honors chemistry: chapter 9 - harlem school district 122! 26 13.2 forces of attraction what...
TRANSCRIPT
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Gases
• Kinetic – movement
• Kinetic theory – explains the
behavior of particles in terms of
their motions
3
Kinetic Theory of Gases
Kinetic Theory of Gases
1. Negligible volume
2. Move in rapid, constant straight-line motion
3. Collide elastically
• No net loss in kinetic energy
4. Far apart with no attractive or repulsive force
4
• Kinetic Energy – energy of motion
– KE = ½ (m)(v)2
• m = mass
• v = velocity
• Temperature – measure of average kinetic energy.
– Molecules within a substance don’t all move at the
same speed
• They have a constant average speed!
Kinetic Energy and Temperature
6
Kinetic Energy and
Temperature
• The higher the temperature, the
faster moving the particles
• At 0 K all molecular motion
would stop
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Explaining the Behaviors of Gases
• Low Density
– – Gases are far apart
• Low mass
• High volume
volume
massDensity
• Compression
– If gases are far apart,
they can be pushed
together
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• Expansion
– If gases are moving in constant straight-line
motion, they will expand to fill any container.
Explaining the Behaviors of Gases
9
• Diffusion – tendency of molecules to move from
an area of high concentration to low concentration
until uniformly mixed.
– If particles are in rapid straight-line motion, the will
spread out.
– Perfume sprayed in a room
Explaining the Behaviors of Gases
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• Effusion – rate at which gases escape from
a small hole in a container.
Explaining the Behaviors of Gases
Both diffusion and effusion depend upon the
velocity of the particles!
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Graham’s Law
Graham’s Law
The rate of effusion of a gas is inversely
proportional to the square root of the gases
molar mass.
RateA Molar MassB
RateB Molar MassA
=
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Graham’s Law Derived
If both gases are in the same
room, then they have the same
temperature.
Would they be moving at the same velocity?
KEKE BA
2
BB
2
AA Vm2
1 Vm
2
1
No, the lighter gas is faster!
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What happens if we play with this
equation?
A
B
2
B
2
A
m
m
V
V
2
BB
2
AA Vm2
1 Vm
2
1
A
B
B
A
m
m
V
V You get
Graham’s law!
Multiply both sides by 2 and get the velocities and masses on the same side.
Get rid of the squared values by taking the square root of both sides.
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Graham’s Law Problems
1. Finding Rate Ratios (How much faster…)
1. Decide what is “A” (listed first) and “B”
2. Setup equation appropriately and solve.
2. Finding Molar Mass
1. Establish rate ratio
2. Which gas is faster?
• This gas becomes the bigger number in the rate ratio and establishes what is “A” and what is “B”.
3. Setup equation appropriately and solve.
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Graham’s Law
How much faster does oxygen diffuse (or
effuse) than nitrogen.
RateA Molar MassB
RateB Molar MassA
=
Which is A? O2
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Graham’s Law
If O2 is A, then N2 would be B. The rate of
oxygen to nitrogen would be the x.
oxygen
nitrogen
MM
MM x
32.0
28.0 x
X = .935 Oxygen is .935 times faster
than nitrogen (less than 1
means it’s slower)!
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• Pressure – simultaneous collisions of billions of
molecules with an object
• Pressure =
– The force is created by the collisions
• More collisions higher pressure
• Harder collisions higher pressure
– Thumbtacks and snowshoes take advantage of
this to work.
Gas Pressure
force
area
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Gas Pressure
• Vacuum- empty space means no collisions
• Atmospheric pressure - air molecule collisions – We live at the bottom of an ocean of air
– If you go up a mountain, atmospheric pressure decreases because the depth of air above you is less.
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• Units of pressure
–Pascal (Pa)-SI unit
–Atmosphere (atm)-Typical sea level
pressure
–Millimeters mercury (mm Hg)-
barometric reading
Gas Pressure
760 mm Hg = 1 atm = 101.3 kPa = 14.7 psi
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Dalton’s Law of Partial
Pressures
Dalton’s Law
At constant volume and temperature, the total
pressure exerted by a mixture of gases is
equal to the sum of the individual (partial)
pressures.
Ptotal = P1 + P2 + P3 • • • •
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Dalton’s Law
• Dalton’s Law Correction
– Because most gases are COLLECTED OVER
WATER, you will often use this law to correct
the pressure.
Ptotal = Pgas + Pwater Solving for the gas yields:
Pgas = Ptotal – Pwater This is usually
the pressure used
in a gas law
problem (next
chapter)!
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13.2 Forces of Attraction
What separates gases from solids
and liquids?
(What do solids and liquids have that
gases don’t)
Attractive Forces
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Forces of Attraction
There are two types of attractive forces
1. Intramolecular forces – forces within a
compound that hold it together.
2. Intermolecular forces – forces between
compounds that hold these compounds
together.
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Intramolecular Forces
Force Model Basis of
Attraction
Energy
(kJ/mol) Example
Ionic Opposite
charges 4000-400 NaCl
Covalent
Shared
electron
pair
1100-150 H2
Metallic
Metal
cations and
delocalized
electrons
1000-75 Fe
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Intermolecular Forces
Forces Basis of
Attraction
Energy
(kJ/mol) Example
London
dispersion
forces
Momentary
dipole – shifting
of electron
cloud
40-.05 H2 to H2
Dipole-
Dipole
Partial charges
of polar
molecules 25-5 HCl to HCl
Hydrogen
Bonding H bonded to an
N, O, or F 40-10
H2O to
NH3
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• Molecules are close together, but not so
tightly packed they can’t move around.
• Unlike gases, liquids are held together
by attractive forces (intermolecular
forces – forces between molecules)
– This gives liquids a fixed volume
Nature of Liquids
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• Molecules are not moving fast enough to break attractive force and become a gas
• Not very compressible – too close together
Nature of Liquids
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Nature of Liquids
• Density
– Intermolecular force holds liquids together causing a
small volume – this means the density is higher than as
a gas
• Fluidity
– Molecules aren't so close together they can’t move, but
intermolecular forces make them less fluid than gases
– Viscosity – resistance of a liquid to flow
• Decreases with increase in temperature because the
motion allows the molecules to overcome
intermolecular forces
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Nature of Liquids
• Surface Tension
– Energy needed to
increase the surface area
of a liquid
– Caused by an imbalance
of intermolecular force
on the surface of a liquid
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Nature of Liquids
• Capillary Action – Cohesion – force of attraction between identical
molecules
– Adhesion – force of attraction between different molecules
• The adhesive force between water and glass is stronger than the cohesive force between water molecules
– Meniscus
– Capillary action
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• Tightly packed - held together by
stronger intermolecular forces than
liquids or gases
• Do not flow - molecules vibrate
around fixed points
• Incompressible
The Nature of Solids
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The Nature of Solids
• Density
– Generally the solid form of a
substance is the most dense
• ~10% more dense than liquids
• Packing tightly together decreases
volume
– H2O is the exception
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• Crystalline
– Most solids are crystalline
– Crystal - the atoms, ions or molecules are
arranged in a very orderly, repeating 3D
pattern known as the crystal lattice
The Nature of Solids –
Types of Solids
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The Nature of Solids –
Types of Solids
• Unit Cell
– Smallest group of particles that retains the
shape of the crystal.
• Allotropes
– 2 or more different molecular forms of the
same element in the same state.
• Diamond and graphite
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Types of Crystalline Solids
Type Unit Particle
Characteristics
of
Solid
Example
Atomic atoms •Soft
•Very low melting point
•Poor conductivity Noble gases
Molecular molecules •Fairly soft
•Low to medium-high mp.
•Poor conductivity I2, H2O, CO2
Covalent
Network
Atoms connected by
covalent bonds
•Very hard
•Very high mp.
•Poor conductivity
Diamond (C)
and Quartz
(SiO2)
Ionic Ions •Hard and brittle
•High mp.
•Poor conductivity NaCl
metallic Metal cations
surrounded by mobile
valence electrons
•Soft to hard and malleable
•Low to high mp.
•Excellent conductivity Any metal
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• Amorphous
– No ordered internal structure
– No definite melting point, melt over a
range or gradually “soften”
– Example Glass, rubber, and many
plastics
The Nature of Solids –
Types of Solids
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1. Melting point - temperature that a solid turns
into a liquid
– As the solid is heated, the molecules vibrate faster
until these vibrations break the intermolecular forces
that hold the solid structure together
– The stronger the intermolecular force the higher the
melting point
– Ionic compounds have stronger intermolecular
forces than molecular compounds
Phase Changes Requiring Energy
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Phase Changes Requiring Energy
2. Vaporization – process of a liquid changing to a gas or vapor
– Evaporation - Liquid to gas below boiling point
• Happens only on surface
• Only particles moving fast enough to overcome intermolecular forces escape to become a gas
• Cooling process - as high speed particles escape this lowers the average kinetic energy and therefore the temperature
• Heating liquid makes more high speed particles and rate of evaporation increases
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– Vapor pressure - vaporized particles will
exert an outward pressure • increasing temperature will increase vapor pressure
Phase Changes Requiring Energy
47
– Boiling Point - when vapor pressure equals
atmospheric pressure
• Unlike evaporation, boiling happens at a set
temperature and is not limited to the surface
• Temperature of liquid never gets above the
boiling point
Phase Changes Requiring Energy
48
3. Sublimation – change of a solid
directly to a gas
– Solid CO2 is called “dry ice” because it
sublimates from a solid to a gas without
becoming a liquid
– Ice sublimes and this is how “frost free”
refrigerators work
Phase Changes Requiring Energy
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Boiling Point Curve
0
20
40
60
80
100
120
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17
Tem
pe
ratu
re
(C)
Time (s)
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1. Condensation – gas or vapor becomes a
liquid
– slow moving gas particles that attract each
other and become a liquid
– high speed particles that enter the liquid
become trapped
2. Deposition – gas to solid (snowflakes)
3. Freezing – liquid to a solid
Phase Changes Releasing Energy
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What happens to a liquid in a closed container?
– At first high evaporation and low condensation
Phase Change Equilibrium
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• Closed Container Continued
– Over time evaporation rate will stay constant (stays at
constant temp in room) and condensation rate
increases as the number of gaseous particles increases
– Eventually the rates will become equal and dynamic
equilibrium is reached (particles are still traded but in
equal amounts so no net change)
Phase Change Equilibrium
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Phase Diagram
• Phase Diagram – diagram that shows
relationship between solid, liquid and
gas states for a substance in a closed
container at different temperatures and
pressures
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• Triple point - temp and pressure where all three states of matter co-exist in equilibrium
• Critical point – above this temperature, the substance can only exist as a gas
Phase Diagram
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Plasma: The Fourth State of Matter
• Gaseous mixture of electrons and
positive ions
• Generally occurs only at high
temperatures where the kinetic
energy is enough to strip electrons
off of gaseous atoms
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• Happens to a small degree (very
unstable and only lasts for a brief
moment in time) at lower
temperatures in fluorescent lights,
neon signs, lightning and the
“northern lights.”
Plasma: The Fourth State of Matter
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Graphite Diamond
Two
allotropes of
carbon
http://www.creative-chemistry.org.uk/molecules/carbon.htm