Transcript Unit 7 - States of Matter

Unit 7
Overview
 States and Properties of Matter
 Kinetic Molecular Theory
 Phase Changes
 Diagrams
 Phase Diagrams
 Heating/Cooling Curves
 Heat of Fusion/Vaporization
 Vapor Pressure
States of Matter
 Based upon particle arrangement
 Based upon energy of particles
 Based upon distance between particles
3 States of Matter
 SOLIDS — have rigid shape, fixed volume; external
shape can reflect the atomic and molecular
arrangement
 Reasonably well understood
 LIQUIDS — have no fixed shape and may not fill a
container completely
 Not well understood
 GASES — expand to fill their container
 Good theoretical understanding
Solids
 Definite shape & definite volume
 Particles of solids are tightly packed, vibrating about a
fixed position.
 Not easily compressible
 Little free space between particles
 Do not flow easily
 Particles cannot move/slide past one another
 Infinite number of free surfaces
Solids (Types)
Crystalline Solids:
highly regular
arrangement of their
components
Amorphous solids:
considerable disorder
in their structures
(glass, plastic).
Allotropes
Some chemical elements can exist in two or more different forms. Atoms
of the element bond together differently.
Diamond, carbon atoms bond
together in a tetrahedral lattice
Graphite, carbon atoms bond
together in sheets
Liquids
 Indefinite shape but definite volume
 Particles are tightly packed, but are far enough apart to
slide over one another.
 Not easily compressible
 Little free space between particles
 Liquids flow easily
 Particles can move/slide past one another
 Have one free surface
Surface Tension
 Intermolecular cohesive attraction causing liquid to
minimize its surface area
 Mostly in polar molecules and liquid metals
Capillary Action
Attraction of the surface of a liquid to the
surface of a solid, which causes the liquid to
rise in a tube
Illustration of capillary
action for large and small
bore capillaries.
Viscosity
Liquids are fluids – they FLOW. Viscosity is the
resistance to flow.
Example: Syrup has a higher viscosity than water.
Glycerine, also called
glycerol, is a liquid with a
high viscosity
High viscosity = strong intermolecular forces
Viscosity decreases as temperature increases.
Gases
 Indefinite shape and indefinite volume
 Particles are very far apart and move freely
 Easily compressible
 There is a great deal of free space between particles
 Flow very easily
 Particles randomly move past one another
 Gases have no free surfaces
Plasma
 1879 - Sir William Crookes, an
English physicist, identified a
fourth state of matter, now
called plasma
 Plasma is by far the most common form of matter
 Plasma in the stars and in the tenuous space between them
makes up over 99% of the visible universe and perhaps most
of that which is not visible
Plasma
 An ionized gas
 Very good conductor of
electricity
 Composed of ions
 Affected by magnetic fields
 Indefinite shape and
indefinite volume
 Particles can move past one
another.
 Easily compressible
 Great deal of free space
between particles.
Products
manufactured
using plasmas
impact our daily
lives:
•Computer chips and integrated
circuits
•Computer hard drives
•Electronics
•Machine tools
•Medical implants and prosthetics
•Audio and video tapes
•Aircraft and automobile engine parts
•Printing on plastic food containers
•Energy-efficient window coatings
•High-efficiency window coatings
•Safe drinking water
•Voice and data communications
components
•Anti-scratch and anti-glare coatings
on eyeglasses and other optics
Kinetic Molecular Theory
ki⋅net⋅ic
1.
pertaining to motion.
2.
caused by motion.
3.
characterized by movement: Running and
dancing are kinetic activities.
Origin:
1850–55; < Gk kīnētikós moving, equiv. to kīnē- (verbid s. of kīneîn to move) +
-tikos
Source: Websters Dictionary
Kinetic Molecular Theory
Atoms and molecules are constantly in motion.
 SOLIDS — little movement between particles
 LIQUIDS — more space between them than a solid
does, but less than a gas
 GASES — molecules are moving in random patterns
with varying amounts of distance between the particles
Changes of State
Changing states requires a change in the energy of a
system. Changing states may also be due to the change
in pressure in a system.
Changes of
State
Phase Diagrams
Phase Diagrams
 (T) Triple Point: temperature and pressure at which all
three phases exist simultaneously in equilibrium
 (C) Critical Point: temperature and pressure beyond which
the liquid and solid phase is distinguishable (supercritical
liquid)
 molecules of a substance have too much kinetic energy to stick
together
 (page 453 of book = supercritical fluid)
Phase Diagrams
Water
Carbon
dioxide
Carbon
Heat of Fusion (formation)
 Energy that must be put into a solid to melt it
 Needed to overcome forces holding it together
 Heat of fusion given off when liquid freezes

Intermolecular forces within solid more stable and have lower
energy than forces within liquid so energy is released during
freezing
Heat of Vaporization
 Energy that must be put into a liquid to turn it into a gas
 Energy needed to overcome forces holding liquid together
 Heat of vaporization given off when gas condenses

Intermolecular forces become stronger when gas condenses so as gas
becomes liquid (more stable), energy is released
 Heat of vaporization larger than heat of fusion
 Many more intermolecular forces must be overcome in
vaporization than melting
 (intermolecular forces severed in vaporization but many carry
over between solids and liquids)
Heating/Cooling Curves
 As heat added to a substance in equilibrium,
temperature of substance can increase or the
substance can change phases, but both changes cannot
occur simultaneously
 Horizontal line at melting point is heat of fusion
 Horizontal line at boiling point is heat of vaporization
Heating/Cooling Curves
Vapor Pressure
 The pressure exerted by molecules as they escape the
surface of a liquid and become a gas
 As temperature increases, vapor pressure of a liquid
increases
 When vapor pressure of liquid increases to point where
equal with the surrounding atmospheric pressure, the
liquid boils
 The weaker the intermolecular forces, the easier it is
for molecules to escape the surface and turn to gas
 Weaker intermolecular forces = higher vapor pressure
Vapor Pressure - Equilibrium
 As number of gas molecules increases, higher probability
that gas molecule will hit surface of liquid and be
recaptured
 When there is even exchange of liquid and gas molecules,
vapor pressure becomes constant (dynamic equilibrium)
 Appears that nothing is happening in the system
 Vapor pressure of a liquid is the pressure exerted by its vapor
when the liquid and vapor states are in equilibrium
 When liquid/solid phase is in equilibrium with the gas
phase, the pressure of the gas equals the vapor pressure of
the substance