Chapter 15 Liquids and Solids
In chapter 13 we studied gases. Gases are the simplest form of matter to study
because there are no interaction between the molecules to make things messy. In this
chapter we will start to study liquids and solids, the condensed forms of matter. In
these condensed state molecules literally begin to touch each other and interact with
each other, so trying to understand what is going on at the molecular level becomes
much more interesting.
Here we will start by studying some of the general properties of liquids and solids, and
the transitions between these states, then toward the end of the chapter we will look at
what happens at the molecular in these condensed state.
15-1 Molecules in liquids and solids
use table 15.1?
Gases - Molecules not in contact distance between molecules 100x larger than
molecule, lots of space between molecules so can translate, rotate and vibrate,
basically no interaction between molecules. That is why compressible and can
move to fill container
Liquids - Molecules are in contact (distance between molecules ~ same as
molecul itself), but position is not fixed, so can translate slowly (with lots of
collisions) rotate (also hindered by collisions) and vibrate. Not very
compressible, but also not rigid
Solids - Molecules in contact and position fixed, so do not translate and usually
don’t rotate. Just sit there and vibrate. Not compressible, and the solid is rigid
15-2 Heating curves
Let’s plot the temperature of water as we heat it as a constant rate, starting with
the material at -10oC. What do you see?
Sketch something like figure 15.2 on board
This kind of a plot is called a Heating Curve
Notice in the solid, liquid and gas states we have a straight line with a slope.
What would this represent? (The heat capacity of the s, l, or g)
But then we have flat lines connecting these regions at 0 and 100
What does this represent?
The energy needed to transform a solid into a liquid, the molar enthalpy of
fusion, or the heat needed to change a liquid into a gas, the molar entropy of
vaporization.
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Key Definitions:
Molar enthalpy of fusion ÄH fus = the heat energy required to change 1 mole of
solid into a liquid at the melting point of that material
Molar enthalpy of vaporization ÄH vap = the heat energy required to change 1
mole of liquid into a gas at the boiling point of that material
15-3 Fusion and vaporization
The enthalpy of fusion for water is:
H2O(s)6H2O(l)
ÄH fus = 6.01 kJ/mol
In the general case we would write
X(s) 6X(l) ÄH fus
The heat needed to melts a material is then
q = nÄH fus
What is going on here. There are intermolecular interactions that hold each
molecule or atom in its place in the solid. The ÄH fusion represents the (+)
energy added to the material to break just enough of these interactions that the
molecules can begin to move around just enough to let the material as a whole
become a liquid. We do not add enough energy to remove all intermolecular
interactions, otherwise we would go directly to the ga phase
The enthalpy of vaporization for water is:
H2O(l)6H2O(g)
ÄH vap = 40.65 kJ/mol
In the general case we would write
X(l) 6X(g) ÄH vap
The heat needed to melts a material is then
q = nÄH vap
Here we are adding more energy to remove the last of the intermolecular
interactions that hold the molecules together. The ÄH vap represents the energy
needed to remove all these interactions so the molecules no longer have any
interaction with each other, and the material becomes a gas
It is possible to go directly from a solid to a gas without going through a liquid
phase. This process is called sublimation
Again we can define a ÄH sub which is the energy required to change one mole of
solid directly to a gas
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If you have ever used dry ice (solid CO2) you will have observed that the solid
disappears without a liquid phase. You may also have observed ice cubes in
your refrigerator getting smaller with time. The are both examples of sublimation
Sample calculations
How much energy is required to melt 50g of ice into water at 0oC?
q=n ÄH fus
n=50/18.02 = 2.77 mole
q = 2.77 mol x 6.01 kJ/mol = 16.68 kJ
How much energy is required to take 50g of ice at -10oC and change it into water
at 60oC?
This is a little more interesting. You actually have three different steps
Step 1: ice @ -10 6 ice @ 0 use the heat capacity of H2O (s)
Step 2: ice@06water @ 0 Use the ÄH fusion
Step 3: water @ 0 6 water @ 60 use the heat capacity of H2O (l)
Since we are using 50 g of water in all three steps, let’s start with
n=50/18.02 = 2.77 mole
Step 1:
q = n CP ÄT
ÄT = 0-(-10) = 10
q1 = 2.77 mol x 37.7 J/K@mole x 10 K
=1044 J = 1.044 kJ
Step 2:
q2 =n ÄH fus
n=50/18.02 = 2.77 mole
q2 = 2.77 mol x 6.01 kJ/mol = 16.68 kJ
Step 3:
q3 = n CP ÄT
ÄT = 60-(0) = 60
q3 = 2.77 mol x 75.3 J/K@mole x 60 K
=12500 J = 12.5 kJ
qtotal = q1 + q2 + q3
= 1.04 + 16.68 + 12.5 = 30.2 kJ
How much energy must be removed from 50 g of water at 60o C to change it to
ice at -10oC?
The reverse of the above process, so -30.2 kJ!
Think about what you would have to do to take this same 50 of ices at -10 to
change it into steam at 400oC!
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15-4 Intermolecular forces
The value of ÄH vap or ÄH sub reflects how strongly the molecules are held
together in the liquid or solid. The larger the ÄH the stronger the forces
So what are the forces holding molecules together? In this section will talk about
three different forces, Charge-charge electrostatic forces, Dipole-dipole
interactions and hydrogen bonds, and London (of van der Waals) forces
I will tackle them in order from strongest to weakest.
The strongest force is the charge-charge electrostatic interaction that we
talked about in chapter 6 when we discussed ionic compounds. The electrical
forces between ions is strong an relatively long range, and this is reflected in the
high ÄH vap of ionic compounds ($100 kJ) and the high melting points and boiling
points of ionic compounds when compared to other types of compounds
The next strongest force is called the dipole-dipole interaction or dipoledipole attraction. This is a force observed in molecules that contain a dipole
(and this is why I tried to make sure you could identify this kind of molecule in
chapter 7)
While a molecule with a dipole had no net charge, so there is no charge-charge
interaction, one end of the molecule has a partial positive charge compared to
the other end (Make diagram of H2O on board) The partial charge + charge on
one molecule is attracted to the slightly negative charge on the other end of a
second molecule, and this is what makes the intermolecular attraction between
the molecule.
ÄH vap of polar molecules are typically ~ 10-20 kJ, making this interaction about
10x weaker than the charge-charge interaction. They are also 10-20x weaker
than a typical bond energy, emphasizing the fact that intramolecular forces
(between molecules) are much weaker than intermolecular forces (forces within
the molecule)
When you combine the small hydrogen atom with a strongly electronegative
atom like N, O or F, you get a particularly strong dipole-dipole interaction, to
which we give the special name hydrogen bond. Hydrogen bonds are 2-3 x
stronger that regular dipole-dipole forces, and any compound that has H
attached to N, O, or F will have this extra strong attractive force.
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It is this hydrogen bond that give water many of its special properties, like high
melting point and boiling point. It is also responsible for the relatively open
structure of ice, that breaks down on heating so liquid water can be denser than
solid water, just the opposite of what is seen in most solids that are more dens
than their corresponding liquids
The last intermolecular force that we observe is call the London force, and it
occurs in any molecule that has electrons (is all molecules) This force is
sometimes called an induced dipole-induced dipole interaction.
Let’s use a couple of He atoms to illustrate how this force works
With only 2 electrons in an He atom, and both of those electrons are wizzing
around in an orbital, through random chance the two electrons will be on one
side of the atom at one time. The, for a brief instant before the electrons move
on (at .01 the speed of light) , the molecule has a dipole. This dipole will
influence the electrons in the neighboring He atom, so they will shift toward the +
temporary dipole of the first atom, so you have induced a dipole in the second
atom. You now have a very brief dipole-dipole interaction that, very temporarily
make the two atoms attracted to each other
(Make sketch like figure 15.8 & 15.9 on board)
Because of there very fleeting nature, London interactions are very weak. Also,
since one molecule has to be almost touching the second molecule to be able to
induce the dipole, the London forces are also very short range.
Since you need electrons for a London force, the more electrons you have, the
stronger the London force. So in general the larger a molecule or atom, the
stronger the London force. Table 15.4 See how He and Ne have ÄH vap 10 or
100x smaller than dipole interaction (~10-20) but something like I2 with good size
and lots of electrons is about equivalent to a dipole interaction.
Practice problems & clicker questions
Rank the following atoms or molecules in order from largest to smallest
intermolecular forces, and predict the relative MP and BP of these chemicals
(Choose several)
NaCl, MgS, LiF, CH2O CH2OH, H2O,Cl2, Xe, CH4 CH3-CH2-CH3, CH3-O-CH3
CF4
15-5 Properties of liquids
Several properties of liquids that depend on intermolecular forces
Key definition:
Viscosity - a measure of the resistance to flow in a liquid
Can be measured several ways, one is to simply see how fast a given volume of
liquid flows through a tube. Another is to drop on object in a liquid as see how
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fast it get to the bottom. Everyday example: 40w vs 5 w motor oil. Tied directly to
intermolecular forces, the higher the forces, the higher the viscosity. Generally
observed that at T8 v9. Why? As T8 molecules move around more, as move
around more the intermolecular interactions 9 so v9
Key Definition:
Surface tension - a measure of the energy needed to increase the surface area
of a liquid
What does that mean? Its pretty hard to explain. Lets start with our
intermolecular forces Sketch like 15.13 on board When a molecule is in the
interior of a liquid, it has intermolecular forces all around it pulling it in all
directions. However when it is at the surface there is nothing to offset the forces
trying to drag the molecule back from the surface into the interior, so there is a
net force taking all the molecules at the surface back into the bulk liquid.
This force trying to make all the molecules at the surface move back into the
liquid has the net effect of trying to make the surface area a minimum. You see
this with water drops and Hg drops
You can either think of surface tension as a measure of this force, trying to drag
the molecules back into the bulk solvent, or you can think of it as a measure of
the energy needed to increase the surface area as you distort it by placing
something on it (Figure 15.12)
Again the higher the intermolecular forces, the higher the surface tension
You can change the surface tension of a liquid by adding solutes that tend to
aggregate on the surface and lower the surface tension. These king of materials
are called surfactants
Detergents are the most common example of a surfactant I can think of. How do
they work? If you have a greasy stain on your clothes and wash them with pure
water, the water and the grease do not interact because they have different
intermolecular forces holding them together (See if you can figure out which
ones!). In fact if you pour water on the grease spot you may even see the water
‘bead up’ as it tries to minimize its contact with the grease. When you add
detergent, it lowers the surface tension of the water. One the water molecules
are not so strongly attracted to each other , they can penetrate into the grease
more easily and begin to dissolve the grease instead of avoiding it.
Surface tension occurs at an interface because the intermolecular forces on the
liquid side are not compensated for at the surface of the liquid. But what
happens at a liquid-solid interface, when the solid can do some of the same
interactions as the liquid?
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Key Definition
Capillary action arises from the attraction of a liquid to a solid capillary tube
Here we now have to differentiate between the cohesive interactions holding the
liquid together and the adhesive forces that may arise between the liquid and
the solid. Both of these forces involve the intermolecular forces we are studying
here.
You have already observed capillary action 1000's of times in every day life.
Why do paper towels absorb water? How can water rise in a plant without a
pump?
The interplay between adhesive and cohesive forces at an interface is very
visible in the meniscus of a liquid in a buret. Why does the meniscus curve
down when the buret is filled with water? Why does it curve up when it is filled
Hg?
Let’s take our discussion of properties of liquids back to the forces that hold
liquids together. What forces should we see in liquids?
(Not charge-charge too strong, observed in solids)
so that leaves dipole-dipole and London
Why do some liquids mix and some don’t? Why do some materials dissolve in
water, but not in gasoline? Polarity like-dissolves like. Again that is why I had
you always determining the polarity of things back in chapter 6
15-6 Vapor Pressure
Now let’s talk about what happens to a liquid as it vaporizes
Sketch like 15.19 on board
To become a vapor molecule, a molecule at the surface of a liquid must have
enough kinetic energy to break free of the surface and go into the gas phase.
As you learned in the chapter 13, at any temperature the molecules of a
substance have a range of kinetic energies. So at a particular temperature a
certain % of molecules will have enough KE to go into the vapor phase.
Initially, with no molecules in the vapor phase the partial pressure of the liquid in
the gas rises quickly (figure 15.20 on board) but as more an more molecules get
into the vapor phase, some of those molecules randomly hit the surface of the
liquid and return to the liquid phase. At some point the rate of molecules
entering the gas phase = rate of molecules returning to liquid phase
(condensing) and you reach what is called the equilibrium vapor pressure
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Key definition:
The equilibrium vapor pressure is the partial pressure of a gas that is in
equilibrium with its liquid at a specific temperature.
Another thing you learned in chapter 13 was that at the temperature increases,
the average kinetic energy of the molecules in a sample increases, so the
number of molecules entering the vapor phase will increase. Thus the
equilibrium vapor pressure is a function of temperature, and as T8, VP8
Figure like 15.23 on board
The plot of VP on the Y and T on the X axis is called a vapor pressure curve.
As you can see different liquids have difference vapor pressure curves, and the
curve seems to rise exponentially. In chapter 23 we will come up with a
mathematical model for this curve, but for now we won’t go into the math in
depth.
Liquids or solids that have a high vapor pressure are called volatile and have
relatively low intermolecular forces, ones with low vapor pressure are called nonvolatile and have higher intermolecular forces
On figure 15.23, what is the significance of the line at 760 torr?
760 torr is the atmospheric pressure. When the vapor pressure of a liquid is
above this value the gas can readily form bubbles and leave the liquid because
its pressure is at or above the air around the liquid. When the VP is below this,
the pressure of the air prevents bubble from forming and the liquid doesn’t form
bubbles
This is the way we define the boiling point of a liquid
Key Definition:
The normal boiling point of a liquid is the temperature at which the vapor
pressure of a liquid is equal to 1 atm.
While the normal boiling point of water is 100oC, what is the BP here in
Spearfish? The atmospheric pressure in Spearfish is generally around .88 atm,
so when you boil water the vapor pressure of the water matches the external
pressure at a lower temperature so the water begins to boil earlier, about 96oC
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Sample calculation using table 15.7
I am going to put .1000 g of water in a 1L flask, and seal it so no water can
escape. If I heat the flask to 50oC will any liquid be left in the container?
PV=nRT
P = nRT/V
T = 273 + 50 = 323K
n = 1g x 1 mole = .0556 mole
18 g
P =( .0556 mole x .08206 L@atm/K@mol x 323K )/1L = .148 atm
The vapor pressure of water at 50oC is .122 atm
Since our theoretical pressure is higher than it should be, liquid water will
condense until the VP in the container = .148
How much will condense?
If the VP or water at 50oC is .122, then the most water that can be in 1 liter of
gas is:
PV=nRT
n=PV/RT
= (.122 atm x 1L)/(.08206 L@atm/K@mol x 323K)
=.0046 mole
.0046 moles x 18 g/mole = .0826 grams of water in the vapor phase
.1 - .0826 = .018 g liquid water remains
15-7 Relative Humidity
Now that you know what vapor pressure is, you can understand what relative
humidity and dew point are. However we don’t worry too much about these
parameters relative in the chem lab, so I think I will skip this section. Feel free to
read it if you are interested.
15-8 Phase Diagrams
You just examined the vapor pressure curve, you can do a similar curve to plot
when a solid and a gas are in equilibrium (Sublimation curve) or solid and liquid
are in equilibrium (melting point curve)
Rather than having 3 separate curves, we can put them into a single diagram
called a phase diagram Figure 15.24 or on board. Point out the three different
curves, work your way from solid to liquid to gas at 1 ATM. Also point out the
breaks on both the X and Y axis
Let’s look at some of the more interesting pieces of this diagram
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The Critical Point
Let’s start with the upper right hand end of the vapor pressure curve. Rather
than going up and to the right forever, the line actually has a well defined
stopping point, call the critical point
What is happening here is that you reach a temperature above which the gas
cannot be changed into a liquid and the liquid and the gas become
indistinguishable. The state is called a supercritical fluid (More on this in a bit)
Key definitions:
Critical temperature - the temperature above which a gas cannot be compressed
into a liquid, bot matter how high you make the pressure.
Critical Pressure - The pressure required to liquify a gas at its critical temperature
Melting point curve
At a given pressure, if you are colder than the line you have a solid if you are
warmer than the line, you are a liquid. If you are right on the line you have an
equilibrium between solid and liquid so you solid is trying to melt and you have
slush
Key Definition:
The temperature at which solid and liquids are in equilibrium and the total
pressure is 1 atm is called the normal melting point or normal freezing point.
It is hard to see in this figure, but water’s melting point curve is a little unusual in
that it has a negative slope. We discussed earlier about how solid water is less
dense than liquid water. Thus when you are near this melting curve in the solid
form, and you apply pressure to compress water, it shifts to it’s denser form, the
liquid. Not many materials do this
Sublimation curve
The sublimation curve is where the solid can go directly to gas. This is how
freeze-drying works. If you take anything that contains water and cool it below
273.16K (~0o C) , and reduce the pressure to < .006 atm, the solid can sublime
directly to a gas, and you can remove the remove the water!
Triple point
Notice there is one point on this diagram where all three phases can coexist at
one temperature and pressure. This point is called the triple point
Key Definition:
The triple point is the point at which three phases can coexist at one temperature
and pressure
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For contrast let’s look at the phase diagram for CO2 Figure 15.25
This has a more normal + slope (Again hard to see in this figure)
at 1 atm transition from solid to gas occurs at -78, so what is the temp of
something you just pulled out of dry ice? (-78)
Triple point 5.1 atm, -56.6oC, critical point at 72.8 atm 31oC
Many fire extinguishers contain CO2. If they contain liquid CO2 what would the
internal pressure have to be?(close to 70 atm)
When release liquid goes to gas, cools off fast. The white fog is actually water
NOT CO2! CO2 is also heavier than air so settles (Should CO2 be heavier than
air? How can tell (Air is mostly N2; N2 has MW of 28, CO2 has a MW of
12+16+16= 44)
Another practical application. We just talked about the critical point. When you
get a fluid up into this region between liquid and gas you have what is called a
super-critical fluid. This may sound like some exotic came talk, but it has very
day-to-day applications. How do you decaffeinate coffee? One way is to extract
with thing like chloroform, or CH2Cl2 (dichloromethane). Not very appetizing.
Another way is with supercritical CO2! In fact this is how the bulk of
decaffeinated coffee is produced. Put it in the CO2 at high pressure and RT and
the caffeine leaches out of the coffee bean. Drop the pressure and , boil off the
CO2 and you bean is decaffeinated! What do you do with the caffeine you have
left? You sell it to the people making soft-drinks!
Caffeine
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Practice problems and/or clicker questions
Using the phase diagram for C, pick out the triple point, critical point etc.
15-9 Crystal Structures
We will now turn our attention to solids.
Solids are usually broken down into two groups, crystalline and noncrystalline.
We will start our discussion with crystalline solids because they have a regular,
repeating structure that is easy to understand. We will end the chapter with brief
discussion of a couple of forms of noncrystalline solids, amorphous solids and
liquids crystals
Crystalline solids
The distinguishing feature of a crystalline solid is that it has a regular repeating
structure, a structure we call the crystal lattice.
This regular repeating structure allows us to use diffraction to probe and identity
the underlying repeating structure.
This text does not go into much depth on how diffraction works, so let me add a
little extra material
X-ray Analysis of Solids
How do we know that the crystals have the structure they do?
Found by X-ray diffraction, that is shine X-ray light at the crystal and see how it
bounces off. Look at figure 10.10 Zumdahl`.
The whole idea is based on constructive and destructive interference. Let’s sit at
angle theta with respect to our crystal plane. Notice in case 1 the wave forms
going to the second atom travel exactly two wavelength more. What happens if
the distance between the atoms is different or the angle is different? Wave out
of phase, and cancel out.
Without getting into the math, which is elementary geometry if you want to think
about it we have the relationship
Key Equation:
në = 2dsin è
This is called the Bragg equation.
Essentially what we do is to shine X-rays at a crystal then look for all the angle
places where X-rays appear around it. Can reconstruct distances between
atoms form this.
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Practice Problem:
If you use an X-ray with a 1.54D wavelength to examine a sample of silicon and
observe an angle of 14.22o for the first order reflection (n=1), what is the distance
between layers in the silicon crystal?
1(1.54D) =2d sin(14.22o)
1.54 = 2d(.2456)
1.54D/(2X.2456) =d; d=3.13D
The Crystal Lattice
Figure 15.28
The smallest subunit of a crystal lattice that can be repeated over and over to
make the entire crystal is called the unit cell.
There are a wide variety of 3-D units cells, but here we will only discuss the cubic
unit cell, because that is the one that is sen in most metallic elements
Figure 15.29
There are three different kinds of cubic unit cells, the simple cubic cell, which
has an atom at each cornet of the cube, a body-centered cubic unit cell, that
has an additional atom in the very center of the cube, and a face-centered
cubic unit cell that has the atom at each corner and an atom in the center of
each face of the cell.
Let’s look a little more closely at each. Notice that is the simple cubic unit cell,
each atom is at a corner of the cell, and only 1/8 of that atom is actually in the
cell With 8 corners and 1/8 of an atom in the cell. Taht make a total of 1 atoms
/cell
Interestingly, of all the metals, only polonium uses this type of unit cell
With the body-centered unit cell, we have 8 x 1/8 corner atoms + the one in the
middle, so we have 2 atoms/unit cell. Metals that use this type of unit cell
include Ba, Cs, K, LI, Mo, Ta, U, & V
In the face-centered cell we have 8 x 1/8 + 6x1/2 for a total of 4 atoms/unit cell.
Metals that use this lattice include Ag, Au, Cu, Ni, Pb,Sr, Pt, & the noble gases.
Calculations
There is an equation I want you to be able to use, and it is fairly simple to derive
You know that density = mass /volume
Now lets substitute in the molar mass and the molar volume
density = M/Vmol ; M = molar mass, Vmol = molar volume
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Let’s get Vmol on one side of the equation
Vmol = M/density
The volume of one atom = molar volume/Na
Volume of 1 atom = Vmol/NA = M/(density × NA)
And if the unit cell has n atoms, the volume of atoms in the unit cell is
(n M) /(Density ×NA)
Key Equation:
Using this equation you can go in lots of different directions.
Example problem 1
Cu crystallizes in a face centered cubic lattice and has a density of 8.96 g/cm3
Calculate the volume of the Cu unit cell.
A face centered cubic lattice had 4 atoms in the unit cell so n=4
Vatoms in unit cell = (4atoms/cell x63.55g/mol)/(8.96 g/cm3 x 6.022x1023 atoms/mole)
= 4.71 x1023 cm3/cell
Getting to the actual volume of the atom is a bit trickier, and involves
remembering some basic geometry. I don’t think I will do this to you on a test so
let’s not go that far. You can see how to do it in your text if you are really
interested.
Example 2
Cerium crystallizes in a face-centered cubic lattice. Given that the density of
Cerium is 6.770 g/cm3, and the length of an edge of the units cell is 516.10 pm
determine Avogadro’s number
Volume of unit cell = (n M) /(Density ×NA)
Rearranging this equation to find NA
NA = (nM)/(Density x Volume of unit cell)
if the length of an edge of the unit cell is 516.10 pm the volume of the unit cell is
(516.10 pm)3
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Let’s convert that to cm
1 pm x 1x10-12 m x 1cm
= 1x10-10 cm
-2
1 pm
1x10
516.10 pm x 1x10-10 cm = 5.1610x10-8 cm
pm
Volume of unit cell = (5.1610x10-8 cm)3 = 1.37458 x10-22 cm3
n, the number of atoms in the face centered unit cell is 4
so
NA = (nM)/(Density x Volume of unit cell)
=(4x140.116)/(6.770 g/cm3 x 1.37458 x10-22 cm3)
=6.0227x1023 Sounds familiar!
15-10 Crystal Forces
Crystal structures determined by:
size of atom, ion, or molecule making up lattice
type of forces used to hold crystal together
Ionic Crystals
Strong Charge-charge interactions holding lattice together
have ions at lattice points
High MP, low BP
Relative size of ions determined lattice used
Figure 15.30
(Cl>Na) NaCl Chloride is face centered cubic
(Cl<Cs) CsCl Chloride is simple cubic
Molecular crystals
entire molecules at lattice points
mix of Hbonding/ dipole-dipole and London Forces holding together
Lower MP and Higher VP than Ionic crystals
Covalent Network crystals
Covalent bonds holding together
atoms at lattice points
Can be all one type of atom (Cdiamond or Cgraphite)
Or can be different atoms (SiO2)
Metallic Crystaline (See next section)
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Amorphous solids - not crystalline
no precise crystal structure
do not melt at a single sharp temperature
Usually soften over a wide range of T
Can think of like liquids that froze to fact to make proper oriented crystals
Rubber, plastic, glass
15-11 Electrons in Metals
Metallic Crystalline solids
Bonding between atoms in metals is something novel
Figure like 15.35 on board
When made a molecular orbital from 2 atomic orbital in nonmetals
had big separation between bonding and nonbonding orbitals
In metal crystal have many more metals
have a wide band of many bonding and non-bonding orbitals
Figure like 15.35,15.36, and 15.37
Nonbonding orbitals are the conduction band,
Bonding orbital make the valence band
Lots of orbital and very little gap in between can move electrons around
easily
Call metallic covalent bond or delocalized covalent bond
Allow metals to move around, conduct heat and electricity
Also explains why nonmetals don’t conduct
and starts to explain semiconductors (with small bad gap)
15-12 Liquid Crystals
So we have studied crystals, what is a liquid crystal?
Figure 15.39
Liquid crystal has properties of both liquid and crystal.
Illustrate with rodlike, hydrophobic molecules
Like to use london force to align side by side
can get Smectic phase where have layers
or Nematic phase where layers disappear
so have phase transition between the two
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15-13 Colloids
In a solution, every individual atom or molecule of the solute is completely
dispersed in the solvent.
In a suspension you have visible (you may have to use a microscope) particles
dispersed in a solvent and, given enough time, the particles will settle to the
bottom.
In between we have colloids or colloid dispersions
particles are too small to bee seen, even with a microscope
But are not individual atoms or molecules
will not settle to the bottom
Not an ‘official’ state of matter like s,l,g but still has special properties
particle size ~ 1-1000 nm (~ size of starch, proteins, DNA)
With such small size, particles have very large surface area and if surface is
attracted to solvent, that is what will keep particle suspended indefinitely
A little bit of nomenclature
Dispersed phase - the particles that are dispersed
Dispersion medium - the medium in which they are dispersed
Just like solutions are not all aqueous, but can include gases and solids,
dispersions can have several different phases.
See table 15.9 for the proper name of each kind of dispersion
One kind of dispersion that plays an important role in everyday life is an
emulsion. As you can see an emulsion is a dispersion of a liquid in a liquid.
Mayonnaise is an emulsion of the fats and oils of egg yolk in water. Emulsions
of oils and water are usually not stable, unless a detergent is present. On the
molecular level soaps and detergents are large hydrophobic molecules with a
charged or polar end. As shown in figure 15.43 the oil water emuslion is made
up of thousands of tiny micelles where the detergent essentially forms an
interface between the hydrophobic oil and the hydrophillic water. It is this
interface that keeps the emulsion stable. You will see this same kind of structure
over and over again in biochemistry in the structure of proteins, membranes,
cells, and even DNA itself.
How do you tell experimentally if you have a colloidal solution? The dispersed
particles will scatter light much differently that a true solution. The scattering of
light by a dispersion is called the Tyndall effect. Demo with laser and starch