chapter 13 states of matter. 13.1 gases the kinetic-molecular theory describes the behavior of gases...
TRANSCRIPT
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Chapter 13
States of Matter
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13.1 Gases
The Kinetic-Molecular Theory
Describes the behavior of gases in terms of particles in motion
Size Motion Energy
Gas Behavior
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Particle Size
Particle size is small relative to the space that surrounds them
The distance between particles is so large that no attractive or repulsive forces exist between gas particles
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Particle Motion
Gas particles are in constant, random motion Particles move in a straight line until they
collide with another particle or the walls of a container
Gas particle collisions are elastic. Although particles in collision can transfer kinetic to eachother, the total kinetic energy of the colliding particles remains constant.
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Particle Energy
All particles in a gas have the same mass, but not the same kinetic energy.
Temperature is the measure of the average kinetic energy of particles in a sample of matter
2
2
1mvKE
Mass
Velocity
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Explaining the Behavior of GasesLow DensityParticle size is small relative to the space that surrounds them As compared to liquids or solids, gases have much smaller densities due to the fact that fewer particles occupy the same volume
Compression and ExpansionParticle size is small relative to the space that surrounds them
The large amount of space surrounding the particles allows particles the room to move closer together as they are compressed
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Diffusion and Effusion
Gas particles are in constant, random motionGas particles tend to move from a high area of concentration to a low area of concentration. The rate at which they diffuse is dependent on their mass.
B
A
MolarMass
MolarMass
RateB
RateAComparison of the diffusion rates of
two gases
Graham’s Law of Effusion
Gas escaping through a small hole
MolarMassteEffusionRa
1
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Practice Problem
What is the ratio of effusion rates for Nitrogen (N2) and Neon (Ne)?
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By: Jordan Gaffin, Sam Bear, and Keith Zubrow
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What is Pressure?
Pressure is defined as force applied
per unit area. We measure air and
atmospheric pressure with a
barometer.
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Measuring Air Pressure
Italian physicist Evangelista Torricelli was the first to demonstrate that air exerted pressure.
He invented a device called the barometer that assisted him in measuring pressure.
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Gas Pressure
A manometer is an instrument used to measure gas pressure in a closed container.
It is a flask that is connected to a U-tube that contains mercury.
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Unit Compared with 1 atm
Compared with 1 kPa
Kilopascal (kPa) 1 atm = 101.3 kPa
Millimeters of mercury (mm Hg)
1 atm = 760 mm Hg I kPa = 7.501 mm Hg
Torr 1 atm = 760 torr 1 kPa = 7.501 torr
Pound per square inch (psi or lb/in)
1 atm = 14.7 psi 1 kPa = .145 psi
Atmosphere (atm) I kPa = .009869 atm
Units Of Pressure
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Units of Pressure
The SI unit of pressure is the pascal (Pa). 1 pascal = the force of 1 Newton per square
meter. The pressures measured by barometers and
manometers can be reported in millimeters of mercury (mm Hg).
There is also a unit called the torr which is equal to 1mm Hg.
Often air pressure is reported in a unit called atmosphere (atm).
1 atm = 760mm Hg
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Conversions
To convert the different measures of pressure, we use factor label.
Use the given measurements: 1 atm= 101.3 kPa = 14.7 psi = 760 torr
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Dalton’s Law of Partial Pressures
This law states that the total pressure of a mixture of gases is equal to the sum of the pressures of all the gases in the mixture.
Dalton’s law of partial pressures can be summarized as:
Ptotal = P1+P2+P3+... Pn
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The P’s in Daltons Theory
Ptotal is the total pressure of a mixture of gas.
All the other P’s (P1, P2, P3, Pn…) are the partial pressures of each gas in the mixture.
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Example
A mixture of Oxygen, Carbon Dioxide, and Nitrogen has a total pressure of .97 atm. What is the partial pressure of O2 if the partial pressure of CO2 is .70 and the partial pressure of N2 is .12?
P.97atm = P N2 .12 atm + P CO2 .7atm + P O2 x
P O2 = .97 atm - .70 atm - .12 atm
P O2 = .15 atm
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Section 13.2Forces of Attraction
By Dorothy Raginsky, Rima Naseer, Luke Morreale, Emil George, Greg Klein
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Table of Contents
Differences between intramolecular forces and intermolecular forces
Dispersion Forces with examplesDipole-dipole Forces with examplesHydrogen Bonds with examples
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Intramolecular Forces
The attractive forces that hold particles together in ionic, covalent, and metallic bonds.
Intra- means “within”.
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Intermolecular Forces
They hold together identical particles or two different types of particles.
There are three types of intermolecular forces: Dispersion forces Dipole-dipole forces Hydrogen bonds
Some intermolecular forces are stronger than others, but all are weaker than intramolecular forces.
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Dispersion Forces
They are weak forces that result from temporary shifts in the density of electrons in electron clouds.
Dispersion forces are sometimes called London forces after the German American physicist who first described them, Fritz London.
They are the weakest intermolecular force. They are the dominant force of attraction
between identical, nonpolar molecules.
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Examples
Fluorine, chlorine, bromine, and iodine exist as diatomic molecules.
Recall that the number of nonvalence electrons increases from fluorine to chlorine to bromine to iodine. Because the larger halogen molecules have more electrons, there can be a greater difference between the positive and negative regions of their temporary dipoles, and, thus, the stronger dispersion forces. This difference in dispersion forces explains why chlorine and fluorine are gases, bromine is a liquid, and iodine is a solid at room temperature.
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Dispersion Forces
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Dipole-dipole Forces
The attraction between oppositely charged regions of polar molecules are called dipole-dipole forces.
Polar molecules contain permanent dipoles; that is, some polar regions of a polar molecule are always partially negative and some regions are always partially positive.
Neighboring polar molecules orient themselves so that oppositely charged regions line up.
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Examples
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.
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Dipole-dipole Forces
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Hydrogen Bonds
A hydrogen bond is one special type of dipole-dipole attraction that occurs between molecules containing a hydrogen atom bonded to a small highly electronegative atom with at least one lone electron pair.
For a hydrogen bond to form, hydrogen must be bonded to either a fluorine, oxygen, or nitrogen atom. These atoms are electronegative.
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Examples
In a water molecule, the hydrogen atoms have a large partial positive charge and the oxygen atom has a large partial negative charge. When the water molecules approach, a hydrogen atom in one molecule is attracted to the oxygen atom on the other molecule.
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Hydrogen Bonding
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Section 13.3 Liquids
Ryan Desch, Dave Derr, Colin Drummond,Kyle Giordano, Jake
Long, Pratt Templeton
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Density and Compression
At one atmosphere of air pressure liquids are more denser than gases.
Change in volume for liquids is much smaller than gas because liquid particles are already tightly packed.
An enormous amount of pressure must be applied to reduce the volume of liquid by even a few percent.
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Fluidity
Fluidity- The ability to flow
A liquid is considered as fluid because it can flow.
Liquids can defuse through other liquids due to their fluidity.
Liquids are less fluid than gases.
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Viscosity Viscosity- A measure of the
resistance of a liquid to flow The particles in liquid are
close enough for attractive forces to slow their movement.
Viscosity decreases with temperature (Increase in temperature = increases the average kinetic energy of molecules).
For Example- When you pour a tablespoon of cooking oil into a frying pan the oil tends not to spread across the bottom of the pan until you heat the oil.
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Surface Tension Surface Tension- The
energy required to increase the surface area of a liquid by a given amount.
The stronger the attractions between particles, the greater the surface tension.
Surfactants- Compounds that lower the surface tension of water.
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Capillary Action When liquid is placed in a narrow container the surface forms a concave meniscus (The surface dips in the center).
This movement of the liquid, such as water, is called capillary action or capillarity.
Cohesion- The force of attraction between identical molecules
Adhesion- The force of attraction between molecules that are different.
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Solids
SolidsJoe BruchBria CollinsMelanie SpivackSiobhan BradyLaura Betterly
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Density
Solids are more dense than liquids and gases All the solids remain a compact phase, where
the arrangement of molecules is very ordered and well defined. With this in mind it is surprising that there exists some solids (few) less dense than liquids.
P=m( mass)/v (volume). This is the equation for density
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Crystalline Solids Are arranged in fixed geometric patterns or lattices. Ex. ice methanol and sodium chloride They have an orderly arranged units and are practically
incompressible Crystalline solids, also, show a definite melting point They passed rather sharply from solid to liquid state There are various crystalline forms which are divided into
six crystal systems or shapes: cubic, tetragonal, hexagonal, rhombic, monoclinic, and triclinic
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Molecular Solids
a solid composed of molecules held together by relatively weak intermolecular forces
are soft, and generally have low melting and boiling temperatures
Most solids are nonconducting when pure and are insoluble in water, but soluble in non-polar solvents
Ex. Sulfur, ice, sucrose, and carbon dioxide solid
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Metallic Solids
More than 80 elements in the periodic table are metals
High densities Large atomic radii Low ionization energies Low electronegativities Usually, high deformation Malleable Ductile Thermal conductors
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Covalent Network of Solids
-Have no discrete molecular units
-Held together by conventional covalent bonds
-Continuous network of bonded atoms
examples – diamond and quartz
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Amorphous solids
is a solid in which there is no long-range order in position- molecules are arranged in a random manner
common window glass and cotton candy and plastic are examples
Amorphous solids also unlike crystalline solids. They do not have definite melting points
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Phase ChangesBy, Taylor Brink, Lindsay
Coloracci, Amanda Couch, Emily Rorer, Callie Wendell, and Jeff
Wright
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Endothermic Changes (within)
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Melting
Hot flows into cold Heat disrupts the hydrogen bonds holding
the ice together These bonds are strong, therefore, lots of
energy is required Melting Point- temperature at which a
solid becomes a liquid- hard to determine exact point
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Vaporization vs. Evaporation
Vaporization liquid changes into a gas (vapor) the gradual input of energy equals the amount of
molecules escaping from the surface Vapor pressure- pressure exerted by a vapor over a
liquid Surface molecules are less attracted to other molecules
Evaporation molecules escaping from the surface better in warmer temperatures because there is more
energy all evaporates over time depending on the amount of
water and energy
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Boiling Process
Boiling point- temperature where vapor pressure equals the atmospheric pressure
At this temperature, all molecules have enough energy to vaporize, where as only the surface molecules could before
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Sublimation
Transition from a solid to a gas, skipping a liquid examples:
• Carbon Dioxide (dry ice)• Air fresheners• Ice in a freezer
Happens when the pressure is low
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Exothermic Changes(outside)
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Condensation
Transition from gas to liquid Hydrogen bonds begin to form, and
the vapor looses energy examples
• Sides of a cold glass• Dew• Fog• Rain
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Deposition
Transition from gas to solid Vapor settles on to a solid
Examples • Frost• Snowflakes
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Freezing
Transition from liquid to solid Water looses energy (heat) and
hydrogen bonds form, stopping the motion of water molecules
Freezing point- temperature where liquid turns to solid
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