogen bonding – the surface tension is
reduced. One company's product claims that a 0.005% solution of their wetting agent
reduces the surface tension of water by 60%!
Cohesive Forces: Intermolecular forces between like molecules.
Adhesive Forces: Intermolecular forces between unlike molecules.
The meniscus in a cylinder is a result of the adhesion of the water to the glass. In
small diameter tubes this adhesion is responsible for capillary action; the water rises
within the tube. If mercury is substituted for water the meniscus is inverted because the
cohesive forces are far greater than the adhesive force between Hg and the glass.
Capillary action can be used to explain the movement of water in plant tissues but to
move water up at 200 foot tree requires a more significant force – We'll treat this in
Chapter 13.
Viscosity - The resistance to flow. Large cohesive forces create an "internal
friction" which reduces the rate of flow. Like surface tension, an increase in
temperature decreases viscosity.
Vapor pressure - The pressure of a vapor in dynamic equilibrium with its
liquid. It is dependent on two and only two things: the nature of the liquid (its polarity,
size, shape) and temperature. The conversion of liquid molecules to the gas phase is
called vaporization or evaporation (the reverse process is condensation). If the
vaporization is high at room temperature (weak intermolecular forces) we say the liquid
is volatile and we can infer that the enthalpy/heat of vaporization is low. As
temperature increases, vapor pressure increases.
A plot of temperature v. vapor pressure generates a vapor pressure curve.
The temperature at which the vapor pressure of a liquid equals that of standard
atmospheric pressure is the normal boiling point. At this temperature, vaporization
occurs anywhere in the liquid – not just the surface. A reduction of external pressure,
however allows a liquid to boil far below its normal boiling point. Conversely, high
confining pressures increase the b.p.
If water, or any other liquid for that matter is heated in a sealed container a
situation eventually arises where the meniscus disappears and one cannot distinguish
the liquid from the vapor phase. The point at which this occurs is called the Critical
Point. The temperature at this point is called the critical temperature, Tc, and the
pressure at this point is called the critical pressure, Pc. Practically, if you wish to
liquefy a gas at room temperature by increasing the pressure alone, the gas you're
working with must have a Tc that is above room temperature. (CO2 has a Tc of 31°C,
therefore by compressing the gas you may liquefy it. CH4 has a Tc of -82°C, therefore
no matter how hard you squeeze, you will be unable to liquefy it.)
A vapor pressure curve is a useful set of data but scientists like straight lines. It
so happens that if you plot the natural logarithm of vapor pressures against the
reciprocal Kelvin temperature you get a linear relationship. The most useful form of this
equation is called the Clausius - Clapeyron equation and is as follows:
A single accurate measurement of a vapor pressure at a given temperature, and the ∆H
P2  Hvap  1 1 

ln 
  

P1
R T1 T2 
for vaporization (Look up the enthalpy of formation for the vapor and subtract the
enthalpy of formation for the liquid) and 760 mm for P2, allows one to calculate the
boiling point of the liquid!
Intermolecular forces in solids define properties such as melting (fusion) point,
(freezing point being the reverse term), sublimation and deposition. A pure
substance when heated or cooled will show no temperature change during its phase
change, only before and after. This can be explained because all the energy involved is
related to the position of molecules and breaking or forming intermolecular bonds (P.E.
not K.E.)!!
A Phase Diagram shows what state of matter (g, l, and/or s) is found at a given
pressure and temperature. Study the diagrams of both carbon dioxide and water. The
triple point of a substance is the unique pressure and temperature where all three
states of matter can coexist in a dynamic equilibrium - the rates of all phase changes
are equal so no net change.
Van Der Waals Forces
1) Dispersion or London Force - The electron cloud may at times be asymmetrical
and thus create an instantaneous, momentary dipole (the atom/molecule briefly
achieves a partial negative and partial positive pole). If a neighboring atom is close
enough and moving slow enough (low temp.) its electron cloud may shift in response
(an induced dipole) and the two may attract. This intermolecular attraction is known
as the Dispersion Force. It is the force commonly associated with nonpolar molecules.
The ease with which a nonpolar molecule can achieve an instantaneous dipole is called
polarizability. The strength of this force depends on molecular shape and size - the
bigger the molecule, the more electrons and the greater the polarizability. (ex. Long
chains of nonpolar hydrocarbons have higher b.p.s, viscosity, surface tension than short
chains). The role of shape is nicely revealed when measurements of volatility and
boiling point are taken on isomers. They of course have identical molecular weights but
their atom arrangement is altered - shape defines function. I have a nice demonstration
of this phenomenon.
PLEASE NOTE: All molecules will have random electron asymmetries – it's just that it is
the only force that can account for attraction between nonpolar molecules.
Dipole-Dipole - If a molecule is a polar molecule, as determined by electronegativity
differences among the atoms and the shape of the molecule, than it may line up with
other molecules with some degree of order. If you could, you would see a partial
positive of one molecule near a partial negative of another molecule and so on and so
on.... Assuming that you are comparing molecules with similar molecular weights, the
dipole-dipole intermolecular force is stronger than the Dispersion Force.
Hydrogen Bonding - A mega version of the Dipole- Dipole Force. It is strictly defined
for the intermolecular bond between a hydrogen that is polar covalently bonded to a N,
O, or F in one molecule to a highly electronegative atom on a different molecule
(Always an N, O or F -- but Cl and S are sometimes tolerated and result in "weak"
hydrogen bonding - lovely). Hydrogen bonds are reasonably strong intermolecular
forces. If the molecule is large enough and shaped appropriately it is possible for
intramolecular hydrogen bonding. Lastly, sometimes hydrogen bonding gets overblown
with respect to its strength. To be absolutely sure, covalent bonds are immensely
stronger. Remember that when you heat water to 100°C, the hydrogen bonds are
broken but you still have water. You almost always break intermolecular bonds before
breaking intramolecular bonds.
Network Covalent Structures: A case where intramolecular and intermolecular forces
are one and the same. Diamond is pure carbon in an extensive tetrahedral crystal
lattice. Each carbon is bonded to four other carbons (sp3 hybridization) in a ever
extending 3D structure - a macro molecule if you will.
Graphite has carbon sp2 hybridized with an additional p orbital perpendicular to the
plane. The alternating double bonds create a delocalized electron cloud from the p
orbital overlap and explains the excellent conductivity of graphite along the planes of
carbon. The lubricating properties of graphite are also explained by the relatively weak
bonds (dispersion forces) between the layers of graphite - They slide off each other with
great ease.
Interionic Forces - The attractive force (similar to the lattice energy) between a pair of
oppositely charged ions increases with increased charge on the ions and with
decreased ionic sizes. To dissolve an ionic compound, the energy of hydration must be
similar to the lattice energy. Generally, high charged ions with small radii do not
dissolve well in water.
Crystal Structures are of great interest but I will limit the treatment of this section. To
adequately consume this information would require a great deal of time, patience and a
$100 dollar crystal model kit. We at best have the middle requirement. Work problems
and study the chapter well.
Table 11-7 (page 470) is a nice organization of information on crystalline solids.