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![Page 1: Liquids and solids 10.1 Intermolecular forces u Inside molecules (intramolecular) the atoms are bonded to each other. u Intermolecular refers to the.](https://reader035.fdocuments.in/reader035/viewer/2022062511/55155d9b550346486b8b47a9/html5/thumbnails/1.jpg)
Liquids and solidsLiquids and solids
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10.1 Intermolecular forces Inside molecules (intramolecular)
the atoms are bonded to each other.
Intermolecular refers to the forces between the molecules.
These are what hold the molecules together in the condensed states.
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Intermolecular forces Strong
•covalent bonding• ionic bonding
Weak•Dipole dipole•London dispersion forces•Hydrogen bonding
During phase changes the molecules stay intact.
Energy used to overcome forces.
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Dipole - Dipole Remember where the polar
definition came from? Molecules line up in the presence of
a electric field. The opposite ends of the dipole can attract each other so the molecules stay close together.
1% as strong as covalent or ionic bonds.
Weaker with greater distance. Small role in gases.
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Hydrogen Bonding Especially strong dipole-dipole
forces when H is attached to F, O, or N
These three because-•They have high electronegativity.•They are small enough to get
close. Affects boiling point.
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Water
+
-+
Hydrogen Bonding Clip
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CH4
SiH4
GeH4SnH4
PH3
NH3 SbH3
AsH3
H2O
H2SH2Se
H2Te
HF
HI
HBrHCl
Boiling Points
0ºC
100
-100
200
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London Dispersion Forces Non - polar molecules also exert
forces on each other. Otherwise, no solids or liquids. Electrons are not evenly distributed
at every instant in time. Have an instantaneous dipole. Induces a dipole in the atom next to
it. Induced dipole - induced dipole
interaction.
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Example
H H H HH H H H
+ H H H H
+ +
LD Video
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London Dispersion Forces Weak, short lived. Lasts longer at low temperature. Eventually long enough to make liquids. More electrons, more polarizable. Bigger molecules, higher melting and
boiling points. Exist in all molecules, much weaker
than other forces. Also called Van der Waal’s forces.
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#36
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#38
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#39
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#40 (In Webassign)
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10.2 Liquids Many of the properties due to
internal attraction of atoms.•Beading•Surface tension •Capillary action
Stronger intermolecular forces cause each of these to increase.
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Surface Tension
Molecules in the middle are attracted in all directions.
Molecules at the top are only pulled inside.
Minimizes surface area.
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Capillary Action Liquids spontaneously rise in a
narrow tube. Intermolecular forces are cohesive,
connecting like things. Adhesive forces connect to
something else. Glass is polar. It attracts water molecules.
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Beading If a polar substance is
placed on a non-polar surface. •There are cohesive,•But no adhesive
forces. And Visa Versa
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Viscosity How much a liquid resists flowing. Large forces, more viscous. Large molecules can get tangled
up.
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Model for Liquids Can’t see molecules so picture
them in motion but attracted to each other.
With regions arranged like solids but•with higher disorder.•with fewer holes than a gas.•Highly dynamic
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Phases The phase of a substance is
determined by three things. The temperature. The pressure. The strength of intermolecular
forces. Changes of State Video
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10.3 Solids Two major types. Crystalline - have a regular
arrangement of components in their structure. (table salt)
Amorphous - those with much disorder in their structure. Said to be “frozen in place”. (glass, plastic)
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10.4 Bonding Models for Metals
Why do metal atoms stay together? How does their bonding effect their properties?
Two Models: Electron Sea Model: A regular array
of metals in a “sea” of electrons. Band (Molecular Orbital) Model:
Electrons assumed to travel around metal crystal in MOs formed from valence atomic orbitals of metal atoms.
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Electron Sea Model
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10.5 C & Si - Atomic Network Solids
Composed of strong directional covalent bonds that are best viewed as a “giant molecule”.
brittle do not conduct heat or electricity carbon, silicon-based graphite, diamond, ceramics, glass
Diamond - hardest natural substance on earth, insulates both heat and electricity.
Graphite - slippery, conducts electricity.
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Diamond - each Carbon is sp3hybridized, connected to four other carbons. Carbon atoms are
locked into tetrahedral shape.
Strong bonds give the huge molecule its hardness.
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Each carbon is connected to three
other carbons and sp2 hybridized.
The molecule is flat with 120º angles in fused 6 member rings.
The bonds extend above and below the plane.
Graphite is different
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This bond overlap forms a huge bonding network.
Electrons are free to move through out these delocalized orbitals.
The layers slide by
each other.
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10.6 Molecular solids. S8, P4, CO2, H2O Different molecules have different forces between them. These forces depend on the size of the molecule. They also depend on the strength and nature of dipole
moments. Non-Polar: Large molecules (such as I2 ) can be solids
even without dipoles. Polar: Dipole-dipole forces are generally stronger than
L.D.F. Hydrogen bonding is stronger than Dipole-dipole forces. No matter how strong the intermolecular force, it is
always much, much weaker than the forces in bonds. Stronger forces lead to higher melting and freezing
points.
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10.7 Ionic Solids The extremes in dipole dipole forces-
atoms are actually held together by opposite charges.
Huge melting and boiling points. Atoms are locked in lattice so they
are hard and brittle. Every electron is accounted for so
they are poor conductors - good insulators.
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What type of solid will each substance form? #68
Diamond Quartz PH3 NH4NO3
H2 Ar Mg Cu KCl C6H12O6
SF2
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10.8 Vapor Pressure Vaporization - change from
liquid to gas at boiling point. Evaporation - change from
liquid to gas below boiling point.
Heat (or Enthalpy) of Vaporization (Hvap) - the
energy required to vaporize 1
mol at 1 atm.
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Vaporization is an endothermic process - it requires heat.
Energy is required to overcome intermolecular forces.
Responsible for cool earth. Why we sweat.
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Condensation Change from gas to liquid. Achieves a dynamic equilibrium with
vaporization in a closed system. What is a closed system? A closed system means matter
can’t go in or out. What the heck is a “dynamic
equilibrium?”
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Dynamic equilibrium When first sealed the molecules
gradually escape the surface of the liquid.
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Rate of Vaporization = Rate of
Condensation Molecules are constantly changing
phase “Dynamic” The total amount of liquid and
vapor remains constant “Equilibrium”
Dynamic equilibrium
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Vapor pressure The pressure above the liquid at
equilibrium. Liquids with high vapor pressures
evaporate easily. They are called volatile.
Increases with increasing temperature. Decreases with increasing
intermolecular forces. •Bigger molecules (bigger LDF)•More polar molecules (dipole-dipole)
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Changes of state The graph of temperature versus
heat applied is called a heating curve.
The temperature a solid turns to a liquid is the melting point.
The energy required to accomplish this change is called the Heat (or Enthalpy) of Fusion Hfus
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-40
-20
0
20
40
60
80
100
120
140
0 40 120 220 760 800
Heating Curve for Water
IceWater and Ice
Water
Water and Steam
Steam
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-40
-20
0
20
40
60
80
100
120
140
0 40 120 220 760 800
Heating Curve for Water
Heat of Fusion
Heat ofVaporization
Slope is Heat Capacity
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Melting Point Melting point is determined by the
vapor pressure of the solid and the liquid.
At the melting point the vapor pressure of the solid =
vapor pressure of the liquid at 1 atm
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Boiling Point Reached when the vapor pressure
equals the pressure of the surrounding atmosphere.
Normal boiling point is the boiling point at 1 atm pressure.
Super heating - Rapid heating above the boiling point allows liquid state to exist above normal boiling point.
Supercooling - Rapid cooling below the freezing point allows liquid state to exist below the normal freezing point.
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Practice
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10.9 Phase Diagrams A plot of temperature versus pressure for a
closed system, with lines to indicate where there is a phase change.
Critical temperature: temperature above which the vapor can not be liquefied.
Critical pressure: pressure required to liquefy at the critical temperature.
Critical point: critical temperature and pressure (for water, Tc = 374°C and 218 atm). Liquid & solid are indistinguishable.
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SolidLiquid
Gas
Triple Point
Critical Point
Temperature
Pre
ssur
e
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Temperature
SolidLiquid
Gas
1 Atm
AA
BB
CCD
D D
Pre
ssur
e
D
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SolidLiquid
Gas
This is the phase diagram for water.
The density of liquid water is higher than solid water.
Temperature
Pre
ssur
e
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Solid Liquid
Gas
1 Atm
This is the phase diagram for CO2
The solid is more dense than the liquid The solid sublimes at 1 atm.
Temperature
Pre
ssur
e
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Like most substances, bromine exists in one of the three typical
phases. Br2 has a normal melting point of -7.2°C and a normal boiling point of 59°C. The triple point for Br2 is -7.3°C and 40 torr, and the critical point is 320°C and 100 atm. Using this information, sketch a phase diagram for bromine indicating the points described
above. – Based on your phase diagram, order the three phases from
least dense to most dense.
– What is the stable phase of Br2 at room temperature and 1 atm?
– Under what temperature conditions can liquid bromine never exist?
– What phase changes occur as the temperature of a sample of bromine at 0.10 atm is increased from -50°C to 200°C?