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Chapter 12

CHM120H5 Chapter Notes - Chapter 12: Covalent Bond, Electronegativity, Surface Tension


Department
Chemistry
Course Code
CHM120H5
Professor
Judith C Poe
Chapter
12

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Chapter 12
Intermolecular Forces, Liquids, Solids, and Phase Changes
12.1:Overview of Physical States and Phase Changes
Phase a physically distinct, homogenous part of system. Each physical state is a phase
Interactions b/w the potential energy and the kinetic energy of the particles give rise to the properties of each
phase:
o the potential energy in the form of intermolecular forces tends to draw the molecules together.
According to Coulomb’s Law, the electrostatic potential energy depends on the charges of the particles
and the distances b/w them
o the kinetic energy is associated w. the random motion of the molecules tends to disperse them. It is
related to the their average speed and is proportional to the absolute temperature
These interactions explain phase changes (changes in physical state from one phase to another)
A Kinetic-Molecular View of the Three States
1. Intramolecular forces exist within each molecule, the chemical behaviour of all three states is the same b/c
each consists of the same molecule
2. Intermolecular Forces exist b/w the molecules, they physical behaviour of these states is different b/c the
strengths of the forces differ
Each phase is accompanied by a standard enthalpy change, given in kilojoules per mole
Δfus < Δvap , it takes less energy to reduce the intermolecular forces for molecules to move out of their
fixed position (melt a solid) than to separate them completely (vaporize a liquid)
The heat of sublimation (solid gas), Δsubl , is a combination of melting and vaporizing. Hess’ Law
states that it equals the sum of the heats of fusion and vaporization
12.2:Quantative Aspects of Phase Changes
Stage 1: gaseous water cools
-at the starting temp Ek > Ep (energy of attraction)
-as the Ek decreases (due to decrease in temp) Ep become more important (attractions)
The heat (q) is equal to the amount (moles) x the molar heat capacity (C)(gas) x ∆T [ q = n x Cwater x ∆T]
Stage 2: Gaseous Water Condenses
-intermolecular attractions cause the slowest of the molecules to form microdroplets then liquid
-during phase changes, the temp and Ek are constant, at the same temp the molecules move farther b/w collisions
but their average speed is the same
-has a lower Ep and changes from gas to liquid, the heat is the amount x the negative heat of vaporization [q = n x -
Hvap]
-this stage contributes the greatest portion of the total heat released b/c of the large decrease in Ep
Type
Exo/Endo
Definiton
Condensation
Exothermic
Gas Liquid (temp drops, gas particles come
together to form a liquid)
- Δvap
Vaporization
Endothermic
Liquid Gas (increase in temperature =
higher K.E)
Heat of Vaporization:
Δvap
Freezing
Exothermic
Liquid Solid (decrease in temp, decrease in
K.E)
- Δfus
Melting(Fusion)
Endothermic
Solid Liquid (increase in temp and K.E)
Heat of Fusion:
Δfus
Sublimation
Endothermic
Solid Gas
Deposition
Exothermic
Gas Solid
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Stage 3:Liquid Water Cools
-the molecules in the liquid state continue to lose heat, and decreases in K.E (goes from 100ºC to 0ºC)
- The heat (q) is equal to the amount (moles) x the molar heat capacity (C)(liquid) x ∆T [ q = n x Cwater x ∆T]
Stage 4: Liquid Water Freezes
-at 0ºC, the sample loses Ep as increasing intermolecular attractions cause the molecules to align
- q = n x -∆Hfus
Stage 5: Solid Water Cools
-the solid furthers solidifies as temperature decreases
- q = n x Cwater x ∆T (C = for solid water)
-According to Hess’ Law the total amount of heat released is the sum of the heats released for the individual stages
-Within a phase, the heat flow is accompanied by a change in temperature, thus increase in Ek
-During a phase change, heat flow occurs at a constant temperature which is associated w. a change in Ep
The Equilibrium Nature of Phase Changes
-in a closed container phase changes are reversible and reach equilibrium
Liquid Gas Equilbria
1. Open System: nonequilbrium process, as the temperature increases the molecules have high enough kinetic
energy to evaporate
2. Closed System: two processes take place: some molecules escape and then collide w. the surface (the slower
ones condense)
3. Disturbing the System at Equilibrium
Decrease in Pressure = increase in volume, the rate of condensation temporarily falls below the rate of
vaporization (forward process is faster-liquid to gas) b/c fewer molecules enter the liquid than leave it
Increase in Pressure: decreasing volume, rate of condensation exceeds rate of vaporization b/c more
molecules enter the liquid than leave it (reverse process is faster-gas to liquid)
When a system in equilibrium is disturbed, it counteracts the disturbance & equilib is re-established
How does a liquid boil?
in an open container, the weight of the atmosphere bears down on a liquid surface. As the temp rises,
molecules move more quickly through the liquid. At some temp, the average Ek of the molecules in the liquid is
great enough for them to form bubbles of vapour in the interior and the liquid boils.
once boiling begins, the temp of the liquid remains constant until all of the liquid is gone
Solid-Liquid Equilbria
Melting Point the temperature at which the melting rate equals the freezing rate
pressure has a little effect b/c liquids & solids are incompressible
Solid-Gas Equilibria
a substance sublimes rather than melts b/c the intermolecular attractions are not great enough to keep the
molecules near each other when they leave the solid state
12.3: types of Intermolecular Forces
A. How Close Can Molecules Approach Each Other?
Bond Length and Covalent Radius:
o Bond length: the shorter distance, is b/w two bonded atoms (i.e two Cl atoms) in the same molecule
o Covalent Radius: ½ the bond length distance
Van der Waals distance and Radius:
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