1. States of Matter

2. The Kinetic-Molecular Theory
Describes the behavior of gases in terms of particles in motion.
Gases are composed of small particles separated from one another by empty space. This spacing means no significant IMF exist among particles.
Gas particles are in constant, random motion.
Temperature is a measure of the average kinetic energy of particles in matter.

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4. Particles collide
with one another
and the sides of
the container.
Energy can be
transferred upon
collision, but the
net result is energy
is conserved =
elastic collisions.

5. Properties of Gases Low density:
A great deal of space exists between gas
particles. Since density is a measure of
m/V, for a given mass of a gas, the
density will be thousands of times lower
than a comparable mass of a solid.

6. Compression and expansion:
The large amount of space that exists
between gas particles allows them to
be easily pushed into a smaller volume.
Diffusion and effusion:
The random motion of gas particles
causes gases to mix until they are
evenly distributed. This process is
called diffusion and depends on the
mass of the particles involved. The
lighter the particle, the faster it moves.

9. Effusion is the escape of a gas through a tiny opening.
Experiments to measure the rates of effusion of different gases indicated an inverse relationship between the rate of effusion and the mass of the gas.
In other words, the lighter the gas, the
faster the particles will move.
r1 = molar mass2
r2 molar mass1

10. Calculate the ratio of the effusion rates for
nitrogen gas and neon.
Calculate the molar mass of butane.
Butane’s rate of diffusion is 3.8 times
slower than that of helium.

11. Gas Pressure Force per unit area.
Gas particles exert pressure when they collide with the sides of their container.
Atmospheric pressure is the force exerted by the particles surrounding the earth.
Air pressure varies with your position on earth. At the top of a mountain, the atmospheric pressure is less than that at sea level since the column of air pressing down on you is less.

12. Measuring Pressure A barometer measures atmospheric pressure.
The height of the mercury in the column is generally 760 mm.
The forces exerted on the column of mercury are the force of gravity downward and the upward force due to the air.

13. A manometer is a device used to measure gas
pressure in a closed container.

14. Units of Pressure SI unit is the pascal (Pa). 1 Pa = 1 N/m2
Common units of pressure are the psi (pounds per square inch) used to measure tire pressure and the atmosphere (atm)
1 atm = 760 mm Hg = kPa

15. Perform the following conversions:
3.5 atm to mm Hg.
450 mm Hg to kPa.
99.5 kPa to atm.

16. Dalton’s Law of Partial Pressures
Examine the flask to the left.
What gas particles are present in the flask?
Dalton’s law states that the total pressure of the gas in the mixture is the sum of all the pressures.

17. Gases “collected over water”
The best application of Dalton’s law to lab
situations is when a gas is collected by bubbling the gas into a bottle containing water.
The pressure inside the flask is the sum of the
pressures of the collected gas and water vapor.
Pdry gas = Patm – Pwater vapor P of the water vapor is T dependent.

18. Forces of Attraction
The state of a substance depends on the forces of attraction within and between its particles.
Intramolecular attractive forces are the ionic, covalent or metallic bonds within the substance.
Intermolecular attractive forces exist between particles of the same or different substances.

19. London Dispersion Forces
Dispersion forces are weak forces caused by temporary shifts in the density of electrons surrounding the nucleus. The strength of dispersion forces is directly related to the number of electrons that a substance has.
Weakest of the three IMF and are significant primarily in large nonpolar molecules.

21. Dipole-Dipole Forces
Attractive forces between oppositely charged regions of polar molecules.
Stronger than dispersion forces but molecules must have comparable mass.

22. Hydrogen Bonding Strongest of the 3 IMF
Occurs when hydrogen is bonded to a small, highly electronegative atom.
Exists when H is bonded to F, O or N.

23. The properties of water, a molecule essential
for sustaining life are due to hydrogen bonding.
The low density of ice as compared to liquid
water, the existence of water as a liquid at
room temperature, the surface tension of water,
and the high specific heat of water are all
properties directly related to hydrogen bonding.

26. Liquids KM theory applies to liquids when IMF are considered.
Liquids assume the
shape of their
container and flow.

27. Liquids can be compressed slightly
Properties of Liquids
Density and compression:
Liquids are more dense than their
vapor. Higher density due to the
IMF holding the particles together.
Liquids can be compressed slightly
but a great deal of pressure must be
applied.

28. Fluidity
Gases and liquids are able to flow.
Liquids can diffuse through each
other, but at a rate more slowly
than gases.
Viscosity
A measure of the resistance of
a liquid to flow.

29. Surface Tension
Defined as the energy required to
increase the surface area of a
liquid by a given amount.
Capillary Action
Adhesive and cohesive forces
account for the behavior of
liquids in small tubes called
capillary tubes.

31. Forces of attraction
between the water
and the glass are
greater than the
forces between
water molecules, so
the water rises
along the glass.

32. Solids Strong attractive forces exist between solid particles.
More order exists in
solids than liquids; as a result, solids
do not flow.

33. Particles of a solid are arranged in a particular way.
Properties of Solids
Density
Since particles are more closely packed, the density of solids is greater than those of liquids.
Crystallinity
Particles of a solid are arranged in a particular way.

34. Types of Unit Cells
Simple Cubic
Body-centered Cubic
Face-centered Cubic

35. Types of Crystalline Solids
Molecular solids: held together by dispersion forces; fairly soft, low melting points.
Covalent Network solids: very hard solids with high melting points.
Ionic solids: positive and negative ions held together by electrostatic forces of attraction; hard, brittle, high melting points.

36. Types of Crystalline Solids
Metallic solids: held together by metallic bonds; electrons are free to move around nuclei; soft to very hard; low to very high melting point; excellent thermal and electrical conductors; malleable and ductile.

37. List the names for the following phase transitions:
solid  liquid
liquid  gas
solid  gas
liquid  solid
gas  liquid
gas solid

38. Endothermic Phase Changes
Endothermic phase changes require the input of energy.
Melting: solid  liquid.
Vaporizing: liquid  gas.
Subliming: solid  gas.

39. Melting point: temperature at which the
forces holding the crystal together are
broken and it becomes a liquid.
Evaporation: vaporization occurring
only at the surface of a liquid.

40. Vapor Pressure Pressure exerted by a vapor over a liquid.
The temperature at which the vapor pressure of the liquid equals atmospheric pressure is called the boiling point.

41. Exothermic Phase Changes
Exothermic phase changes release energy.
Condensation: gas  liquid.
Deposition: gas  solid.
Freezing: liquid  solid.

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