Transcript Ch. 10 States of Matter

Ch. 10 States of Matter
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Ch. 10.1
Ch. 10.2
Ch. 10.3
Ch. 10.4
The Nature of Gases
The Nature of Liquids
The Nature of Solids
Changes of State
Ch. 10.1 The Nature of Gases
• Kinetic theory
– States that the tiny particles in all forms
of matter are in constant motion
– Makes 3 basic assumptions:
• Gases are composed of tiny particles with
very little volume and a great deal of empty
space between them; no attractive forces
• The particles move rapidly in random motion;
they travel independently in straight paths
and collisions occur
• All collisions are perfectly elastic – total
kinetic energy is conserved
Ch. 10.1 The Nature of Gases
• Gas pressure
– Force exerted by a gas per unit surface
area
• Atmospheric pressure
– Results from the collisions of air molecules with
objects
– Decreases as you move higher in the atmosphere
– Barometers are used to measure atmospheric
pressure (dependent on weather)
– SI unit is the pascal (Pa); two older units are mm
HG and atm
– One standard atm is the pressure necessary to
support 760mm Hg in a mercury barometer at 25oC
Ch. 10.1 The Nature of Gases
• Gas pressure
– Modern barometers are aneroid
• Measure the # of collisions of air molecules
• It is important to measure gases under a standard
condition
– STP is defined as OoC and 101.3kPa or 1 atm
– Kinetic energy and Kelvin temperature
• When heated, particles absorb some energy as
potential, the rest is kinetic and speeds up the
particles
• Particles have varying amounts of energy, so
average kinetic energy is used
• At absolute zero, 0 K, the motion of particles
theoretically ceases
– This has never been achieved in a lab
Ch. 10.2 The Nature of Liquids
• A model for liquids
– Particles that make up a liquid vibrate and
spin as they move from place to place
• this motion contributes to the kinetic energy
• Most particles do not have enough energy to
escape the intermolecular forces and enter the
gaseous state
• Liquids and solids are known as the condensed
states of matter
Ch. 10.2 The Nature of Liquids
• Evaporation
– Conversion of a liquid to a gas or vapor
is called vaporization
• When the process occurs in a liquid that is
not boiling, it is called evaporation
– Liquid particles with enough kinetic energy break
away from the surface and enter the gas phase
– Some particles collide with air molecules and
return to the liquid phase
• Liquids evaporate faster when heated
• Evaporation is a cooling process
– When particles with high kinetic energy escape,
the overall energy of the liquid is lowered
Ch. 10.2 The Nature of Liquids
• Evaporation in a closed container
has different results
– Particles that escape collide with the
walls if the container and produce a
vapor pressure
• the pressure of the gas above the liquid
– Particles will continue to leave the
liquid, but some will re-enter
• A dynamic equilibrium exists (see pg. 276)
• Rate of evaporation = rate of condensation
– Increasing temperature increases the
vapor pressure
Ch. 10.2 The Nature of Liquids
• Vapor pressure can be measured with a
manometer
• Boiling point
– The temperature at which the vapor pressure
of a liquid is equal to the external pressure
• Normal boiling point occurs under standard pressure
• Boiling occurs more easily at high altitudes, and
more slowly at low altitudes
– Boiling is also a cooling process
• The temperature never rises above boiling point
• Liquids can only boil faster if heat increases
• Vapor pressure is at the same temperature, but it has
a much higher potential energy
Ch. 10.3 The Nature of Solids
• A model for solids
– Particles in solids tend to vibrate about a fixed
point
• They are dense and incompressible
• They do not flow or take the shape of their containers
– Melting point
• When solids are heated, the particles gain kinetic
energy
• Organization of the particles is disrupted
• The melting point is the temperature at which a solid
becomes a liquid
– Disruptive vibrations overcome the strong forces
holding the particles together
– Melting and freezing are reverse processes
Ch. 10.3 The Nature of Solids
• Crystal structure and unit cells
– Crystals have sides or faces
– There are 7 types of crystal systems
– The shape of a crystal depends on the
arrangement of the particles
– The smallest group of particles within the
crystal system is called a unit cell
• There are 3 types of unit cells
– Some substances can take more that one
crystal form – these are called allotropes
Ch. 10.3 The Nature of Solids
• Amorphous solids
– Not all solids are crystalline, some are
amorphous
• their atoms are randomly arranged
• Rubber, plastic and asphalt are amorphous solids
• Glass is another type
– Sometimes called supercooled liquids
– Structure is intermediate between a crystalline solid
and a free-flowing liquid
– Does not melt, but gradually softens
– Breaks in irregular shapes
Ch. 10.4 Changes of State
• Phase diagrams
– Give the conditions of temperature and
pressure at which a substance exists as a
solid, liquid and gas
• The point at which all three curves meet is known as
the triple point
• Sublimation
– Change of state from a solid to a gas without
passing through the liquid phase
– Solids have a vapor pressure
• If the vapor pressure is high enough, a solid can
change directly to a gas
• Dry ice (solid CO2), and mothballs (naphthalene) are
examples