States of Matter and Intermolecular Forces Chapter 11 11-1 States and State Changes.
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Transcript of States of Matter and Intermolecular Forces Chapter 11 11-1 States and State Changes.
![Page 1: States of Matter and Intermolecular Forces Chapter 11 11-1 States and State Changes.](https://reader036.fdocuments.us/reader036/viewer/2022070413/5697bfd81a28abf838caf34e/html5/thumbnails/1.jpg)
States of Matter and
Intermolecular ForcesChapter 11
11-1 States and State Changes
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Solids• Particles have an orderly,
fixed arrangement• Fixed volumes and shapes
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Liquids• Particles move easily
past one another (have more energy)• Fixed volume, no fixed
shape
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Viscosity• Ability to Flow• Honey is very viscous
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Surface Wetting
Adhesion
• Stick to something else
Cohesion
• Stick to each other
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Capillary Action• The movement of
water up through a tube – because of adhesion and cohesion
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Surface Tension• Cohesive forces• Causes liquids to minimize surface
area• That’s why water drops are round
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Gas• Particles are
independent• Far apart• No fixed volume or
shape• Gases and liquids are
fluids
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Changing State
• Freezing – liquid becomes a solid• Melting – solid becomes a liquid• Evaporation – liquid becomes gas• Condensation – gas becomes liquid• Sublimation – solid becomes gas• Deposition – gas becomes solid
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Temperature, Energy, and State
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Evaporation• High energy particles change to gas• Causes the substance to cool
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Boiling Point
• The temperature at which bubbles of vapor rise to the surface• Also depends on atmospheric
pressure
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Intermolecular Forces11-2
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Attraction between Particles• Takes energy to separate particles (change state)• The stronger the force, the more energy it takes
• The boiling and melting point is a good measure of the strength of the force
• Strong force of attraction = high boiling point
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Force of attraction in Ions• Higher force of attraction then between molecules• High melting points
• Smaller ions larger force (NaCl > KCl)• Larger charge larger force (CaF2 > NaCl)
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Intermolecular Forces• The Force of Attraction between molecules
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Types of Intermolecular Forces
• Dipole-Dipole Forces• Hydrogen Bonds• London Dispersion Forces• All are short range• Little effect on gases• Many gases have low
boiling point (that is why they are gases)
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Polar Molecule
• A molecule that has an unequal distribution of charge• One end slightly positive, One end slightly negative• Caused by difference in electronegativity of the atoms
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Dipole-Dipole Forces• Interaction between polar molecules• Positive end of one molecule attracts the negative end of another
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Dipole-Dipole Forces and Boiling Point• The more polar the molecules, the stronger the force
between them, the higher the boiling point
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Hydrogen Bonds• When a hydrogen atom of one molecule is attracted to an atom of a different
molecule• Water
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Hydrogen Bonds• Can create a larger
difference in electronegativity• Also hydrogen is small
and has only 1 electron • Which increases the
bond strength• Which increases the
boiling point
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Hydrogen Bonds and Water• Water has unique
properties, because of hydrogen bonds• Can form multiple
hydrogen bonds Strong intermolecular forces
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Solid water is less dense than liquid water
• Ice Floats• Ponds freeze
from top down• Expanding ice
cracks rocks and concrete
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London Dispersion Forces
• The force that hold non-polar molecules together• The weakest of the
intermolecular forces• Explains why some non-
polar molecules are not gases
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London Dispersion Forces• Nonpolar molecules can become temporary dipoles (electrons move from side
to side)• Causes molecules to attract each other
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London Dispersion Forces• Nearby molecules always attract• The more
electrons, the stronger the force
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Energy of State Changes
11-3
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Enthalpy• The total energy of a system
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Entropy
• A measure of system’s disorder
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Enthalpy of Fusion• The energy added during melting or removed during
freezing• AKA the heat of fusion
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Entropy of Fusion• The increase of entropy when a solid melts
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Enthalpy of Vaporization• The energy added during evaporation
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Entropy of Vaporization• The increase of entropy when a liquid evaporates• Much larger than entropy of fusion
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The molar enthalpy of fusion• The heat energy needed
to melt 1 mol of a substance
For water it is 6.01 kJ/mol
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The molar enthalpy of vaporization
• The heat energy needed to evaporate 1 mol of a substance
For water it is 40.67 kJ/mol
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Phase Equilibrium11-4
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System• A set of components that are being studied
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Phase• A region that has the same
composition and properties throughout
Lava lamp – Two phases of liquid- Different chemical compositions
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Phase
• Water – Two phases, same chemical composition - Different States
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Dynamic Equilibrium
• The net amount of substance in a given phase stays the same• Eg. The rate of evaporation
equals the rate of condensation
Which of these?
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Vapor Pressure• The pressure exerted by a gas in equilibrium with a liquid
• Boiling point – The temp at which vapor pressure equals the external pressure
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As temperature increases, vapor pressure increases
• Normal Boiling Point – when vapor pressure equals the atmospheric pressure
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Phase Diagrams• A graph of the
relationship between the state of a substance and its temperature and pressure
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Phase Diagrams
• 3 lines• Vapor pressure for
liquid-gas equilibrium A-B• Liquid-solid
equilibrium A-D• Solid gas equilibrium
A-C
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Triple Point
• The temperature and pressure at which all three states are in equilibrium
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Critical Point• The temperature and pressure at which the gas and liquid states become
identical
• Called a supercritical fluid
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Supercritical Fluid
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