Chapter 11 States of Matter; Liquids and Solids Dr. Peter Warburton [email protected] .
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Transcript of Chapter 11 States of Matter; Liquids and Solids Dr. Peter Warburton [email protected] .
Chapter 11States of Matter; Liquids and Solids
Dr. Peter [email protected]://www.chem.mun.ca/zcourses/1011.php
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Gases
Gases are compressible fluids.
This behaviour arises because the gas molecules are in
constant random motion through mostly empty space.
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Liquids
Liquids are incompressible fluids.
This behaviour arises because the molecules of the liquid are in
constant random motion without much empty space to
move around in.
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Solids
Solids are incompressible and rigid.
This behaviour arises because the molecules of the solid can
only vibrate because they have almost no empty space to
move into.
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Figure 11.2
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Ideal gas law
We’ve treated gases in the past as behaving ideally, so that
PV = nRT
However, this gas law ASSUMES that molecules:
1) Have NO SIZE (no repulsive intermolecular forces)
2) DO NOT have attractive intermolecular forces
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All molecules have intermolecular forces (IMFs)
Repulsive IMFs (real molecules have size) will limit the compressibility of a
group of molecules.
SQUEEZE!
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All molecules have intermolecular forces (IMFs)
Attractive IMFs cannot be overcome at sufficiently low temperatures (molecules
have low average kinetic energies).
Lower T!
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Changes in state
We can often change the physical state of a substance (called a phase transition) by
changing the temperature and/or pressure by suitable
amounts.
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Phase transitions
From solid to liquid is fusion.
Change in enthalpy (or heat) of fusion is Hfus
From liquid to gas is vaporization.
Change in enthalpy of vaporization is Hvap
From solid to gas is sublimation.
Change in enthalpy of sublimation is Hsub
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Phase transitions
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Enthalpy of phase transitions
The enthalpy change of a phase transition tells us how much heat must
be added to (or removed from) a substance in a given phase so it changes
phase.Since ALL of the the energy is involved in the phase change, the temperature MUST REMAIN CONSTANT during the
change.
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Enthalpy of phase transitions
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Vapour pressure
Vapour pressure is the partial pressure (the part of the total
pressure that comes from a given substance) of the vapour (gas) above the liquid (or solid) phase
measured
at EQUILIBRIUM
at a GIVEN TEMPERATURE
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Equilibrium and vapour pressure
Vapour pressure should be measured when the vapour pressure
has STOPPED CHANGING.
This equilibrium means that the rate of molecules leaving the liquid (or solid) phase is BALANCED EXACTLY by the
rate of the molecules in the vapour joining the liquid (or solid) phase.
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Equilibrium and vapour pressure
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Temperature and vapour pressure
The vapour pressure depends on how easily molecules can overcome
attractive intermolecular forces that keep it in the liquid (or solid) phase.Higher temperatures mean each molecule, on average, has more
kinetic energy that COULD ALLOW the molecule to “escape” from the
attractive forces.
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Temperature and vapour pressure
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Temperature and vapour pressure
Since all groups of molecules AT THE SAME TEMPERATURE have the
SAME AVERAGE KINETIC ENERGY, then molecules that have WEAKER
intermolecular forces are MORE VOLATILE (have GREATER vapour
pressures at the GIVEN temperature) than molecules with STRONGER
intermolecular forces.
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Temperature and vapour pressure
In this Figure, the MOST VOLATILE (weakest IMFs)
liquid is on the left, while the LEAST
VOLATILE (strongest IMFs) is on the right.
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Boiling point
The boiling point of a liquid is the temperature
where the vapour pressure IS THE SAME
AS the external pressure.
Therefore, if the external pressure changes, the
boiling point temperature ALSO
changes.
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Normal boiling point
The normal boiling point of a liquid is the temperature where
the vapour pressure IS THE SAME AS an external pressure of
exactly1 ATMOSPHERE.
1 atm = 760 mmHg = 101.325 kPa
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Freezing (or melting) point
The freezing point of a liquid is the temperature where at which the phase transition from liquid to
solid occurs.
Since melting (fusion) is the exact opposite transition (solid to liquid)
the freezing point and melting point are IDENTICAL.
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Freezing and boiling points
Since the normal freezing and normal boiling points of a PURE substance are fixed
properties, measuring them is an easy and useful first step in
identifying unknown compounds.
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Intermolecular forces
The forces between molecules are electrostatic. They depend on charges (like charges repel, opposite charges attract), and the distance between the
charges.
Larger charges (like those found on ions) and smaller distances between
molecules tend to lead to stronger forces.
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Intermolecular forces
At everyday conditions, the forces between molecules tend to be
weakly attractive overall.
These intermolecular forces are generally called van der Waals
forces. However, vdW forces can be subdivided into two different groups.
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Dipole-dipole forces
We’ve already seen that some molecules have a permanent
molecular dipole.
Since full charges are not involved in molecular dipoles, these dipole-dipole
intermolecular interactions are relatively weak as compared to ionic bonds, where full charges are involved.
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Dipole-dipole forces
The larger the dipole moment (the molecules are more polar), the stronger the IMFs tend to be.
:
Cl-H
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London (dispersion) forces
Nonpolar molecules still interact with each other despite their lack of
permanent dipoles.
In a nonpolar molecule, “on average,” the electrons do not prefer one part of the
molecule over the other. However, at any given instant, they might not be
evenly distributed and so the molecule ends up having a “temporary dipole.”
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London (dispersion) forces
When a molecule with a temporary dipole comes close to another molecule, the electrons of the second molecule will try to move away from the negative partial charge of the first
molecule, leading to the second molecule having a temporary induced dipole.
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London (dispersion) forces
London forces tend to be very weak because the partial charges
tend to be small and fleeting.However, ALL chemical species
have London forces between them.
Variations in the strength of London forces depend on two factors.
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London (dispersion) forces
Polarizability – is the ability for electrons to move freely within the
molecule. The more freely electrons can move, the larger the induced
dipole can be.
Larger molecules and atoms are more polarizable, and generally
have larger London forces.
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London (dispersion) forces
Shape – The shape of a molecule plays a part in determining how the electrons can move in a
molecule. More compact shapes are usually more symmetrical and allow less contact
between molecules. They generally have smaller induced dipoles with weaker London forces.
Stronger London forces
Weaker London forces
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Hydrogen bonding
A hydrogen bond is an attractive
interaction between a hydrogen atom bonded to a very
electronegative atom (O, N, and F), and an unshared electron
pair on another electronegative
atom.
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Hydrogen bonding
Hydrogen bonds are really just a very special case of dipole-dipole forces.
H-F, H-O, and H-N bonds are very polar (larger partial charges)
Also, because the hydrogen is very small, it is possible for another molecule to
approach it very closely (short distance).
Hydrogen bonds are relatively strong intermolecular forces
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Hydrogen bonding (Figure 11.18)
With London forces, boiling points will increase
with molecular size (polarizability).
We EXPECT boiling points to follow the trend
CH4 < SiH4 < GeH4 < SnH4
and so on across the periodic table.
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Hydrogen bonding (Figure 11.18)
This trend is generally true, except for NH3, H2O, and HF because of the hydrogen bonds
(stronger than London forces) that can occur.
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IMFs and bonding at a glance
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Problem 11.37
For each of the following substances, list the kinds on intermolecular forces expected:
BF3 CH3CHOHCH3 HI
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Problem 11.43
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Classification of solids
Ionic solids are formed by regular arrangements of cations
(positive) and anions (negative). Because of the strong ionic bonds involved, the melting
points are usually high, and the solids are brittle and hard.
NaCl is an example.
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Classification of solids
Molecular solids are separate molecules held together through the mainly weaker dispersion, dipole-dipole and hydrogen
bond intermolecular forces. Weaker forces mean lower melting points and softer consistency.
Also, since electrons can’t easily move from one molecule to another, the solids
are nonconducting of electricity. Ice and sugar are examples.
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Classification of solids
Covalent network solids are formed by a large array of atoms. These solids are
formed from repeating units, but are almost better considered to be “one large molecule”.
If the arrays are three-dimensional, the solids are rigid, and so the melting points are
usually high, and the solids are hard but usually not brittle.
Quartz and diamond are examples.
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Classification of solids
Covalent network solids can also have structures of loosely held rigid sheets or
chains.
Graphite can be used as a lubricant because the sheets of carbon can easily
slide past each other.
Asbestos forms chains.
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Diamond and graphite
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Classification of solids
Metallic solids are large arrays of metal nuclei found in an “electron sea”.
Since the forces vary widely amongst metals we see widely variable melting
points and hardnesses. However, since electrons can move from one atom to another, the solids are conducting of
electricity and heat. Silver and iron are examples.
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Classification of solids
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Classification of solids