chapter 12 · 2017. 1. 27. · 12.1 kinetic molecular theory • describes the behavior of matter...
TRANSCRIPT
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Chapter 12
States of Matter
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12.1 Kinetic Molecular Theory
• Describes the behavior of matter in terms
of particles in motion
– Makes several assumptions about the size,
motion, and energy of gas particles
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Assumptions of the Kinetic
Molecular Theory
1. Gases consist of small particles
2. Gases take up little volume relative to the
volume of empty space around them so
size of individual particles is zero
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3. Gas particles move in constant, random
straight lines until they collide with other
particles or with the walls of the container
– Collisions are elastic –
– Collisions cause gas pressure -
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4. Gas particles do not attract or repel each
other
5. The average kinetic energy of gas
particles is directly proportional to the
KELVIN temperature of the gas
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12.2 Gases
• Gases expand, diffuse, exert pressure,
and can be compressed because they are
in a low density state consisting of tiny,
constantly moving particles
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Explaining the Behavior of Gases
• Kinetic molecular theory helps explain the
behavior of gases
– Blowing up a balloon
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• Gases have low densities
– Why?
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• Gases can be compressed and can
expand
– Explain
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• Gases diffuse and effuse
– Diffusion = gas particles move from an area of
high concentration to low concentration
• Ex
– Effusion = gas particles escape through tiny
openings
• Ex
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12.3 Forces of attraction
• Intermolecular forces (dispersion forces,
dipole-dipole forces, and hydrogen bonds)
determine a substance’s state at a given
temperature
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Intermolecular forces
• Inter- means between or among
• Intermolecular forces can hold together
identical particles or two different types of
particles
• Weaker than intramolecular forces (bonds)
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Dispersion Forces
• Weak forces that result from temporary
shifts in the density of electrons in electron
clouds
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• Exist between all particles
– Weak for small particles
– Get stronger as the number of electrons
involved increases
– F2
– Cl2
– Br2
– I2
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Dipole-dipole forces
• Attraction between oppositely charged
regions of polar molecules
– Polar molecule =
• Neighboring polar molecules orient
themselves so that oppositely charged
regions align
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Hydrogen Bonds
• Dipole-dipole attraction that occurs
between molecules containing a hydrogen
atom bonded to a flourine, oxygen, or
nitrogen atom
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• Explain why water is liquid at room
temperature while compounds of similar
masses are gases
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12.4 Liquids and Solids
• The particles in solids and liquids have a
limited range of motion and are not easily
compressed.
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Liquids
• Kinetic molecular theory also applies to
liquids and solids
– Must take intermolecular forces into account
to apply it
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• Density and compression
– Much denser than gasses
• Due to intermolecular forces holding particles
together
– Incompressible
• Why can you compress a gas but not a liquid?
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• Fluidity – both gases and liquids are
classified as fluids because they can flow
and diffuse
– Liquids diffuse more slowly because
intermolecular attractions interfere with the
flow
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• Viscosity - measure of the resistance of a
liquid to flow
– Attractive forces – stronger intermolecular
forces = higher viscosity
– Particle size – larger molecules = higher
viscosity
– Temperature – lower temperature = higher
viscosity
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• Surface tension – the energy required to
increase the surface area of a liquid by a
given amount
– Caused by intermolecular forces pulling down
on the particles on the surface of a liquid
which stretches it tight like a drum
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• Stronger the attraction between particles
in a liquid = greater surface tension
• Surfactant – lowers the surface tension of
water by disrupting hydrogen bonds
between water molecules
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• Cohesion – force of attraction between
identical molecules
• Adhesion – force of attraction between
molecules that are different
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Solids
• Solid particles have as much kinetic
energy as liquids or gasses but much
stronger attractive forces between
particles
– Limit the motion of particles to vibrations
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• Density of solids –
almost always greater
than density of liquids
– Exception = water
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• Crystalline solids – solid whose atoms,
ions, or molecules are arranged in an
orderly, geometric structure
– Unit cell = smallest arrangement of atoms in a
crystalline solid that has the same shape as
the whole crystal
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• Categories of crystalline solids
– Classified based on the types of particles they
contain and how they are bonded together
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• Molecular solids
– Molecules are held together by dispersion
forces, dipole-dipole forces or hydrogen
bonds
– Most are not solid at room temperature
– Poor conductors
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• Covalent network solids
– C or Si, can form multiple covalent bonds
which allow it to take many forms
– Allotrope – element that can exist in different
forms at the same state
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• Ionic solids
– Made of cation + anion
– Each ion is surrounded by ions of the
opposite charge
– High melting point
– Brittle
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• Metallic solids
– Positive metal ions surrounded by a sea of
mobile electrons
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Amorphous solids
• Particles are not arranged in a regular,
repeating pattern
• Does not contain crystals
• Forms when molten material cools too
quickly for crystals to form
– Glass
– Rubber
– Some plastics
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12.5 Phase changes
• Matter changes phases when energy is
added or removed
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Phase changes that require energy
• Melting
– Heat flows from an object at a higher
temperature to an object at a lower
temperature
– Ice absorbs heat which does not raise
temperature but is used to break hydrogen
bonds
– When hydrogen bonds are broken molecules
can move further apart into the liquid phase
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• Melting point – temperature in which
forces holding a solid together are broken
and it becomes a liquid
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• Vaporization – process by which liquid
changes to vapor
– Vapor – gaseous state of a substance that is
normally liquid at room temperature
– Evaporation – when vaporization occurs only
at the surface of a liquid
– Vapor pressure – the pressure exerted by a
vapor over a liquid
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– Boiling – temperature at which the vapor
pressure of a liquid equals the atmospheric
pressure
– Energy being input causes molecules to move
around more and vaporize
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• Sublimation – changing from solid to gas
without becoming a liquid
– Dry ice
– Moth balls
– Solid air fresheners
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Heating Curve
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Phase changes that release energy
• Freezing
– Heat flows out of warmer object into cooler
object
– Molecules slow down & become less likely to
flow past one another
– Intermolecular forces cause the molecules to
become fixed into set positions
– Freezing point – temperature in which a liquid
becomes a solid
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• Condensation – process by which a gas or
vapor becomes a liquid
• Deposition – substance changes from gas
or vapor to solid without first becoming a
liquid
– frost
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Phase Diagrams
• Temperature and pressure both effect the
phase of a substance
– Have opposite effects
• Phase diagram – graph of pressure vs
temperature that shows which phase a
substance will be in under different
conditions.
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• Triple point = point at which all three
phases exist at the same time