States of Matter
Chapter 12
It is easy to distinguish between solids and liquids by looking at their physical properties...BUT it is very difficult to distinguish gases from one another.
James Maxwell and Ludwig Boltzmann proposed the kinetic molecular theory which describes the behavior of matter in terms of particles in motion.
(Refresh: kinetic energy is energy in motion, potential energy is stored energy.)
*Several assumptions were made by the theory:
1. Particle Size­ gases consists of small particles that are separated from one another by empty space. Because they are far apart, they experience no significant attractive or
repulsive forces.
2. Particle Motion­ gas particles are in constant random motion and move in a straight line until they collide with other particles or the walls of their container.
­Collisions between gas particles are elastic­these are collisions where no kinetic energy is lost, just transferred between particles.
3. Particle Energy­ Mass and velocity determine the kinetic energy KE = 1/2 mv2
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KE = 1/2 mv2
KE = kinetic energy
m = mass
v = velocity, both speed and direction of motion
Temperature­ a measure of the average KE of the particles in a sample of matter.
How Gases Behave
*Low Density (m/v) ­this is because they take up more space *Compression & expansion ­The large volume of gas particles allows the particles to be compressed or expanded if given additional space. ­Example: if you squeeze a foam pillow you can compress it, if you release it, it will expand again. *Diffusion & effusion ­gas particles can easily flow past one another because of their low attraction to one another. Gases are often entering areas already occupied by another gas. ­The random movement of particles causes the gases to mix until they are evenly distributed.
­Diffusion is the term used to describe the movement of one material through another from areas of high concentration to areas of low concentration.
­Effusion ­ the term used to describe an escapting gas through a tiney oepning such as a balloon with a pin hole in it.
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Graham's law of effusion
­Graham conducted experiments to measure rates of effusion for different gases at the same temperatures
Graham's law of effusion states that the rate of effusion for a gas is inversely proportional to the square root of its molar mass.
Rate of effusion *The rate depends mainly on the mass of the particles involved.
­Lighter particles will diffuse more quickly
Recall: Gases at the same temperature have the same KE but masses vary from gas to gas so the velocitys must also change (mass and velocity are inversely proportionate)
Using Grahm's law, we can set up a proportion to compare the diffusion rates for 2 gases.
Try Using it! Practice Problems 1 & 2 on page 405
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Pressure ­ the force/ unit area
­Example: The bottom of a boot is much larger than the area of a high­heeled shoe. This means the overall pressure on the soft surface of snow is less with a boot than with a heel. Moral of the story: wear boots not heels in the snow!
Earth is surrounded by an atmosphere that contains particles that move in every direction thus they cause pressure in every direction. This is atmospheric pressure or air pressure. It varies from place to place on Earth.
­Because gravity is greater at the surface­ there are more particles here than in higher altitudes as well as more pressure.
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Evangelista Torricelli­ the first to demonstrate that air exerted pressure.
­He invented the barometer­ an instrument that is used to messure atm pressure.
*Atmospheric pressure at sea level usually has mercury up to a height of about 760mm.
­exact height is determined by 2 forces: gravity pushing downward and air pressure pushing upward.
­Changes in air temperature or humidity can cause air pressure to vary greatly.
Manometer­ an instrument used to measure gas pressure in a closed container.
­The difference in the height of the mercury in the two arms is used to calculate the pressure of the gas in the flask.
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Units of Pressure
­SI unit is the Pascal (Pa), it is derived from the SI unit of force called the Newton (N). ­1 pascal = to a force of one N/m2
­many fields of science use more traditional units of pressure
­engineers use pounds/square inch (psi)
­barometers and manometers use mm Hg
­torr and bar are also used
­At sea level the average air pressure is 101.3 kPa when the temperature is 0oC
­Air pressure is often reported in atmospheres (atm)
­1atm = 760 torr or 760 mm Hg or 101.3 kPa
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Dalton's law of partial pressures
Dalton found that each gas in a mixture exerts pressure independently of other gases that are present.
­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.
PTotal = P1 + P2 + P3 + ....Pn
­The portion of the total pressure contributed by a single gas is its partial pressure.
­depends on the number of moles of gas, size of the container, and the temperature of the mixture.
­Homework pg 410: #8­10, 12­13
­Conversion Worksheet
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12.2 Forces of Attraction
­The attractive forces that hold particles together in ionic, covalent, and metallic bonds are called intramolecular forces.
­intra means within
­Other forces called intermolecular forces (inter meaning between or among) 3 types: dispersion forces, dipole­dipole forces, and hydrogen bonds
­in general intermolecular are weaker than intramolecular
1. Dispersion forces (aka London forces) ­ weak forces that result from temporary shifts in the density of electrons in electron clouds.
Remember: electrons in the cloud are in constant motion if they collide with another electron from a different cloud they repel. The electron density around each nucleus is for a moment greater in one region of the cloud and forms a temporary dipole. If they get closer enough dispersion forces exists
These exist between all particles, they become stronger as the size of the particles increase (more electrons means a greater difference between the positive and negative regions of their temporary dipoles­stonger dispersion.)
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2. Dipole­dipole forces ­attractions between oppositely charged regions of polar molecules.
Remember­ polar molecules contain permanent dipoles (parts are always partially positive and partially negative)
­Small polar molecules with large dipoles have stronger forces than dispersion BUT in many polar molecules dispersion forces dominate dipole­dipole forces.
3. Hydrogen Bonds ­ a special type of dipole­dipole that occurs between molecules containing a hydrogen atom bonded to a small, highly electronegative atom with at least one lone pair of electrons.
­must be bonded with F, O, N
­dominates the other two forces in strength
­The presence of hydrogen bonds explain why water is a liquid at room temperature, while compounds of comparable masses are gases.
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­Study Guides 12.1 & 12.2, ­page 414: #14­16
12.3 Liquids and Solids
­Kinetic molecular theory also applies to liquids and solids­ you must consider the forces of attraction between the particles and their energy of motion.
Properties of Liquids:
­take the shape of their container
­incompressible
­more dense than gases
­have a fixed volume­can't expand to fill their container
*According to the Kinetic molecular theory individual particles do not have fixed positions but the attraction between particles limit their range of motion so they remain closely packed in a fixed volume.
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1. Fluidity ­ gases and liquids are classified as fluids
­liquids diffuse more slowly than gasses because intermolecular forces interfere with the flow
2. Viscosity ­ a measure of the resistance of liquid to flow
*Ex: honey and syrup have higher viscosity than water
*The viscosity of the liquid is determined by the type of intermolecular forces in the liquid, the size and shape of the particles, and the temperature.
*The stronger the intermolecular forces, the higher the viscosity
*The larger the particle size the higher the viscosity
­Longer chains of molecules have higher viscosity than smaller molecules, assuming attractive forces are the same.
*Temperature ­viscosity decreases with an increase in temperature.
­The added heat energy makes it easier for the molecules to overcome the intermolecular forces that keep the molecules from flowing.
3. Surface Tension: particles at the surface have no attractions above them, unlike the attractions between the particles below the surface.
­surface tends to have the smallest possible area and acts as though it is stretched tight
­surface tension­ is 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.
*Why are drops of water round? the surface area of a sphere is smaller than the surface area of any other shape with a similar volume.
*Soaps and detergents decrease the surface tension of water by disrupting the hydrogen bonds between the water molecules
­Surfactants­ compounds that lower the surface tension of water
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4. Cohesion and Adhesion
*cohesion­ describes the force of attraction between identical molecules
*adhesion­ the force of attraction between molecules that are different
Because adhesive forces are stronger, water rises along the inside walls
5. Capillary Action­ the tendency of water to be drawn upwards in thin tubes.
­Explain's a paper towel's absorbancy & trees ability to get water to the branches
SOLIDS ­Properties
­definite shape & volume
­Kinetic Molecular Theory states that a mole of solid has the same energy as a mole of liquid at the same temperature.
­There are strong attractive forces between the particles in the solid, limiting the motion to vibrations around fixed locations in the solid
­Not fluid
1. Density ­ more dense than most liquids.
­Majority of the time if both the solid and liquid form are present the solid will sink in the liquid
­Water is unique. It is less dense as a solid. As water freezes the molecules form hydrogen bonds, making them less
closely packed.
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2. Crystalline solid­ a solid whose atoms, ions, or molecules are arranged in an orderly, geometric structure. The locations of particles in a crystalline solid can be represented as points on a framework called a crystal lattice.
Unit cell­ smallest arrangement of atoms in a crystal lattice that has the same symmetry as the whole crystal (building block determines the shape)
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Categories of crystalline Solids
*based on types of particles and how they are bonded together
1. atoms­ only noble gases
2. molecular solids ­ held together by intermolecular forces
­most are not solids at room temperature, only if they have large molecular masses
­poor conductors of heat and electricity.
3. Covalent network solids ­ones that can form multiple covalent bonds. ­allotrope­ an element that can exist in different forms in the same state.Ex: carbon can be diamond or graphite in the solid state.
4. Ionic ­ ions are always surrounded by ions of opporsite charges. Type and ratio determine the structure and shape of the crystals
5. Metallic ­ have wide variety of physical properties due to large variance in bonds between cations and electrons.
*Know the characteristics of each & be able to give an example.
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Amorphous solid­ one in which the particles are not arranged in a regular repeating pattern.
­no crystals
­often forms when a molten material cools too quickly ­Examples include: glass, rubber,and many plastics
Homework page 424: #18­25
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12.4 Phase Changes
Review:
*Exothermic­ releasing energy
*Endothermic­ taking energy in
6 Different Changes in State
start from:
change to:
Name:
solid
liquid
melting
liquid
solid
freezing
liquid
gas
evaporation
gas
solid
gas
liquid
gas (skipping liquid phase)
solid (skipping liquid phase)
condensation
sublimation
deposition
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Fill in the chart, Be sure to include whether the process is exothermic or endothermic.
Time
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­Phase change diagram practice
­Yellow phase change diagram sheet (FOR A GRADE)
­Online Lab
­When energy is added or removed one phase can change into another
­Two variables combine to control the phase of a substance: Temperature & Pressure
­As Temp goes up, pressure goes down
Phase Diagram: a graph of pressure verses temperature that shows in which phase a substance exists under different conditions
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Point A is the triple point. This is the point on the graph where all lines meet. This represents the temperature and pressure at which all 3 phases of a substance coexist.
­All 6 phase changes can occur at the triple point.
Point B is the critical point. This indicates the critical temperature and pressure above which water can not exist as a liquid. *Phase Diagrams are different for each substance because normal boiling and freezing points of substances are different.
Classwork page 430: 27­33
Transparency Worksheet pg 92­93 & 100­101
Lab finding triple point
Review
Test
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