Transcript Chapter 5 Homework:5.17, 5.18, 5.20, 5.24, 5.28, 5.29, 5.30, 5.37, 5.39, 5.51, 5.66, 5.74, 5.84

Chapter 5
Homework:5.17, 5.18, 5.20, 5.24,
5.28, 5.29, 5.30, 5.37, 5.39, 5.51,
5.66, 5.74, 5.84
Gases, Liquids, Solids
• Various forces hold matter together
• The strength of these forces decide the
state of matter
• There are attractive forces that hold
molecules together
• They are weaker than the forces that hold
ions together
• These attractive forces counter act Kinetic
Energy in molecules
• With out these forces, the kinetic energy that
particles posses would keep them moving
mostly in random, disorganized ways
• The kinetic energy increases with
temperature
• Therefore, the higher the temperature, the
greater the tendency of particles to fly around
randomly
• The physical state of matter thus depends on
a balance between the kinetic energy and the
attractive forces
• At high temperatures, molecules posses a
high K.E. and move so fast that the
attractive forces are too weak to hold them
together
• This is the gas state
• At lower temperatures, a gas condenses
to form the liquid state
• Molecules are still moving past each other
in the liquid state, but do so much slower
• At even lower temperatures, molecules no
longer posses the velocity to move past
each other
• This is the solid state
• Molecules have a certain number of
neighbors and these neighbors do not
change
• The strength of the attractive forces
does not change in the 3 states!!!
• Only the Kinetic Energy changes!!
• Most substances can exist in all three
states
• A solid, when heated to a sufficient
temperature usually melts and forms a
liquid
• That temperature is called the Melting
Point.
• Further heating causes the temperature to
rise to the point at which the liquid boils
and becomes a gas
• That temperature is the Boiling Point.
• Not all substances exists in 3 states.
– Examples:
Gas Pressure
• We live under a blanket of air that presses
down on us
• The amount it presses down on us
changes
• We measure the air pressure with a
barometer.
• Pressure is most commonly measured in
millimeters of mercury (mmHg)
• In can also be measured in torr
• At sea level, the average pressure of the
atmosphere is 760 mmHg.
• This is commonly called 1 atm
• 1 atm = 760 mmHg
= 760 torr
= 101,325 pascals
= 28.96 in Hg
- A Manometer is used to measure the
pressure of a gas in a container.
Gas Laws
• Gas Laws- relationships observed under
different temperatures, volumes, and
pressures
• Boyle’s Law- define the relationship
between pressure and volume by stating,
“At constant mass and temperature, the
volume of a gas is inversely proportional to
the pressure.”
Charles’s Law
• Defines relationship between temperature
and volume by stating, “At constant mass
and pressure, the volume of a gas is
directly proportional to the temperature in
Kelvins (K).”
Gay-Lussac’s Law
• Defines relationship between temperature
and pressure by stating, “At constant mass
and volume, the pressure is directly
proportional to temperature in Kelvins.
Combined Gas Law
• Problem: 3.00 L of a gas is at 2.00 atm.
What would the volume be in the pressure
was increased to 10.15 atm
Avogadro’s Law
• Defines relationship between mass and
volume by stating, “Equal volumes of a
gas at the same temperature and pressure
contain equal numbers of molecules.”
• The temperature and pressure of gases
does not matter when we are comparing
them
• However, chemist have selected some
standards for convenience:
1 atm = standard pressure
273 K (0oC) = standard temperature
• If a sample is at both standard
temperature and pressure, we abbreviate
by saying it is at STP
Another Standard
• The volume of one mole of gas (6.022 x
1023 molecules) at STP is 22.4 L
• This is called the Molar Volume
Ideal Gas Law
• Combining Avogadro’s law with the Combined
Gas law, we can write an equation that is valid
for any pressure, volume, temperature, and
quantity of gas.
• This is the IDEAL GAS LAW
PV=nRT
P= pressure, V= volume, n= moles, T=temp
R= Ideal Gas Constant = 0.0821 L atm mol-1 K-1
Ideal Gas constant
• The value of R can be calculated by using
the standards mentioned earlier, one mole
of any gas at STP occupies 22.4 L
• The ideal gas law holds for all ideal gases
• There are no ideal gases!!!
• The gases that we are surrounded by and
work with are real gases
• However, under most experimental
conditions, real gases behave sufficiently
like ideal gases that we can use the ideal
gas law
• Thus we can use PV=nRT to calculate any
of the variable as long as we know the
other 3.
Example
• If there are 5.0 g of CO2 gas in a 10.0 L
cylinder at 350 K, what is the gas pressure
within the cylinder?
Gas Mixtures
• In a mixture of gases, each of the gas
molecules act independently of all others
• Dalton’s Law of Partial Pressures- the total
pressure, PT, of a mixture of gases is the sum
of the partial pressures of each gas.
• Partial Pressure of a gas in a mixture is the
pressure that the gas would exert if it were
alone in the container
Example
• 5.0g of O2, 10.0 g of CO2, and 4 g of N2
are placed in a 10.0L vessel at 300K.
What is the total pressure of the vessel?
Kinetic Molecular Theory
• This theory explains the relationship
between the observed behavior of gases
and the behavior of individual gas
molecules within a gas
• There are six assumptions that it makes:
1) Gases consist of particles, either atoms
or molecules, constantly moving through
space in straight lines, in random
directions, and with various speeds
2) The average kinetic energy of gas
particles is proportional to the
temperature in kelvins
3) Molecules collide with each other, like
billiard balls, bouncing off each other and
changing direction
4) Gas particles have no volume
5) There are no attractive forces between
gas particles
6) Gas particles collide with the walls of the
container, and these collisions constitute
the pressure of the gas.
These 6 assumptions give us an idealized
picture of a gas
Real Gases
• In real gases, there are attractive forces
and the particles (atoms or molecules)
have volumes
• This is why there are no ideal gases
• But at STP, most real gases behave like
ideal gases
Forces of Attraction
• In general, the closer the molecules are, the
greater effect of intermolecular forces
• At room temperature and 1 atm, molecules are
so far apart, we can basically ignore the
intermolecular forces
• But as temperature decreases and/or pressure
increases, the molecules get closer together so
the effect of the intermolecular forces increases
• It is these forces that cause
Condensation and Solidification
• Condensation- change from gas to liquid
• Solidification- change from liquid to solid
• There are three types of intermolecular
forces:
– London Dispersion Forces
– Dipole-Dipole Forces
– Hydrogen Bonding
London Dispersion Force
• All molecules have intermolecular forces
• We know this because all gases can be
condensed to a liquid
• These forces that every molecule posses
are called London Dispersion Forces
• They occur due to the spontaneous
uneven distribution of electrons at any give
moment in time.
Dipole-Dipole forces
• These are the attraction between the
positive end of one molecule an the
negative end of another molecule
Hydrogen Bonding
• These are a lot like Dipole-Dipole forces,
only much stronger
• The positive end is a hydrogen bonded to
an Oxygen or Nitrogen and the negative
end is a lone pair of electrons on an
Oxygen or Nitrogen of another molecule
Strengths
• Hydrogen Bonds are the strongest of the
three forces
• The strength of a hydrogen bond range
from about 2-10 kcal/mol
• Even though they are the strongest,
relative to regular bonds, they are very
weak.
• The typical covalent bond between an
Oxygen and Hydrogen is 119 kcal/mol
• Hydrogen bonding is very important in
biological molecules
• They are what hold the double helix of
DNA together.
Properties of Liquids
• As the pressure on a gas increases, the
intermolecular forces have more of an
impact because the molecules pack closer
together
• Once the molecules pack so tightly that
almost all the molecules touch or almost
touch one another, the gas condenses to a
liquid
• Liquids have a set volume, regardless of
the container
• They do not expand to fill the container
like a gas
• Liquids also have very little space between
the molecules, so they are much tougher
to compress, unlike gases
• It is so tough that liquids are said to be
incompressible
• The density of liquids are greater than gas
because the same mass occupies a much
smaller amount of space, or volume
• The position of the molecules in a liquid is
random
• There is some empty spaces that the
molecules can slide into
• So the molecules are constantly sliding
past one another, changing their position
with respect to their neighbors
• This causes the liquids to be fluid,
meaning they have definite volume, but
not definite shape
• Liquids have surface properties as well
• One of these is called Surface Tension
• The strength of the surface tension is
directly proportional to the strength of the
intermolecular forces it possess
• Water has a very high surface tension
because of the very strong hydrogen
bonding among water molecules
Vapor Pressure and Boiling Point
• An important property of liquids is their
tendency to evaporate
• Evaporation:
• Equilibrium- A condition in which two
opposing physical forces are equal
Vapor Pressure
• Vapor Pressure- the pressure of gas in
equilibrium with its liquid form in a closed
container
• The vapor pressure is a physical property of a
liquid and changes with temperature
• As you increase the temperature of a liquid, you
increase the KE of its molecules so more can
escape to the vapor phase
• So as you increase the temperature, the
vapor pressure increases, until it equals
the atmospheric pressure
• At this point, bubbles of vapor begin to
form under the liquid surface and force
their way upward
• The liquid begins to Boil!!!
• The molecules the evaporate are those
with the highest KE
• The ones left behind have lower KE
• The temperature of a sample is
proportional to the average KE, so as the
average KE drops, the temperature goes
down.
• This is the cooling sensation you get when
exiting a pool. Water evaporates off your
skin!!
Boiling Point
• The boiling point of a liquid is the
temperature at which it vapor pressure is
equal to the pressure of the atmosphere in
contact with its surface
• The boiling point when the atmospheric
pressure is 1 amt is the normal boiling
point
Factors that affect Boiling Points
1) Intermolecular forces- the tighter the
molecules are held together, the more
energy is required to overcome this.
2) Molecular Shape- As surface area of the
liquid particles decreases, contact
between molecules goes down, so the
strength of london forces goes down,
meaning the boiling point will decrease
Solids
• When liquids are cooled, their molecules come
so close together and the intermolecular
attractions become so strong, that the random
motion stops, and a solid forms.
• Formation of a solid from a liquid is called
solidification or crystallization
• All crystals have a regular shape which often
reflects the arrangement of the particles within
the crystal
• Solids almost always have a higher density
than liquids
• Crystals typically have characteristic shapes
and symmetries
• Some compounds have more than one type of
solid state
– Example: Cabon has 5!!
• The packing of the carbon molecules is
different in all 5 states and leads to different
properties
– Diamond- very hard, very dense
– Soot- an amorphous solid meaning its atom’s have
no set pattern and are arranged randomly
Types of Solids
• Know Table 6.4 page 182
Type Make up Characteristics
Ionic
Molecular
Polymeric
Network
Amorphous
Example
Phase Diagrams
• Phase Diagrams- show all the phase
changes for any substance.
• Temperature is plotted on the x-axis
• Pressure on the y-axis