Download Kinetic Energy

Chapter 13 Notes
States of Matter
Kinetic Theory and Gases
• Kinetic Energy—Energy that an
object has due to motion.
• The Kinetic Theory is that tiny
particles form all matter,
and they are constantly in
motion.
Kinetic energy vs. Potential Energy
Kinetic refers to motion—so kinetic energy is
the energy of motion; this is different from
potential energy, which is the possible
amount of energy stored in something.
-the kinetic theory states that tiny
particles form all matter, that are in
constant motion.
Kinetic Theory and Gases
1) A gas is composed of particles that
are small, hard spheres with
insignificant volume and no particle
interaction.
Kinetic Theory and Gases
2) Particles in a gas are in constant
motion—they travel straight paths
unless they collide with another
particle or their container.
Kinetic Theory and Gases
3) All collisions are considered
elastic—no energy is lost to friction.
Kinetic Theory and Gases
4.) No kinetic energy is lost when
gas particles collide
– elastic collisions occur w/ other gas particles
or with the wall of the container
– energy can be transferred in collision but the
total kinetic energy of the 2 particles does not
change
Kinetic Theory and Gases
5.) All gases have same average kinetic
energy at the same temperature
• kinetic energy of motion (molecules are always
moving)
• Temperature is a measure of the average kinetic
energy of the particles in a sample of matter (at a
given temp., all gases have the same avg. KE)
Temperature
↓ in temp = ↓ in K.E. (molecules slow
down; theoretically, if you could lower
the temp enough motion would
cease)
Temperature
Kelvin scale is a direct measure of average kinetic energy
(eg. particles at 200 K have 2x as much nrg as at 100 K)
K= oC + 273 (0 oC = 273 K)
Which has more kinetic energy and does most damage to a
brick wall - a big pickup truck or a Honda Prelude?
wt. = 15,000 lbs
wt. = 3000 lbs
Kinetic Energy = ½ mv2
– big molecules move more slowly, lightweights move faster
– gases move from hi concentration → lo concentration
– rate they move depends on kinetic energy (in other words,
the size and velocity of particles)
Kinetic Theory and Gases
Effusion = gas escapes thru tiny opening
ex: hole in tire, air effuses and tire goes flat
ex: helium in balloon overnight vs air in balloon
Diffusion = gas A mixing with (moving thru) gas B
ex: perfume sprayed in one room, noticed in
next rm
ex: rotten egg
Kinetic Theory and Gases
Graham’s law of effusion:
Rate of effusion= 1/(sq. root of molar mass)
Graham’s law of diffusion:
Rate of diffusion=
Rate A = (sq. root of molar mass B/ molar mass A)
Rate B
Behavior of Gases
• Kinetic-molecular theory  a great deal of
space exists between gas particles
– Large amount of empty space between
the particles allows compressibility and
expansion of gas particles
Gas Pressure
• Kinetic theory explains the existence
of gas pressure.
• Gas pressure—the force exerted by a
gas per unit surface area.
Gas Pressure
• The force of one molecule hitting an
object is relatively small, but the result
of billions of particles of air hitting a
surface at once is significant.
Gas Pressure
pressure = force / unit area
To increase pressure (force/area):
1. more particles per unit area
a. decrease volume of container (↓ area)
b. add more particles
2. increase temp: ↑ speed of particles causing
↑ collisions
What happens as you increase altitude
(climb a mountain)?
Gravity pulls air particles in toward earth.
The air at higher altitudes has less air above
pushing down and fewer air molecules in a given
space. Atmospheric pressure decreases as you
gain altitude. Pilots gauge their altitude by
measuring pressure.
Atmospheric Pressure
A barometer measures atmospheric pressure.
The SI unit for pressure is the pascal (Pa).
Atmospheric pressure at sea level is about 101.3
kilopascals (kPa). Other units of measurement
are atmospheres (atm), mm Hg, and pounds per
square inch (psi).
1 atm = 101.3 kPa = 760 mm Hg = 14.7 psi
Comparison of Pressure Units
Units of Pressure (p390)
1 atm = the average atmospheric pressure at sea
level
kilopascal
Torricelli
mm mercury
inches mercury
pounds / in2
1 atm = 101.3 kPa
1 atm = 760 torr
1 atm = 760 mm Hg
1 atm = 29.9 in Hg
1 atm = 14.7 psi
Pressure conversion problems
1. Convert 190 mm Hg to atm
2. The pressure at the top of Mt Everest is 4.89
psi. How many mm of Hg is this? in. of Hg?
How many atm?
What is an absence of particles
called?
• A vacuum!
• No particles = no pressure
• Atmospheric pressure is the amount
of pressure from the particles in the
atmosphere colliding with objects.
STP
STP = Standard Temperature and Pressure
Since temperature and air pressure may vary form
place to place it is necessary to have standard
reference conditions for testing purposes
STP is commonly used to define standard conditions
for temperature and pressure
0oC or 273K
and
1 atm or 760 mm
Dalton’s Law of Partial Pressures
There are mixtures of gases in a container
– each type of gas contributes a fraction of the
particles which will supply a similar fraction of
the pressure
Dalton’s Law of Partial Pressures
At constant vol. & temp., the total pressure exerted
by a mixture of gases is equal to the sum of the
partial pressures
Ptotal= P1 + P2 + P3 + ….Pn
Dalton’s Law of Partial Pressures
example:
Air contains oxygen, nitrogen, carbon dioxide and
trace amounts of argon and other gases. What
is the partial pressure of O2 at 1 atm of pressure
if PN2 = 593.4 mm?
PCO2 = 0.3 mm, and Pother = 7.1 mm ?
Dalton’s Law of Partial Pressures
Partial Pressure
* colliding particles → pressure
* more particles → more pressure
# of particles often measured in moles
Does 1 mol O2 contain the same # of molecules as 1
mol H2?
1 mol = 6.02 x 1023 particles = 22.4 L
Does 1 L of O2 contain the same # of molecules as 1
L H2 ?
End of Daily Notes
Liquids and Kinetic Theory
Particles in a liquid still have kinetic energy—the
particles vibrate and spin and slide past each
other—but not as much as is present in a gas.
One of the differences between the two is that
particles in a liquid are attracted to one another.
The attraction brings the
particles closer together,
and hold it together with
other molecules. This also
gives rise to surface
tension.
Intermolecular forces
• Intermolecular forces- hold together identical
particles (drop of water), carbon atoms in
graphite, and the cellulose particles in paper
• All intramolecular, or bonding forces are
stronger than intermolecular forces
Dispersion Forces
• Dispersion forces weak forces that result from
temporary shifts in the density of electrons in
electron clouds (weakest intermolecular force)
– Example: Oxygen molecules are nonpolar (b/c e- are
evenly distributed); under the right conditions, oxygen
molecules can be compressed into a liquid; the force
of attraction between oxygen molecules is dispersion
or London forces
Dispersion Forces
• Dispersion forces cont:
– e- in an e- cloud are in constant motion
– When 2 nonpolar molecules are in close contact or
when they collide, the e- cloud of one molecule repels
the e- cloud of the other molecule.
– The e- density around each nucleus is, for a moment,
greater in one region of each cloud; each molecule
forms a temporary dipole
– When temporary dipoles are close together, a weak
dispersion force exists between oppositely charged
regions of the dipoles
Dispersion Forces
• Recall your Halogen gases (F, Cl, Br, I) all exist
as diatomic molecules.
– The # of nonvalence e- from fluorine to chlorine
to bromine, to iodine. B/c the larger halogens
have more e-, there can be a greater difference
between positive and negative dipoles and thus
stronger dispersion forces
Dipole - Dipole Forces
• Dipole – Dipole forces  attractions between
oppositely charged regions of polar molecules
since polar molecules contain permanent
dipoles
– Neighboring polar molecules orient themselves so
that oppositely charged regions line up
Dipole - Dipole Forces
For Example:
When hydrogen chloride gas molecules approach,
the partially positive hydrogen atom in one
molecule is attracted to the partially negative
chlorine atom in another molecule.
Hydrogen Bonds
• Hydrogen Bonds  type of dipole-dipole
attraction that occurs between molecules
containing a hydrogen atom bonded to a small,
highly electornegative atom with at least one
lone e- pair
Hydrogen Bonds
For example:
– for a hydrogen bond to form, hydrogen must
be bonded to either a fluorine, oxygen, or
nitrogen atom
– These atoms are electronegative enough to
cause a large partial positive charge on the
hydrogen atom, yet small enough that their
lone pairs of e- can come close to hydrogen
atoms
Rank the intermolecular forces in order
of increasing strength
Dispersion forces 
dipole-dipole forces 
hydrogen bond
Liquids
• Kinetic-molecular theory predicts the
constant motion of the liquid particles
• Individual liquid molecules do not
have fixed positions; forces of
attraction between liquid molecules
limit their range of motion so that the
particles remain closely packed in a
fixed volume
Liquids
• Like gases, liquids can be compressed
• The change in volume is much less than that of
gases b/c liquid particles are already tightly
packed together
Liquids
• Fluidity  ability to flow
– Liquids are less fluid than gases because of
intermolecular attractions
• Viscosity  measure of the resistance of a liquid
to flow
– As temp. increases, viscosity decreases
Solids (least KE)
• The particles in the solid move, but
don’t move around. They vibrate
around a fixed point.
• Most solids are crystalline—they have
definite repeating structure.
• Substances that have more than
one crystalline structure are called
allotropes.
Solids
MOLECULAR solids: covalent molecules
held together by intermolecular
attractions only, weaker than ionic or
metallic bonds so these have lower
melting and boiling points
ex: H2O, CO2, sugar, wax
Solids
COVALENT NETWORK solid : a
crystalline exception to the molecular
norm
ex: diamond
SOLIDS
IONIC Solids: held together by strong
attraction between + and – ions, hi
melting pt., form crystals :
ions arranged in orderly repeating pattern of
unit cells
ex: NaCl, KCl, MgSO4, NaOH
Solids
METALLIC solids: cations in a
sea of valence e-; most have
strong bonds & crystalline
structure & hi melting point
Solids
There are some substances that have
no crystalline structure at all. These
are called amorphous solids. There
atoms are randomly arranged with no
pattern.
Examples are rubber, plastic, glass,
asphalt, etc.
Changes of State
• We have discussed that the state of a substance
does not just depend on the temperature of the
substance, but also the pressure that it is under.
• -A phase diagram shows the conditions at
which a substance exists as a solid, liquid and
gas.
Phase changes that require energy
Vaporization = liquid turns to gas (vapor)
Evaporation = vaporization occurring only at the
surface (cooling process)
Melting = solid becomes a liquid
Vapor pressure = pressure exerted by a vapor over
a liquid
Boiling = vapor pressure equals atmospheric
pressure (cooling process)
Sublimation = solid changes directly into a gas
Evaporation
evaporation—conversion of liquid to a gas
when the surface of a liquid is not boiling
evaporation is a cooling process –
water (or sweat) absorbs heat
kinetic energy rises
surface water escapes the chaos and takes some of
the kinetic energy (aka temp) with it leaving the
cooler (slower moving) molecules behind to absorb
more heat. They suck the heat out till they too
escape.
Melting Point
• The temperature at which a solid
becomes a liquid is the melting point.
• As kinetic energy is added to a solid,
eventually the particles have so much
energy that they overcome the
interaction between particles and
vibrate and spin themselves right out of
their structure.
Boiling Point
the boiling point is the temperature at which the vapor
pressure of the liquid is just equal to the external
pressure.
-bubbles of vapor form throughout the liquid as the
molecules with the highest kinetic energy go from
the liquid phase to a gas and escape into the air.
-boiling and evaporation are both cooling
processes for a liquid. In each case, the molecules
with the highest amounts of kinetic energy are
leaving the liquid and entering the gaseous phase.
Phase changes that release energy
Condensation = gas or vapor becomes a liquid
(reverse of vaportization; H-bonds form in liquid
water and energy is released)
Deposition = gas or vapor becomes a solid without
first becoming a liquid (reverse of sublimation)
Freezing = liquid converts to a crystalline solid
Phase Diagram
The most interesting thing on a phase diagram is a
triple point—the only set of conditions where a
substance can exist as a solid, liquid and gas
simultaneously.
-Looking at a phase diagram, you can see that
there is usually a point where a substance will
go straight from the solid phase to a gas. This is
called sublimation.
Phase Diagrams
Phase diagram for H2O
• Sublimation—going
directly from solid to
gas
• Triple point—the
one set of
conditions where a
substance can exist
as a solid, liquid
and gas
simultaneously.
In Review…
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What is the kinetic theory?
What do we assume about collisions in gases?
What is gas pressure?
What is the phase with the most KE?
What are most solids?
What is the boiling point dependent on?
What is the triple point?
What is sublimation?