There are three physical states that we need to be aware of: solids, liquids, and gases. Depending on the
circumstances, many substances can exist as solids, liquids, or gases. Water vapor is a gas and is a
constituent of the air that we breathe, we drink or swim in liquid water, and we skate or chill our drinks
with solid water, known more commonly as ice. All around us we can pick out many common examples
of substances that exist as solids, liquids, or gases. Gases have many unique and interesting properties
that will be examined later. Now the focus will be on the solid and liquid states of matter.
A phase describes the physical state of the substance at some temperature and pressure. Is water a
liquid or a solid at room temperature? What happens to frozen water at room temperature? How many
phases are present when you put ice cubes into a glass of water? How many phases are present when
you put ice into a glass of Pepsi? Liquids and solids are known as condensed phases as their particles
are close together.
Liquid water
There is energy between particles in the liquid and solid phase as the particles come near one another. In
the gas state, the particles are too far apart to notice the other molecules present, but in liquids and
solids, there is an interaction between the molecules or ions. This is governed by intermolecular forces
which occurs between molecules in the liquid or solid state. These forces give rise to the properties of a
particular substance and dictates the phase that the substance will be in. WHY is O2 a gas a room
temperature? Why is water a liquid at room temperature? And what causes a substance to change its
phase?
There are two types of forces at work when examining a chemical substance:
Intramolecular forces: the bonds that exist between atoms within a molecule. We talked about these
types of attractions for ionic species which result from the transfer of electrons, with covalent species
which balance the attractive forces of the electrons with the nuclei and the repulsive forces of the electron
clouds and the two nuclei with one another, and finally, metallic bonding which incorporated the idea of
a sea of mobile electrons which surrounded positive metal cations. These bonds influence the chemical
properties of the substance.
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Intermolecular forces: the forces that exist between molecules (which are atoms/ions bonded together)
which influence the physical properties of the substance.
For example: consider solid water and liquid water. Their chemical behaviors will be identical as the
intramolecular forces are the same; both are made up of H-O-H molecules in the bent molecular
geometry (tetrahedral VSEPR!!) However, their physical properties will be different based on the
intermolecular forces that the individual H-O-H molecules can undergo in their various physical states.
Whether or not a substance exists in the liquid, solid, or gas phase depends on the balance of the
potential energy of the intermolecular attractions, which tends to draw the molecules together, with the
kinetic energy of the molecules, which tends to drive them apart. The potential energy depends on the
charges of the molecules/ions and the distance between them while kinetic energy depends on the
temperature.
This balance of the potential energy of the attraction and the kinetic energy of dispersion directly affects
the properties that each of the three phase states displays:
Gases: the energy of attraction is very low compared to the kinetic energy of motion. This keeps the gas
molecules far apart from one another (minimizing attraction even further!) This large distance between
molecules allows gases to take up the entire space of the container. Good thing too – because if all the air
only congregated in the front of the room the people in the back of the room would be in big trouble!
Gases are highly compressible and they flow and diffuse through one another very easily. Remember,
think of air. It is made up of O2, N2 Ar, H2O(g), and CO2. They better flow around and diffuse easily or
you exhale and the only air in front of you would be CO2!!
Liquids: the energy of attraction is much greater than in gases and the particles remain in virtual contact
with one another. However, they still have kinetic energy and this kinetic energy allows the particles to
“tumble” around one another. The molecules are held in a discrete location by the boundaries of
attraction. They cannot go anywhere they want to like gases can, but they are free to move as long as
they stay within the volume of the other molecules. This means that a liquid will be able to be contained
within a container and since they can tumble, they will actually take the shape of the container. This also
leaves the liquid with a surface – a “side” so to speak. There is not a lot of free space between the
molecules, which means that liquids really cannot be compressed except very slightly. Good thing too –
because we rely on this premise every time we press down the gas pedal on our cars. There is liquid in
there – called brake fluid – and if liquids compressed like gases did – our cars would not stop quickly.
They do flow and diffuse, but more slowly than gases do. Add cream to your coffee – if you do it slowly
enough you will form two layers based on density – give it a swirl with a spoon and you can mix the
liquids quickly (they will eventually mix together given time as the molecules tumble, but it takes time
and your coffee will get cold ), but simply putting two liquids together does not always mean mixing
(diffusion) – think oil and water!!
Solids: the energy of attraction is high and the energy of motion is very low. Essentially the
molecules/ions are fixed due to the strong attractive forces that exist between the molecules. This means
no tumbling, and no large motion. The molecules are generally closer together in the solid than in the
liquid and since their positions are “locked” this gives the solid a specific shape. Consequently, solids
really do not undergo compression and their particles do not flow. Put two solids in a glass together –
like salt and sugar and they are not going to “mix” together on their own!
Phase changes are determined by the balance – or lack thereof! – of the attractive forces compared to the
kinetic or motion forces that affect molecules. As the temperature increases, the motion of the particles
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increases, thus the KE increases. As the motion increases, the particles are better able to overcome the
attractive forces that they might be experiencing. This allows for a phase change to occur either from the
liquid to the gaseous state or from the solid to the liquid. Conversely, lower temperature favors less
movement, thus allowing the particles to maximize their attraction with one another, favoring the liquid
or solid state.
Condensation: the phase change of a gas into a liquid
Condensation
When a strong vortex is formed, say at a wing tip, the swirl component of the velocity causes the velocities to become
considerably larger than the surrounding flow. In the latter regions, the flow speeds are still on the order of (or even
approximately equal to) those of the freestream. According to the steady state Bernoulli equation, the pressures should become
very small within the vortex. In many cases of interest, the flow density is approximately constant and the temperatures must
also drop well below those of the surrounding flow. If the ambient air is sufficiently humid, the low pressures and temperatures
will cause the water vapor to condense, forming a "cloud" in the low pressure vortex. In many cases, no condensation will be
seen if the aircraft is in steady, level flight. It is only when the pilot cranks the plane into a high-g maneuver that the lift
increases to the point where condensation can occur.
With further cooling, the particles move even more slowly and become fixed in place, as the liquid
becomes a solid (known as freezing). The opposite phase change from freezing is melting (or fusion).
We have come to think of freezing as a temperature thing – but think of it as fixing molecules in place –
like freeze tag. You were not really frozen!! You were however, stuck in a particular spot and
unmovable. Freezing does not mean cold!! For example: gold freezes (solidifies) at 1064oC.
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Vaporization: the phase change of a liquid into a gas
The plume of steam above this cup of coffee is actually a cloud of tiny drops of water. The evaporation of water is an important
mechanism for cooling the coffee; the recondensation of the water delivers the energy to the surrounding air.
There is a symbol, H, which describes the enthalpy or the heat of the system. Is heat given off or must
heat be put into the system for a phase change to occur? As molecules in a gas get closer together and
condense to form a liquid and then freeze to become a solid, heat is released or given off. Thus
CONDENSATION and FREEZING are EXOTHERMIC processes. For the reverse, energy in the form of
heat must be put into the system in order to overcome the attractive forces seen in solids and liquids.
Thus, for solids to liquids (melting) and liquids to gases (vaporization) MELTING and VAPORIZATION
are ENDOTHERMIC processes.
Remember that for exothermic reactions, heat is lost, given off, thus the final amount of heat is lower
than the initial amount of heat. This is defined as a -H value while for endothermic reactions, the heat
was added to the system, thus there is more heat at the end than the beginning thus a +H value. The
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VALUE or the mathematical quantity will be the same when looking at the same species, but the sign
changes to indicate whether heat was given off or absorbed. Measurements of H are done per mole
and measured at a pressure of 1 atmosphere and the temperature at which the change takes place.
Reactant
H2O(l)
H2O(s)
H2O(g)
H2O(l)
Product
H2O(g)
H2O(l)
H2O(l)
H2O(s)
H (kJ/mole)
Hvap = 40.7
Hfus= 6.02
-Hvap = -40.7
-Hfus= -6.02
Endo/Exo
Endothermic
Endothermic
Exothermic
Exothermic
Water is a typical example of the fact that it takes less energy to melt substances (fusion) than it does to
vaporize them (convert liquids to gases). We are not completely disrupting the intermolecular forces
when we melt, but that it takes even MORE energy to separate molecules that have an attraction to one
another when we convert liquids to gases as gases have no intermolecular forces between the molecules!
The three states of water are stable at their various temperatures and at regular atmospheric pressures.
Other substances have phases that are not stable and thus do not naturally exist. CO2 is known to us as a
gas at room temperature and pressure. It can also be turned into a solid and is known as dry ice at
certain temperatures but still at regular pressures. The phase change from solid CO2 to gaseous CO2
does not go through a liquid phase. The solid does not melt! Instead, it goes directly from a solid to a
gas, which is known as sublimation. Other solids, like iodine in the solid form sublime. Ice will also
sublime from the solid to the gas state. Deposition occurs when a gas turns directly into a solid.
Hsubl is the enthalpy change when 1 mole of substance sublimes. It is a combination of solid converting
to liquid and liquid converting to gas all in one step. Thus, Hsubl is equal to the sum of Hfus and Hvap.
This makes Hsubl a positive number, indicating that energy was put into the system to convert the solid
to the gas.
Phase changes occur every day and propagate life on this planet.
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We can examine phase changes in detail by looking at a heating-cooling curve for a particular substance,
which shows the changes that occur when heat is added to or removed from the system at a constant
rate.
STAGE 1: at a T=130oC water is a gas. The gaseous water molecules have high kinetic energy (lots of
motion) and a very low attraction for one another. Therefore, there are no real intermolecular forces
between the individual water molecules. As the temperature falls (heat removed) the molecules begin to
slow down and we are at the interface between the conversion of the gas to the liquid phase.
The heat from H2O(g) @ 130oC to H2O(g) @ 100oC can be indicated by taking the number of moles
of water (n) times the molar heat capacity of gaseous water (Cwater) and the temperature change
(T = Tfinal – Tinitial)
q = heat = n x Cwater x T
q = 2.50 moles x 33.1 J/moleoC x (100oC – 130oC)
q = -2482 Joules = -2.482 kJ
The negative sign indicates that heat is released.
STAGE 2: gaseous water is condensing. As the water molecules start to condense, the molecules get to
spend a little bit more time near one another and the attraction between the molecules increases as the
temperature decreases. This allows for little clusters of molecules to form which drop out of the gas
phase and become liquid. While examining a phase change we are holding the temperature constant.
We are on the temperature line between gas and liquid.
The change from H2O(g) @ 100oC to H2O(l) @ 100oC is the number of moles times the
-Hvap
value (since gas converted to liquid is the opposite of vaporization!)
q = heat = n x -Hvap
q = 2.50 moles x (-40.7 kJ/mole)
q = -102 kJ
This stage results in the greatest amount of heat being released because we are allowing the molecules to
come together from free molecules to clusters. It is a tradeoff of the kinetic energy with the potential
energy of the system. This is also the stage that would take the most energy to convert a liquid to a gas –
thus it is the stage the releases that energy when converting between a gas to a liquid.
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STAGE 3: liquid water is cooling. As the temperature drops, all the molecules are in the liquid state.
The temperature decreasing does not affect the phase again until the freezing point of the substance is
reached.
The heat from H2O(l) @ 100oC to H2O(l) @ 0oC, released is n x Cwater as a liquid x T
q = heat = n x Cwater x T
q = 2.50 moles x 75.4 J/moleoC x (0oC – 100oC)
q = -18850 J = -18.8 kJ
STAGE 4: water cooling to the point of freezing. The molecules in the liquid state have motion, the
tumbling motion that was previously mentioned, and as the temperature continues to drop, this motion
stops and intermolecular forces dominate. The molecules slow and align themselves in the crystal
structure of ice. Molecular motion continues, but only vibration of the atoms in their fixed positions.
As during condensation, the temperature is held constant at the phase change, and the heat
released is n x -Hfusion (opposite of melting!) as H2O(l) 2 0oC is converted to H2O(s) @ 0oC.
q = heat = n x -Hfus
q = 2.50 moles x -6.02 kJ/mole
q = -15.0 kJ
STAGE 5: water, now in the solid state continues to cool. The only motion that the atoms are capable of
is vibration. As the temperature continues to drop, this motion also decreases. The heat released can be
calculated based on the “final” temperature that we examine for continuing to cool the water.
The heat from H2O(s) @ 0oC to H2O(s) @ -40oC, released is n x Cwater as a solid x T
q = heat = n x Cwater x T
q = 2.50 moles x 37.6 J/moleoC x (-40oC – 0oC)
q = -3760 Joules = -3.760 kJ
Hess’s Law states that the enthalpy change for the overall chemical phase changes occurring is the sum
of the individual steps of the processes. Thus, the sum of all the q values from each stage will give us the
overall heat change for converting water as a gas at 130oC to water as a solid at -40oC.
qtotal = q1 + q2 + q3 + q3 + q4 + q5
qtotal = -2.48 kJ + -102 kJ + -18.8 kJ + -15.0 kJ + -3.76 kJ
qtotal = -142 kJ
1.) WITHIN a phase, a change in heat is accompanied by a change in temperature, which in turn
affects the kinetic energy of the atoms/molecules. As the temperature decreases, so does the
motion. The heat lost of gained depends on the amount of substance present, the molar heat
capacity (C) for that phase, and T.
2.) DURING a phase change, the temperature is constant, which in turn affects the potential energy
as the distance between the molecules is changing. At the temperature point which marks the
boundary between phases, both phases are present. The heat lost or gained depends on the
amount of substance and the enthalpy associated with that phase change (Hvap or Hfusion).
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Phase changes are seen everyday. We had one recently that you may remember, when it got cold
enough – we did not have rain - we had snow! We were all sent home from school  But then the
temperature warmed up, the snow melted, and we came to school the next day  We boil water on the
stove, and if you are like me, you forget that you are boiling water and walk into the kitchen and your
pot is empty  All of these changes occur in the open, under normal atmospheric pressures. In a closed
system under controlled conditions, phase changes can reach a type of equilibrium, or balance and can
be reversible.
Liquid-Gas Equilibria: Holding T constant
If a flask containing a liquid is held out in the open, slowly but surely, the molecules will vaporize and
float away. They are “lost” to the original container. They are in the gas state, and as a gas can go
wherever they want, and chances are, should they condense back into a liquid, it is not going to be in the
container! Molecules left behind in the liquid state fill in the place of the molecules that left (were
vaporized) and the process continues until all the liquid is gone.
If the liquid were contained in a sealed flask with all the air removed, the molecules that go into the
vapor state actually make the pressure of the system go up. They are contained in that space and
bouncing around the walls and running into things. More and more molecules will go into the gas phase
until that space is saturated with molecules and that space can hold no more molecules. This means that
the pressure inside the container will stabilize. We have saturated the area above the liquid with
gaseous molecules. As the gas molecules continue to move around the container, some of them will
come in contact with the liquid. If their kinetic energy is not sufficient enough, they will be “sucked”
back into the liquid state by the molecular attractions that they have with other molecules. As they reenter the liquid phase, this leaves room for another liquid molecule to pull away and turn into a gas. At
this point, equilibrium has been reached. The vapor pressure will remain constant. The rate of
condensation will equal the rate of vaporization. As the process continues over and over again, it is said
to be dynamic.
Thus, the molecules in the gas phase are in a state of dynamic equilibrium with molecules in the liquid
phase.
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The pressure exerted by the gas phase molecules is known as the vapor pressure of that liquid at that
temperature. If the size of the container were changed – perhaps to something smaller – the vapor
pressure would stay constant. There would be fewer molecules in the gas phase and but eventually,
dynamic equilibrium would be reached and the vapor pressure would be the same as if a larger
container were used. It is important to note that if you change any of the outside effects on the flask –
perhaps you would open the stop cock and let some of that vapor out, or perhaps put more vapor into
the flask – these changes will result in the system adjusting itself, until eventually equilibrium is
established again. Equilibrium is a state of balance, where everything is equal. It can be thought of as
the lower energy state.
The vapor pressure of a substance depends on the temperature. Raising the temperature makes
molecules move faster, and faster moving molecules have more kinetic energy. This means that the
likelihood of the molecules staying the in liquid state is lower. More molecules will be able to overcome
intermolecular forces in the liquid phase and turn into a gas. In general, the higher the temperature, the
higher the vapor pressure.
The vapor pressure also depends on how strong the IM forces are! If the forces of attraction are weak,
then it will be easier for moving molecules to overcome them and turn into the gas. If the IM forces are
strong, it will be harder and the molecules will need more kinetic energy to overcome the forces in order
to become a gas. In general, the weaker the IM forces, the higher the vapor pressure.
Temperature is the only factor that has a real effect on vapor pressures. Vapor pressures are determined
in a closed system, therefore they are sometimes referred to as equilibrium vapor pressures. If you
change the pressure in the system, the vapor will compensate and reach equilibrium again. This means
that initially if you increase the pressure on the system there will be a drop in the number of molecules,
but eventually the rate of vaporization will = the rate of condensation. If you raise the temperature
however, you are able to put MORE molecules into the vapor phase than would “normally” be there.
More particles in the vapor phase increases the pressure. This means that the rate of vaporization > the
rate of condensation.
In an open container, the atmosphere exerts pressure on the liquid molecules. As the temperature rises,
the molecules begin to move and leave the surface of the liquid more quickly. At some temperature, the
motion of the molecules is great enough to create little pockets or vapor inside the liquid. This
appearance of bubbles then leads to the bubbles erupting into what we call boiling. Thus, the vapor
pressure has exceeded the atmospheric pressure. Once the boiling begins we are stuck in a phase change
– converting liquid → gas and in a phase change the temperature stays constant. The temperature
required for a liquid to boil is directly dependant on the atmospheric pressure pushing down on that
liquid. The lower the atmospheric pressure, the lower the temperature needed for something to boil.
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Concept Test
Rank the following places in order of lowest boiling point
of water to highest boiling point of water
1. Mt. McKinley
2. Death Valley
3. Portland
Solid Liquid Equilibria: changes in T
If we could look at the individual molecules in a solid we would see them locked in place – only
vibrating. As the temperature is slowly increased, the kinetic energy of the molecules begins to increase
until the molecules are not just vibrating, they are moving. This movement gives them flow and we are
at the equilibrium temperature between the solid and liquid phases for that substance, known more
commonly as melting. The opposite of melting is when the liquid turns into the solid, commonly
thought of as freezing and this temperature is the same. At the equilibrium point an equal number of
molecules are melting as are freezing. If the temperature is increased even further, the kinetic energy
imparted to the molecules allows the phase change from the solid to the liquid to occur, until eventually
the sample is completely liquid. Again, while the phase change is taking place, the temperature remains
constant.
Solid Gas Equilibria: solids have a much lower vapor pressure than liquids, meaning they do not really
turn into the gas phase all that readily. However, some solids do undergo this phase change, known as
sublimation, which means they do have a high vapor pressure. Dry ice, solid room deodorizers, moth
balls, and iodine are some common examples of solids that sublime. A solid sublimes rather than melts
because the combination of IM forces and atmospheric pressure is not great enough to keep the
molecules close together in the liquid state.
To describe the phase changes that a particular substance undergoes as the temperature and pressure of
the system change, a phase diagram is constructed. This is a combination of the melting, vaporization,
and sublimation processes. All phase diagrams have the following components:
1.) phase regions: each region corresponds to a particular phase that the substance will adopt at a
given temperature and pressure. The regions indicate the most stable phase for the temperature
and pressure. In general, a solid is stable at low temperatures and high pressures, a gas is stable
at high temperature and low pressure, and the liquid phase is generally intermediate of these
two situations.
2.) lines separating the regions: the lines represent the equilibrium situation where the both phases
exist at the same time.
3.) the critical point: the liquid gas line ends at the critical point. At this temperature, no amount of
pressure can be applied to convert the sample into the liquid phase. The pressure at this
temperature is known as the critical pressure.
4.) the triple point: there is a place on the graph where the liquid, solid, and gas lines come together
and meet. At this location, at this temperature and pressure all three phases are in equilibrium
together. This means that a sample is subliming and depositing, melting and freezing, and
vaporizing and condensing all at the same time!
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Examining a phase diagram helps to explain the properties of the substance. By choosing a phase, you
can determine what pressure and temperature is necessary for that phase to exist.
The particular phase that a substance will be in and what conditions result in a phase change are the
result of intermolecular and intramolecular forces. Intramolecular forces are the bonds that hold the
atoms together in a molecule, while intermolecular forces are the interaction of molecules with other
molecules as a result of partial charges, or the attraction of ions and molecules.
The two types of forces are very different from one another. Bonds – be they ionic, covalent, or metallic
are very very strong and are – at times – very difficult to break. Bonds involve larger charges and atoms
that are closer together. Intermolecular forces are weak compared to intramolecular forces. They
typically involve smaller charges that are further apart. How far apart – or how close together depends
on which atoms are involved in bonding. We have already examined the bonding radii associated with
two atoms coming together to form an ionic, covalent, or metallic bond. This is the distance that
balances out the nuclei electron attraction with the nuclei-nuclei and electron-electron repulsion. There
is another distance that results when two Cl2 molecules come near one another. This is a longer distance
than the Cl-Cl bond. If Cl1 and Cl2 are bonded together to make ClA atom, and Cl3 and Cl4 are bonded
together to make ClB atom, this longer distance would be between Cl1 or Cl2 with Cl3 or Cl4. The
bonding distance, or the covalent radius would be the distance between Cl1 and Cl2 OR Cl3 and Cl4.
See Figure on next page:
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Covalent bond length
van der Waal distance
van der Waal radius
Covalent radius
From examining the distances you can see that they bond radii is less than the van der Waal radii. This
is due to the fact that the bond length allows for slight overlap of the electron clouds in order to
maximize the attraction of the electrons with the nucleus. However, the intermolecular force distance
does not account for electron cloud overlap and thus the atoms will not be as close together, ever, for van
der Waal radii distances. Bond lengths are always shorter than van der Wall lengths. Bond radii are
always shorter than van der Waals radii.
Intermolecular forces are often called van der Waals forces due to the fact that they are the interaction of
bonded molecules with one another. There are several types that we will discuss: ion-dipole, dipoledipole, H-bonding, ion-induced dipole, dipole-induced dipole, and dispersion or London forces.
When an ion and a polar molecule (a molecule that has a dipole moment) get near one another, the
charged species (the ion) will interact with the partial charged species (the polar molecule) such that the
oppositely charged and partial charged ends will attract one another. The most common and probably
the most important example is the dissolving of an ionic compound (ions) in water (a polar molecule).
The ionic compound will break down into ions which will be attracted to either the partially negative
oxygen or the partially positive hydrogens in the water molecule.
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When polar molecules lie near one another, they will orient themselves in such a manner that their
oppositely charged ends are near one another. In a solid, the molecules will lock into place with their
oppositely charged ends attracting one another. In a liquid, again the oppositely charged ends will
attract one another but the molecules are allowed to move and flip. They are not locked into place as in a
solid, instead, it is like a square dance. They are “partnered up” but constantly flip and roll around,
finding new molecules to interact with.
Dipole-dipole forces are very important for explaining some common physical properties – such as
boiling point of molecules. For molecules of similar molecular weight, the greater the dipole moment,
the larger the disparity in the partial charges (the more “ionic in nature the bond becomes), the greater
the forces between the molecules and the more energy it takes to overcome those attractions. Thus, the
greater the dipole moment, the more energy is needed to boil the compound (the higher the boiling
point). This means, that when given the dipole moments, you can predict the order of when the
compounds will boil!
Concept Test
Put the following species in order of increasing boiling point based
on their dipole moments, put the lowest boiling point on the LEFT
and the highest boiling point on the RIGHT
Molecule
CH3CH2CH3
CH3Cl
CH3OCH3
CH3CN
CH3CHO
Molecular Wt (g/mole)
44.09
50.48
46.07
41.05
44.05
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Dipole Moment (D)
0.1
1.9
1.3
3.9
2.7
The special ability of a molecule to form a hydrogen “bond” only occurs when a hydrogen is bonded to a
highly electronegative atom that has lone pairs – meaning an oxygen, nitrogen, or fluorine. ONLY!!!
The N-H, O-H, and F-H bonds (the actual bond here) are very polar, so the electron density is drawn
away from the hydrogen (if we calculated oxidation number on the H – what would it be for every
situation like this??) As a result, this partially positive hydrogen of one molecule is attracted to the
partially negative lone pair on a N, O, or F on another molecule. The hydrogen “bond” or attraction
between the partially positive hydrogen from one molecule to the electron density on an O, F or N, is
shown with a dashed or dotted line between the H and the other atom (O, N, or F). It is not a complete
bond – it is the interaction and attraction of oppositely partially charged species!
As long as the N, O, or F have lone pair electrons and there is a
partially positive H, H bonds can form.
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Concept Test
Which of the following molecules will form H-bonds?
H2N-C≡N
CH2=N=N
N = N
CH2
H bonding is another force that must be overcome in order to change phases. It keeps the molecules
together, specifically in the liquid state where molecules are free to move around and interact.
When comparing boiling points, if a molecule can H-bond, this extra force that must be overcome will
cause an enormous deviation because the H bonds require additional energy to break before the
molecules can separate from one another and enter the gas phase.
The strength of the H bond is only about 5% of a typical covalent single bond, but when combined, the
overall effect of H bonding can be enormous. H bonding helps to hold two strands of DNA together, but
is weak enough to allow the chains to separate for protein synthesis and cell reproduction.
Electrons are not fixed in orbits as Bohr originally proposed, they are actually in some three-dimensional
space which we call orbitals. They can be thought of, also, as clouds. And these clouds surround the
nucleus and are where we can find the electron 90% of the time. These atoms are not hard spheres, these
electron clouds are quite “malleable” in that, if exposed to something negative, like an electric field, the
cloud will distort away. In effect, the field induces a distortion of the electron cloud and the cloud shifts,
resulting in a temporary dipole moment. If the electric field is removed, the cloud returns to normal.
The electric field can be a battery, a charged ion, or even the partial charge of a polar molecule. When an
ion gets near a non-polar molecule, it induces a dipole moment in that non-polar molecule – resulting in
an ion-induced dipole interaction. When a polar molecule with a dipole moment gets near enough to a
non-polar molecule it can induce a dipole moment in the non-polar molecule resulting in a dipoleinduced dipole interaction.
How easy is it to distort an electron cloud of a molecule? That depends on the molecule of course!! The
ease with which an atom’s electron cloud can be distorted is called polarizability. Smaller atoms tend to
be less polarizable than larger atoms. Smaller molecules tend to be less polarizable than larger
molecules. The larger the atom or molecules, the larger the electron clouds and the farther away from
the nucleus the cloud is.
Polarizability increases down a group of atoms (or ions!) because size increases and larger atoms
or ions have larger electron clouds which are easier to distort
Polarizability decreases from left to right across a row for atoms as size of the atom decreases and
the effective nuclear charge can hold onto the electron more tightly, making them less distortable
15
Cations are less polarizable than anions. Cation have lost electrons, thus the remaining electrons
are held onto more tightly making the electron cloud less distortable. Anions are more
polarizable than either their parent atoms or cations as the electron clouds are larger and thus
more able to be distorted.
There is a force inherent in all molecules, polar or non-polar that occurs regardless of other
intermolecular forces that might be going on in the molecule. Atoms are not hard spheres. When they
come together to form molecules or compounds, the resulting conglomeration is still not a hard sphere.
The electrons are in constant motion moving around the atom or molecule. If the electrons are moving,
then there will be momentary oscillations of electron charge. Over time, the electron charge is
distributed uniformly around the nucleus or around the non-polar molecule. But at any moment in time,
there might be an instantaneous dipole moment as the electron move around the atom/molecule. This
instantaneous dipole can affect other atoms or molecules should they be close enough together. As one
molecule has an instantaneous dipole, it can induce a dipole in the neighboring atoms/molecules. The
result is a synchronized motion of the electron in the clouds which causes an attraction between them.
At low temperatures, this attraction among the dipoles keeps the atoms/molecules together. Thus,
dispersion forces are instantaneous dipole-induced dipole interactions.
Dispersion forces are weak, but they exist in any molecule. Therefore, except for small polar molecules
with larger dipole moments, or molecules that have H bonds, the dispersion force is the predominant
intermolecular force between identical molecules. The relative strength of the dispersion force depends
on the polarizability of the particle, which in turn, depends on its size. Dispersion forces increases with
the number of electrons, which correlates closely with molar mass. Generally speaking, particles with
more electrons also have larger molar masses! As molar mass increases, polarizability increases,
dispersion forces increase and boiling points increase. Again, this is another force that must be
overcome in order for a phase change to take place.
Molecules must be very close together for this weak force to take affect
In He-He, bond energy is ~0.1 J/mol
(compare to 102 - 103 kJ/mol for covalent bonds ~ 1 million times smaller!)
As size increases, boiling point increases due to London forces
16
For non-polar substance with identical molecular weights, the strength of the dispersion force may be
influenced by the molecule’s shape. Shapes that allow more points of contact have more space over
which the electron cloud can be distorted, which allows for stronger attractions to be made. Comparing
the straight chain pentane with neo-pentane (a branched form of pentane that contains the same number
of carbons and hydrogens as pentane) shows the different shapes and how that affects dispersion forces.
The straight chain pentane is oblong in shape – like a cylinder. Neopentane however, is more spherical
in nature with its branches coming off the central carbon. The cylindrical form allows for more points of
contact between two identical pentane molecules than the spherical form does. Thus, the straight chain
molecule will have a higher boiling point.
Liquids share the properties of solids and gases and that makes them difficult to understand at times.
Liquids have and are under the influence of intermolecular forces as are solids, yet they can be random
at times since their molecules are not locked into place, like gases.
Surface Tension
In a sample of liquid, intermolecular forces affect the surface of a sample differently than they affect the
interior of the liquid. Interior molecules are attracted by others on all sides. Molecules on the surface are
only attracted by the molecules on the side and the molecules below it. As a result, the molecules on the
surface are pulled inward and downward which causes the surface to behave like a skin. In order for a
molecule to be on the surface, or to increase the surface area of the liquid, the molecule must overcome
the attractions within the interior of the liquid which requires energy. The stronger the forces are
between the particles in a liquid, the greater the surface tension.
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Capillary Action
The rising of a liquid through a narrow space against the pull of gravity is called capillary action. This
phenomenon happens when your blood is tested at the doctor’s office and they prick your finger and
then draw a small sample out for a quick analysis (ex. for anemia). The capillary action results from the
competition between intermolecular forces within the liquid (cohesion) and those between the liquid and
the tube walls (adhesion).
There is adhesion between the walls of glass (SiO2) and water are stronger than the cohesion
between the water molecules, thus allowing the water to creep up the tube. At the same
time, the cohesive forces pull the water down creating the surface tension. That is why
a meniscus forms and it is concave. Mercury (on the left) is the exact opposite. Its cohesive
forces are stronger than its adhesive forces, giving rise to a convex meniscus that
pulls away from the walls of the cylinder.
Viscosity
When a liquid flows, the molecules move past one another. Viscosity is a liquid’s resistance to flow. The
stronger the intermolecular attractions in the liquid, the greater the resistance to flow, the more viscous
is the liquid. Both gases and liquids flow, but liquids are more viscous than gases since they have
stronger intermolecular forces operating over shorter distances.
Viscosity decreases with heating. Heating disrupts intermolecular forces and thus makes the liquid
more likely to flow. Size also affects viscosity. We already talked about the fact that the more points of
contact there are for molecules, the greater the intermolecular attractions between them. Thus, the more
points of contact there are for molecules, the more viscous they are. Longer molecules (the cylindrically
shaped ones) are more viscous than their same molecular weight spherical counterparts (pentane and
neopentane for example).
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Water is an amazing little molecule. Oxygen, O2, has a molecular mass of 32 amu or 32 g/mole and we
know that O2 is a gas at room temperature. Water, H2O, has a molecular mass of 18 amu or 18 g/mole and it
is a liquid at room temperature. We take these simple things for granted. But if O 2 was not a gas and
H2O was not a liquid at standard temperatures and pressures, WE would NOT be here!! So water is
almost half the mass of oxygen – yet it is a liquid. WHY aren’t these two things reversed?? Because of
intermolecular forces! These small, very small, attractions between molecules keep our planet – and our
life (remember DNA!!) in working order. O2 is a nonpolar molecule – therefore the only intermolecular
force that it is subjected to are dispersion forces or other induced dipole forces. This keeps those oxygen
molecules AWAY from one another – making oxygen a gas. Water has a dipole moment. It is a polar
molecule – it is subjected to London dispersion forces (as are all molecules) but it has that oxygen there
with lone pairs and a dipole moment – which means the oxygen is partially negatively charged and the
hydrogens are partially positively charged and this makes water an ideal molecule to participate in H
bonding. Water has a tetrahedral VSEPR geometry and a bent/angular molecular geometry. This shape
and the dipole moment allow for water to engage in 4 hydrogen bonds per 1 molecule of water! That is a
LOT of extra attraction/interaction that must be overcome giving rise to water’s unique physical
properties.
H is represented by the white spheres, O by the red spheres
And the lone pairs on O by the purple spheres
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Water is often called the universal solvent. This is due to its polarity and its ability to H-bond. Water
dissolves ionic compounds through ion-dipole interactions that separate the ions in the solid yet keep
them in solution. The partially negative O will surround the positive cations while the partially positive
H’s will surround the negative anions from the ionic solid. Water can solubilize polar species like
alcohols and sugars by hydrogen bonding with them. Water can even solubilize gases through dipoledipole and London dispersion forces.
Water has a very high – completely unpredicted – specific heat capacity. Heat capacity is defined as the
measure of heat absorbed by a substance to cause a given temperature rise. When energy is put into a
system of water, some of that energy causes molecules to vibrate, some to rotate, and some to increase
the average kinetic energy or speed of the molecules. It is the increase in the average kinetic energy that
we measure by temperature. However, for a given amount of energy put into a system of water, the
overall temperature does not rise all that much meaning water stays liquid and does not evaporate into a
gas! Remember, all those H bonds have to be broken first – before H2O(l) is converted to H2O(g) meaning
that the energy put into the system disrupts the H bonds first. With water covering 70% of our planet
and being one of the most important components for life on this planet – it would not be a good thing for
the oceans – or our bodies! – if a little input of heat caused water to boil.
H bonding also gives rise to the amazingly high surface tension that water has. You can make a pin float
on top of water due to its surface tension. It may take some practice – but you can try it. Except for
some metals and molten salts, water has the highest surface tension of many liquid. Pretty amazing for
such a small molecule! Water also has high capillary action which allows plants to stay hydrated
especially during the dryer seasons.
When water solidifies, the H bonds become frozen in space. In the liquids, the H bonds are free moving,
rotating and flipping and changing as the molecules move and rotate through space. When the
temperature drops, the water molecules and the H bonds become locked in a hexagonal form.
Snowflakes and ice crystals arise because of this hexagonal open structure of the ice.
The large spaces in the ice crystals give rise to the fact that ice actually has a lower density than does
liquid water. Ice floats on water. If the solid were denser than the liquid – which is true for nearly every
other substance, the surface of lakes would freeze and sink, icebergs would sink, then the surface would
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freeze again, and then sink, and then the surface would freeze again and sink. Slowly, our water would
turn into the solid state for the winter months and aquatic life would not survive. Water density changes
slightly as temperature changes and this allows for water to cycle in the system, keeping it all
oxygenated and distributing the nutrients. The water cycle (given in the beginning) helps to maintain
life on earth as well. Water helps with erosion – which can be both a good and bad thing! – it feeds our
plants, and us too! It keeps us cool, and gives us electricity just to name a few things!
Last term you were asked to look up information in the CRC. In the CRC there was a column that listed
the crystalline structure. Depending on the atoms involved, crystal structures can vary immensely.
There are needles, rhombics, hexagonals, just to name a few. You have probably seen some rocks – like
quartz for example, and seen how they differ in their shape. Even though both sugar and salt are white
solids – take a look at them up close in your hand and you can tell a difference in them based on their
shape. Salt crystals are very regular and cubic looking. Sugar crystals are more random and seem
almost broken or chunky.
Solids can be broken down into two categories: the first are crystalline solids which have order and
structure of the atoms of which they are composed. The second are amorphous solids have poorly
defined shapes and the atoms are not arranged in an orderly manner.
If you could see the atoms in a crystal you could see how they are arranged or ordered in that crystal.
There are several different types of crystal lattices to be aware of: simple cubic, body centered cubic, and
face centered cubic. The lattice is a series of repeating units, or atoms, which make up the crystal. The
unit cell is the smallest portion of the crystal, that if repeated in all three dimensions defines the crystal.
It is the smallest portion of the crystal that we use to define if it is simple cubic, body centered, etc . . .
The coordination number of a particle in a crystal is the number of nearest neighboring atoms that
surround that atom in a lattice.
Simple Cubic Cells: the centers of 8 atoms define the corners of the cube. Attractions pulls the atoms
together on an edge – but the atoms do not touch along the diagonal. The coordination number is 6, four
in its own layer, one above and one below.
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Body-Centered Cubic Cells: there is an atom in the center, which is surrounded by 8 atoms in the
corners. The atoms in the corners only touch the central atom. Each atom is surrounded by 8 neighbors,
four above and four below (it is easiest to look at the central atom to see this!)
Face-Centered Cubic Cells: there is no central atom, instead, each face on the cube has an atom in the
center. Again there are 8 atoms on the corners. The corners touch the atoms on the face but not each
other. The coordination number is 12.
One unit cell lies adjacent to another unit cell giving rise to the crystal. In order to determine how many
atoms make up the cell, you have to consider the unit in 3 dimensions. For a simple cubic cell, there are
8 atoms in the corners. Each one of those atoms will be shared with another cube on all sides. Think
about blocks stacked on top of one another. How many blocks will that corner touch?
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If we start surrounding our simple cubic block with other blocks we see that in the first layer, there are a
total of 4 blocks sharing that one corner. BUT – we can build higher! If we place another 4 blocks on top
of these 4, that corner will be shared by a total of 8 blocks!
Thus, each corner atom will be shared amongst 8 other cells. Thus, there are 8 corner atoms and each
will have 1/8 of its atom in 1 cell so 8(1/8) = 1 atom per unit cell!
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Now try for body centered cubic and face centered cubic. Remember, what we learned above – a corner
is always 1/8!
For BCC: the central atom is always contained as a whole unit and is surrounded by 8 atoms on the
corner. Thus the number of atoms in the unit cell are the central atom (1) + 8( 1/8) = 1+1 = 2 atoms per
unit cell.
For FCC: there are eight corners again, but this time there are atoms on the face. Each atom on a face
will only be shared by two cells, half of the face will be in one cell and the other half in the other. This
means the 6 face atoms are halved 6(1/2) + 8(1/8) = 3+1 = 4 atoms per unit cell.
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There are two types of holes in a close-packed structure. Octahedral holes are made from six anions
arranged in an octahedral pattern. These are larger than the tetrahedral holes made from four atoms
arranged in a tetrahedral. In a unit cell of a closed-packed arrangement there are eight tetrahedral holes
and four octahedral holes. Cations fill holes based on their size and the stoichiometry of the structure.
Visualizing the coordination number can be difficult at times as you have to think spatially in 3
dimensions. Below are some diagrams that might help you.
Simple Cubic
Coordination number = 6
Body-Centered Cubic
Coordination Number = 8
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Face-Centered Cubic
Coordination Number = 12
Properties of unit cells can be used to calculate the radii of the atoms that make up the crystal. In order
to perform calculations for unit cells it is useful to know as much information about them as possible.
Remember that the volume of a sphere =
4π r 3
and that the unit cells are defined to be cubed, so the
3
area of a cell is the (edge length)3. Below is a Table of Unit Cell Properties that will be useful for unit cell
calculations.
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Type of Cell
Atoms per cell
1
Coordination
Number
6
% Occupied
Space
52.36
simple cubic
body-centered
2
8
68.02
face-centered
4
12
74.05
Edge length
2r
4r
3
4r/ 2
Concept Test
Iron crystallizes in a body-centered cubic lattice with a unit cell edge of 286 pm.
Estimate the radius of the iron atom:
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The unit cells we find today are a result of the various ways atoms pack together. For particles of the
same size, the higher the coordination number the greater number of particles there are packed into a
certain volume. Therefore, simple cubic has the fewest number of particles, then the body-centered, then
the face-centered which has the most particles per area.
If you think about the simple cubic spheres in a single layer, there will be spaces between those spheres.
These are known as interstitial spaces between the particles. One could, theoretically stack another later
of spheres directly on top of the first layer. This does not seem too stable. How many times have you
seen oranges stacked like strands of pearls in the produce section of the grocery store? Not too often –
they tend to fall over! Stacking in this manner means there is a 52% packing efficiency with 48% empty
space! None of the common metals has a simple cubic lattice.
The body-centered cubic cell has a more efficient packing arrangement than does a simple cubic cell.
The % occupied space is 68% compared to the simple cubic 52%. Fourteen metals, including iron,
chromium, and all the alkali metals crystallize in body-centered cubic lattices.
There is a more efficient packing arrangement that packed the maximum number of atoms or other
spherical particles (like oranges at the grocery store), into a give n volume and is called closest-packing.
When you start with a layer of spherical atoms there is already a built in loss of space between the
spheres. When three spheres meet, there is an interstitial space. Notice the white areas between the
spheres:
When you see oranges stacked at the grocery store you notice that the next layer of oranges is not
directly over the top of the sphere below it, but rather, off-center, it actually drops down into this
interstitial hole area. Not all the small hollows will be able to be filled however, remember spheres take
up space! In fact, only half the small hollows can be filled with another sphere. The third layer will be
identical to the first layer. The fourth layer will be the same as the second. . . etc. Thus, there are two
“different” types of layers in terms of organization. Thus, this is termed an ABABAB arrangement.
What you end up with at the end, is a hexagonal arrangement of atoms, so this is known as hexagonal
close packing.
There is another closest packing which is very similar to the hexagonal close packing. Again the layers
start out with A followed by B which lies in the holes of A. Instead of repeating the A layer, however, a
new arrangement is placed on top. Remember only one-half of the holes could be filled on layer A by
layer B. Thus, if we place the third layer of spheres over top of those holes, we have a layer that is
different than A or B. It is able to lie in the extra holes from A but also is stabilized since the spheres fit
in the spaces in layer B as well. This gives rise to the ABCABC pattern. This is known as cubic closest
packing and is based on the face-centered cubic unit cell.
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The packing efficiency of both closest packing patterns is about 74%. There is no way to pack spheres of
the same size more efficiently. Magnesium, Titanium, and Zinc adopt the hexagonal closest structure
while Nickel, Copper, and Lead adopt the cubic closest structure.
Atomic solids: Individual atoms that are held together by dispersion forces. The noble gases are the
only examples of atomic solids. The IMF are very weak (dispersion forces are weak!) Melting and
boiling points are very low indicating that it is not too difficult to disrupt the dispersion forces in these
solids in order to cause a phase change. The temperatures to melt or boil the solids rise with size –
which is typical for samples under the influence of dispersion forces.
Molecular solids: There are many examples of molecular solids and they can be under the influence of
dipole-dipole, dispersion, and H bonding forces which accounts for their wide range of physical
properties. For the nonpolar substances, dispersion forces are the principle forces holding the molecules
together. Again, melting and boiling points generally increase with increasing size and molar mass. For
polar molecules, dipole-dipole and H-bonding forces dominate. Generally speaking, molecular solids
have higher melting and boiling points than do the atomic solids. But, the IMF are weak compared to
ionic, covalent, and metallic solids, so melting point are lower for molecular compared to those solids.
Ionic solids: The unit cell contains whole charged ions. Because of these complete charges (as opposed
to the partial charges that we see for dipole-dipole interactions), IMF are much stronger in ionic solids
than molecular or atomic. To maximize attractions, the cations are surrounded by as many anions as
possible and vice versa. The larger ion is packed and the smaller ion fills in the holes or gaps. The unit
cell will be able to represent the empirical formula for the ionic compound. It will be the smallest whole
number ratio which retains the chemical composition of the compound. Thus, NaCl indicates a 1:1 ratio
of anion to cation. Ionic solids can adopt several different crystal structures. NaCl for example, adopts
the cubic closest packed.
The properties of the ionic solids are due to the fact that the ions are fixed in position. This also means
that the attractions between the charged particles are locked in place. This creates a very strong “bond”
between the ions. Thus, ionic solids typically have very high melting points and low conductivity (no
free ions roaming around!). When enough energy is put into the solid and it melts, free ions exist and
the liquid is then able to conduct electricity. Ionic compounds are hard and brittle. Forces applied to
them cause the crystal to break rather than bend due to repulsion between the ions.
Metallic solids: previously we talked about metals and how the atoms bond together creating this “sea
of electrons”. Atoms in a metal are held together by metallic bonding in which the valence electrons are
not confined to individual atoms or to individual bonds but are allowed to flow freely through the metal
(termed delocalized). Metallic bonding is sometimes described as a lattice of positive ions bathed in a
sea of mobile electrons; these mobile electrons provide metals with their conductivity and other
distinctive properties.
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Most metals crystallize in one of the two closest packed structures. Metals have high electrical and
thermal conductivity, luster, and malleability.
Which all result from the presence of delocalized
electrons (review!). Metals have a wide range of melting points and hardness which are related to how
well the atoms are packed in the crystal structure and to the number of valence electrons. Group 2A
metals have higher melting points than metals in Group 1A (group 2 are +2 ions and group 1 are all +1
ions). Higher charges mean more attraction between the ions!
Network covalent solids: These are covalent substances that do not rely on intermolecular forces to hold
them together. These are known as network covalent solids. These compounds are held together by
covalent bonds which go throughout the sample, in three dimensions. Thus, their properties will reflect
the strength of only the covalent bonds that hold the atoms together. Two examples of extremely strong
covalent bonds are diamonds and quartz.
Covalent compounds are poor electrical conductors, even when melted or when dissolved in water since
electrical current is carried by either mobile electrons or mobile ions, and covalent compounds do not
have free electrons. Their electrons are localized in bonds or are localized as lone pairs on the atoms.
Also, no ions are present which would allow for mobility of the electrons.
You may read on your own about conductors, semi-conductors, and insulators.
The ideal perfect crystal is only obtained if you allow your crystals to grow very slowly. This generally
means allowing the sample to cool slowly so that the molecules or ions can come together in a very
ordered systematic way (for those people taking organic, re-crystallization is an easy – though not
always effective!! – way to purify your compounds). When crystals form rapidly, there is the greater
chance for defects to occur. Impurities can get trapped in the crystal, or the molecules/ions just won’t
quite come together perfectly. When doing synthetic organic chemistry – making large amounts of a
particular drug for examples – impurities can be a very very bad thing. In the inorganic field, however,
sometimes these impurities can turn out to be quite advantageous.
Sometimes impurities will give a sample improved properties, such as increased strength or hardness, or
perhaps it increases the substance’s conductivity. Alloys, such as brass (copper and zinc) or steel (iron
and carbon) can be harder than the components that come together to make them. The hardness of the
alloy will depend on the ratios of the components that are mixed together. Thus, the alloy can be
changed depending on the % composition of the components.
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