Extension of KM Theory to liquids

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Chapter 11
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• Extension of K-M Theory to liquids solids
• Still composed of particles but they are closer
  together in liquids solids
• Particles are still moving rate is dependent on
  T, but not proportional
• In liquid - particles move over, under around
  each other - held by attractions
• In solid - only vibrations - basically fixed in
  relative position
• Liquids still diffuse only slower.

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We have previously discussed what I call primary
chemical bonds (bonding between atoms within
molecules or ions). Now we will discuss what I
call secondary chemical bonds (attractions
between different molecules and/or ions). These
are also called intermolecular attractions or
more commonly Van der Waals attractions. There
are several types
1. Between 2 non-polar molecules - called London
Forces
2. Between two polar molecules - called
dipole-dipole attraction
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3. Hydrogen bonding
4. Between a polar and a non-polar molecule
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London Forces Also called Dispersion Forces.
Caused by attraction of outer electrons of one
molecule for the charged center of a nearby
molecule. This is extremely temporary as the
molecules are constantly moving and this
attraction is constantly changing. For the
instant of attraction, however, the 2 molecules
get distorted and thus each temporarily is
polarized ( a and end for each molecule
created by the distortion of the electron cloud).
Each of these will then induce distortions in
other nearby molecules, by reason of the
polarization.
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These are generally very weak - about 4 kJ /mole
or less Evidence for the existence of London
Forces Assume non-polar substance. Should
have no attractions (no charged ends). But we
know there must be attractions because all such
substances do have liquid and solid phases, where
intermolecular attractions are significant. Best
explanation is these are temporary, but always
present somewhere in the sample, called
polarizations.
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Two Major Effects of London Forces on M.P. and
B.P Review process of melting boiling,
vis-à-vis overcoming attractive forces Therefore
greater London Forces, higher M.P.s B.P.s,
all other factors being equal.
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1. All other factors being equal, bigger atoms
in a molecule, higher London forces, because
bigger atoms are easier to polarize. (Ex. F2
-188, Cl2-34.7, Br258.0, I2183
2. All other factors being equal, more atoms in
a molecule, larger London Forces, because
molecule is larger and therefore easier to
polarize. (Ex. CH4, C2H6, C3H8 and C4H10 ?-162,
-84.5, -42, 0)
In actuality both of these rules result from same
concept - all other factors being equal, the
larger the molecule, the greater the London
Forces. Be careful to compare similar compounds.
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Hydrogen Bond Essentially this is an extreme
case of dipole-dipole attraction. It occurs
whenever H is covalently bonded to N, O or F
only. Any time that occurs the effects of H-
bonding will be seen - Higher B.P.s and M.P.s
than expected
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Illustration - H2Te, H2Se, H2S, and H2O --gt
B.P.s are -2, -41.5, -60.7 and 100. For more
examples see Figure 11.7 on page 423.
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Hydrogen bonds are so strong that they can
sometimes be almost as strong as a primary
covalent bond. This is the case in H2O. In the
liquid state, H2O exists as a group of several
molecules traveling together and sticking
together because of the H-bond. When H2O
freezes, each group gets locked into a position
in the new crystal that forms, but since this is
an uneven group, this results in empty spaces in
the crystal (you can see these empty spaces if
you look closely at an ice cube. The end result
of all this is that ice is less dense than liquid
H2O ( one of the few substances where the solid
is less dense than the liquid). This is very
good, because ice floats. If not for this, in
winter cold climates, all life in lakes, rivers
and oceans would die, because they would freeze
from the bottom up, rather than from the top
down. They would totally freeze.
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True solids exist as crystals, which are
regularly shaped solid forms. There are several
common shapes. In all cases the particles making
up the crystal (atoms, ions or molecules) try to
get as close as possible to each other. There
are various ways to get this done, but we wont
worry about them.
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Amorphous solids dont form regularly shaped
crystals and thus have slightly different
properties. The most pronounced is that they
dont have exact melting points, but they
gradually get softer and softer until they appear
to be liquid. Glass is the most common amorphous
solid.
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Remember, we have already discussed phase changes
before. Phase changes are transformations from
one phase to another. During phase changes
energy is added or removed. All phase changes
are physical changes. We will discuss many of
these changes in more detail in the rest of this
chapter.
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First lets follow what happens to a solid when
we heat the substance. Solid, as T increases,
particles vibrate faster begin to move apart,
eventually T is reached where particles are
moving fast enough to overcome attractive forces
begin to move around each other (M.P.). (Now a
liquid), as T increases, move faster, density
decreases, (move farther apart, attraction
decreases), some particles attain K.E. great
enough to leave surface of liquid (go into gas
phase), eventually a T is reached when many
particles attain this K.E. volume of liquid
noticeable decreases, noticeable passage into gas
state (B.P). There is no change of T at M.P. or
B.P. (all molecules have sufficient energy, just
need the extra push) (Also at B.P molecules must
reach the surface)
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A diagram of this is called a heating curve or
for the reverse, a cooling curve. See Figure
11.38 on page 486.
Can reverse this. (Condensation freezing).
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Changes from liquid to vapor or vice versa
Molecules in the vapor phase are in a higher
energy state than molecules in the liquid phase.
In the vapor phase, as we learned earlier, the
molecules are basically free from influence from
the other molecules. while in the liquid phase
there are much stronger interactions between the
molecules. When a substance in the liquid phase
absorbs enough energy that molecules have energy
equal to the energy that those molecules have in
the vapor phase, then the liquid starts to
vaporize, transforming some of its molecules into
the vapor phase. Vaporization or Evaporation is
the process of going from a liquid to a vapor.
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Vapor Pressure Definition - Pressure of vapor in
equilibrium above a liquid (closed container).
Equilibrium is situation where 2 opposite
processes are occurring at the exact same rate
(thus in a closed container, eventually rate of
vaporization equals rate of condensation, which
is equilibrium). If container is not closed,
vapor cannot be trapped above the liquid,
therefore equilibrium is never reached and the
liquid eventually evaporates completely. Even in
that situation, we still refer to the Vapor
Pressure of the liquid at a specific temperature.
It is what the pressure of the vapor would be if
container was closed and equilibrium was
established.
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Molar Heat of Vaporization The amount of energy
needed to vaporize 1 mole of a substance. This
is a characteristic value for a substance, just
like density, melting point, etc. Directly
related to the strength of the intermolecular
forces in the liquid.
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Vapor pressure of any liquid increases as its
temperature increases. When the vapor pressure
equals the external pressure, the substance
boils, and the temperature at which this occurs
is called the boiling point. As a liquid is
heated bubbles start to form within the liquid.
This is simply the substance in its vapor state.
The pressure inside the bubble is essentially the
vapor pressure of the liquid at this temperature,
while the pressure on the bubble is essentially
the atmospheric pressure.
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If the atmospheric pressure is gt than the vapor
pressure, the bubble will burst. If the vapor
pressure is greater than the atmospheric
pressure, the bubble will also burst. If the two
are equal, the bubble will rise to the surface
and then burst (at the surface its temp can
increase beyond the temp of the liquid. Hence
the boiling point (BP) is defined as above.
Normal B.P. Temperature at which the V.P.
1.00 atm exactly.
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At any phase change, when the 2 phases are both
present, then they must be in equilibrium with
each other. Just as liquid ? vapor is
vaporization, and the reverse is condensation,
solid? liquid is melting and the reverse is
freezing. Just as the temperature for boiling
and condensation is characteristic (B.P), so too
is the temperature for freezing or melting
(Melting point freezing point)(M.P. F.P).
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Molar Heat of Fusion the energy needed to melt
one mole of a substance. These are generally
lower than Molar Heats of Vaporization (all
attractions dont have to be broken,
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Another equilibrium exists between the solid
phase and the vapor phase. Sublimation is the
process in which molecules go directly from the
solid phase to the vapor phase without first
melting. The reverse is called deposition.
Under ordinary conditions some substances only
sublime, never melting. Examples are dry ice
(CO2), and iodine. Other substances will sublime
under fairly mild conditions (water sublimes at
atmospheric pressure at any T below its F.P.).
There is also a Molar Heat of Sublimation. See
figure 11.39 on page 488 for a summary of all
these phase changes.
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A phase diagram summarizes all these equilibria
and phase changes along with the conditions for
each phase to exist or for each phase change.
Lets look at some of these in Figure 11.40 and
11.41 on page 489 of your text.
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One interesting equilibrium point that can be
seen on each diagram is called the Triple Point
for that substance. This is the only condition
at which all 3 phases exist together in
equilibrium.
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One other characteristic of substances that is
frequently illustrated in phase diagrams (but not
in these) is called the Critical Temperature Tc),
which is the highest temperature that a substance
can exist as a liquid under any pressure. The
minimum pressure that must be applied to a
substance at the critical temperature to liquefy
it, is called the Critical Pressure (Pc)