Transcript of Liquids and Solids - Mrs. Freeman's AP Chemistry
Chapter 10 Notes.notebookNov 151:29 PM
Types of Solids
Crystalline solids: Solids with highly regular
arrangement of their components
Amorphous solids: Solids with considerable
disorder in their structures.
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Lattice: A 3D system of points designating the positions
of the components (atoms, ions or molecules) that
make up the substance
Unit Cell: The smallest repeating unit of the
lattice structure.
XRay Diffraction: The method is to determine the
structure of a crystalline solid. Diffraction patterns of light
are measured and compared
Types of Solids
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Molecular Solids have discrete and relatively small
molecules at the lattice points that are held together by
IMFs
Examples: ice, dry ice, and sugars.
Types of Solids: Molecular solids
Properties: low melting/boiling points, poor conductors, low
densities, typically soft.
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Ionic Solids have ions at the lattice points and held
together by strong electrostatic attractions (stronger than
traditional IMF's).
Examples: Sodium chloride and lithium nitrate
Types of Solids: Ionic solids
Properties Include: High melting/boiling points, brittle (due to
cleavage), poor conductors as solids, good conductors as liquids
or in solutions, high densities.
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Types of Solids: Network Solids
Atomic/Molecular Network Solids: atoms/compounds at
the lattice points and held together by directional covalent
bonds (stronger than IMF and electrostatic)
Examples: graphite, diamond, and quartz (SiO2).
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Closet Packing: A metallic crystal structure where the atoms are
spherical and packed together and bonded to each other equally
in all directions
Metals are held together by nondirectional covalent
bonds where the electrons are delocalized throughout
the structure (weaker than directional covalent bonds)
Types of Solids: Metallic Solids
Properties: High melting/boiling points, good conductors, high
densities, malleable/ductile.
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Alloy: A substance that contains a mixture of elements
and has metallic properties
Substitutional Alloy: Some of the host metal
atoms are replaced by other metal atoms of
similar size (Sterling Silver93% silver 7% copper,
Pewter85% tin 7% copper 6% bismuth and 2%
antimony)
Interstitial Alloy: When some of the interstices
(Holes) in the closet packed metal structure are
occupied by small atoms (Steel contains carbon
atoms in the holes of an iron crystal)
Types of Solids: Metallic Solids
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Lack of rigidity (ability to flow)
(viscosity and surface tension)
High Density (compared to gases)
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Vaporization (Evaporation): Molecules of a liquid
escape the liquid's surface and form a gas
Condensation: Vapor molecules reform a liquid
Equilibrium: Two opposite processes exactly balanced
with each other
Vapor pressure: The pressure of the vapor present at
equilibrium
Liquids
L + G Gas
Heat Curve: A plot of temperature versus time for a
process where energy is added at a constant rate
Endothermic: Reactions require (take in) energy (heat)
Exothermic: Reactions release (give off) energy (Heat)
Breaking bonds is ALWAYS Endothermic
Normal Melting Point: The temperature at which the
solid and liquid states have the same vapor pressure
under conditions where the total pressure is 1 atm
Normal Boiling Point: The temperature at which the
vapor pressure of the liquid is exactly 1 atm
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Heat of vaporization (Enthalpy of vaporization):
The energy required to vaporize 1 mole of a liquid at
a pressure of 1 atm.
Heat of Fusion (Enthalpy of fusion): The change in
energy that occurs at the melting point when a solid
melts
Phase Change
Specific Heat Capacity: The energy required to raise
the temperature of one gram of a substance by one
degree Celsius
q = (m)(C)ΔT
q = energy (Joules or KJ)
m = Mass (grams)
C = Specific Heat Capacity
ΔT = Change in temperature
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Phase Change
CIce = 2.087 J/g
ΔHFus = 6.02 KJ/mol
CWater = 4.184 J/g
ΔHVap = 40.7 KJ/mol
CSteam = 1.996 J/g
How many kilojoules are needed to raise the temperature
of 112g of ice at 18.0 to steam at 109?
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Phase Change
Phase Diagrams: A diagram that represents the
phases of a substance as a function of temperature
and pressure (This is in a CLOSED system)
Triple Point: The point in which all states of matter
exist
Critical Point: The point where the critical pressure
meets the critical temperature. After this point there is
only a "intermediate fluid" that is neither vapor nor liquid
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Phase Change
Phase Diagram Fun Facts:
A liquid boils at the temperature
with the vapor pressure of the
liquid equals the external
pressure.
This is why it takes longer to hard boil an egg in the
mountains than it does at the ocean. Water will boil at a
lower temperature where there is less pressure
(mountains) meaning the egg has to cook longer to
hard boil in the lower temperature.
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Chapter 10 Notes.notebook
Properties of Solutions
Solution: Homogeneous mixture where a solute is
dissolved in a solvent
Dilute: Relatively little solute present
Concentrated: Large amount of solute present
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Properties of Solutions
Ways to Express Concentration:
Molarity: Moles divided by Liters (Mol/L which is
symbolized by M)
Mass Percent (Weight Percent): The percent by mass of
the solute in the solution
Mole Fraction: The ratio of moles of a given component to
the total number of moles
Molality: Moles of solute per kilogram of solvent
symbolized by a m
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Properties of Solutions
Steps of the Formation of a Liquid Solution
1). Separating the solute into its individual components
(expanding the solute)
2). Over coming intermolecular forces in the solvent to
make room for the solute (expanding the solvent)
3). Allowing the solute and solvent to interact to form the
solution
Steps 1 and 2 are commonly endothermic Step 3 is
commonly exothermic
Think about the intermolecular forces!!!
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Properties of Solutions
Enthalpy of Solution (ΔHsoln): The change in energy
associated with the formation of the solution. It is the
sum of the energy changes for all three steps.
ΔHsoln = ΔH1 + ΔH2 + ΔH3
If ΔH has a positive sign the step was endothermic
(energy absorbed) if ΔH has a negative sign the step was
exothermic (energy was released)
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Properties of Solutions
Enthalpy of Solution (ΔHsoln):
ΔHsoln = ΔH1 + ΔH2 + ΔH3
Let's look at dissolving oil in water:
Step 1: Expansion of the solute
medium +ΔH
Oil is a nonpolar substance that only has LDF holding it
together. These intermolecular forces are weak and
only require a small amount of energy to break apart,
however since oil is a larger compound it will take more
energy to "untangle" the molecules from each other
Step 2: Expansion of the solvent
large +ΔH
Water has very strong intermolecular forces
(Hydrogen bonds) that will require a lot of energy to
separate the molecules. This will result in a large
amount of energy
Step 3: solute and solvent interaction
small ΔH
The interaction between water (polar) and oil
(nonpolar) will be negligible causing a very small
release of energy
The medium and large positive step 1 and step 2 and the
very small negative step 3, cause the overall process to
be very large and positive. This means this solution will
unlikely not occur. (Enthalpy is not the ONLY determining
factor!!!)
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Properties of Solutions
Enthalpy of Hydration (ΔHhyd): Combines the
energy of step 2 and step 3 due to the fact that
hydration is the process in which water separates
from itself and surrounds the ions present in solution.
NaCl(s) ⇒ Na+(g) + Cl(g)
ΔH1 = 786 kJ/mol
H2O(l) + Na+(g) + Cl(g) ⇒ Na+(aq) + Cl(aq)
ΔH2+3 (ΔHhyd) =
783 kJ/mol
ΔHsoln = 786 kJ/mol 783 kJ/mol = 3 kJ/mol
The overall reaction is positive, but it still occurs
quite readily. This is due to the fact that enthalpy is
not the only factor when determining whether a
process will proceed readily. Entropy (chaos) is also
taken into consideration.
Let's look at dissolving salt in water:
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Polar solute, Polar solvent
Nonpoar solute, Polar solvent
Nonpolar solute,
nonpolar solvent
Large
Small
Small
Large
Large
Large
Small
Small
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Factors Affecting Solubility
Pressure effect: Pressure does not effect solids and
liquids but does significantly effect gases.
Chapter 10 Notes.notebook