Solids, Liquids, and Gases

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Solids, Liquids, and Gases

• Chapter 16

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Kinetic Theory

• All matter is made of atoms and molecules that
  act like tiny particles
• These particles are always in motion. The higher
  the temperature, the faster they move
• At the same temp., more massive particles move
  slower than less massive particles

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Kinetic Molecular Theory

• Particles in an ideal gas
• have no volume
• have elastic collisions
• are in constant, random, straight-line motion
• dont attract or repel each other

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States of Matter

• Solid definite volume, definite shape
• Liquid definite volume, no definite shape
• Gas no definite volume, no definite shape
• Plasma gas-like, no definiteshape or volume

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Solids

• Have a rigid structure
• Particles have almost no freedom to change
  position, low KE
• Still vibrate around a fixed location

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Types of Solids

• Crystalline - repeating geometric pattern
• covalent network
• metallic
• ionic
• covalent molecular
• Amorphous - no geometric pattern

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Types of Solids
Ionic (NaCl)
Metallic
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Types of Solids
Covalent Molecular (H2O)
Covalent Network (SiO2 - quartz)
Amorphous (SiO2 - glass)
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Liquids

• Take the shape of their container
• Particles are held close together, but not
  attracted as strongly as particles in a solid,
  higher KE
• Classified as a fluid because it can flow

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Liquid Properties

• Surface Tension
• attractive force between particles in a liquid
  that minimizes surface area

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Liquid Properties

• Capillary Action
• attractive force between the surface of a liquid
  and the surface of a solid

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Liquids vs. Solids

• LIQUIDS
• Stronger than in gases
• Y
• high
• N
• slower than in gases

• SOLIDS
• Very strong
• N
• high
• N
• extremely slow

IMF Strength Fluid Density Compressible Diffusio
n
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Gases

• Free to spread in all directions
• Under standard conditions, gas particles move
  rapidly, high KE
• Exert pressure, but can be compressed into a
  smaller volume

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Real Gases

• Particles in a REAL gas
• have their own volume
• attract each other
• Gas behavior is most ideal
• at low pressures
• at high temperatures
• in nonpolar atoms/molecules

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Characteristics of Gases

• Gases expand to fill any container
• random motion, no attraction
• Gases are fluids (like liquids)
• no attraction
• Gases have very low densities
• no volume lots of empty space

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Characteristics of Gases

• Gases can be compressed
• no volume lots of empty space
• Gases undergo diffusion effusion
• random motion

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Describing Gases

• Gases can be described by their
• Temperature----K
• Pressure----atm
• Volume---L
• Number of molecules/moles----

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Temperature

• Always use absolute temperature (Kelvin) when
  working with gases!

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Plasma

• very high KE - particles collide with enough
  energy to break into charged particles (/-)
• gas-like, no definiteshape or volume
• stars, fluorescentlight bulbs, TV tubes

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Thermal Expansion

• Most matter expands when heated contracts when
  cooled
• ? Temp causes ? KE - Particles collide with more
  force spread out
• Examples
• Sidewalk dividers
• Thermometers

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Changes of State

• Changing states of matter requires energy

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Changes of State

• Evaporation
• Condensation
• Melting
• Freezing
• Sublimation

• Liquid -gt Gas
• Gas -gt Liquid
• Solid -gt Liquid
• Liquid -gt Solid
• Solid -gt Gas

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Phase Changes
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Phase Changes

• Evaporation
• molecules at the surface gain enough energy to
  overcome IMF
• Volatility
• measure of evaporation rate
• depends on temp IMF

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Phase Changes

• Equilibrium
• trapped molecules reach a balance between
  evaporation condensation

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Phase Changes
p.478

• Vapor Pressure
• pressure of vapor above a liquid at equilibrium

v.p.

• depends on temp IMF
• directly related to volatility

temp
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Phase Changes

• Boiling Point
• temp at which v.p. of liquid equals external
  pressure

• depends on Patm IMF
• Normal B.P. - b.p. at 1 atm

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Phase Changes

• Melting Point
• equal to freezing point

• Which has a higher m.p.?
• polar or nonpolar?
• covalent or ionic?

polar
ionic
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Phase Changes

• Sublimation
• solid ? gas
• v.p. of solid equals external pressure

• EX dry ice, mothballs, solid air fresheners

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Heating Curves
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Heating Curves

• Temperature Change
• change in KE (molecular motion)
• depends on heat capacity

• Heat Capacity
• energy required to raise the temp of 1 gram of a
  substance by 1C
• Specific heat - different values for each
  substance

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Heating Curves

• Phase Change
• change in PE (molecular arrangement)
• temp remains constant

• Heat of Fusion (?Hfus)
• energy required to melt 1 gram of a substance at
  its m.p.

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Heating Curves

• Heat of Vaporization (?Hvap)
• energy required to boil 1 gram of a substance at
  its b.p.
• usually larger than ?Hfuswhy?

• EX sweating, steam burns

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Heating Curve

• Video Clip