Chapter story: "the bends"—resulted from "un-dissolved" gas that is formed when pressure around the body is released too quickly (which is similar to when a soda can is opened in the air). This property is described by the Henry’s law, cf. 5.11.

Concept not covered by the book, but important
A) Kinetic-molecular theory: the atoms, ions, or molecules of gases, liquids, and solids are in constant motion and that there is no attraction between gas particles.

B) Gas particles
• High speed, ~1,600 km/h at 20 °C (the higher the temp the higher the speed)
• Traveling in straight lines
• Changing directions by collisions
• Short mean free path (short unimpeded path of travel)
• Overall, an aimless path—a random walk

C) Kinetic energy: energy from motion
• Particles of all substances at the same temperature have the same average kinetic energy.
• Temperature is a measure of the average kinetic energy; increase temperature, increasing kinetic energy.
• Kinetic energy of particles has a Gaussian distribution.

Pressure: measure of force on a unit area, i.e., P = force/area.
Rationalize why a sharp object can pierce through the skin easily.

Gas pressure: result of simultaneous collision of gas particles on an object. Atmospheric pressure is an example.
High pressure system— cold air (cold front)
Low pressure system— tropical depression and hurricane
Differences in pressure cause the movement of gases from one place to another.

Atmosphere (atm): a unit of pressure, defined as the pressure required to support 760 mm of Hg (the unit for blood pressure).

The space of the vacuum is larger if the tube is longer.

Other units: 1 torr = 1 mm Hg = 1/760 atm
pascal (Pa): 1 atm = 101.3 Pa
    How many Pa is for 1 mm Hg?
    (1 mm Hg ´ 1/760 atm/mm Hg ´ 101.3 Pa/atm = ______ Pa)

The Unit psi (pound per squared inch) has also been very often used, particularly on the gauge of gas tanks like the one on P. 124.
1 psi = 0.068 atm
      How many atm is the oxygen pressure in a tank of 2000 psi?
     2000 psi ´ 0.068 atm/psi = _____ atm

The Gas Laws: Behavior of gases
—effect of pressure, volume, and temperature on gas
Review yourselves the concepts about "proportionality" and "inverse proportionality", and find some examples!

Effect of adding gas (to a constant volume)
Examples: pumping a tire, a football,...
Result: The more gas is pumped into a fixed volume, the higher the pressure, i.e., pressure is proportional to the number of gas molecules,
i.e., P = constant ´ n (P: pressure, n: number of moles)

What about if the volume is not fixed, and gas is pumped into the volume?
(Volume increases, i.e., V = const. ´ n.)

Effect of changing the size of the container (with a fixed amount of gas molecules)
Examples: getting medicine into a syringe, inhaling air by the lungs
Result: Drug solution is pushed into syringe. The pressure in the syringe is reduced, and the solution is pushed into the syringe by the air, i.e., with a fixed amount of gas, pressure is inversely proportional to volume, i.e., P = const. ´ 1/V.

P and V relationship at constant temperature and constant amount of gas (Boyle's law): The volume of a gas is inversely proportional to its pressure.
i.e., = or P₁ x V₁ = P₂ x V₂

Effect of heating or cooling a gas
Examples: a basketball (or balloon) in the sun
Result: it becomes harder (or bigger)

basketball case—volume is more or less fixed; the higher the temperature, the higher the pressure, i.e., P = const. ´ T.
T and P relationship at constant volume and amount of gas (Gay-Lussac’s Law): Pressure is proportional to its temperature in K.
i.e., = P₁/P₂ = T₁/T₂

balloon case—pressure is more or less fixed (slightly higher than 1 atm); the higher the temperature, the larger the volume,
i.e., V = const. ´ T
T and V relationship at constant pressure and amount of gas (Charles' law): Volume of a gas is  proportional to its temperature in K.
i.e., = V₁/V₂ = T₁/T₂

Absolute zero—the temperature at which particles stop moving and an idea gas has no volume.

Ideal gas: would not liquefy and solidify as cooled, have no volume at 0 K, and obey the gas law (discussed later) at all temperatures and pressures.

Avogadro's law: Equal volumes of gases at the same temperature and pressure contain equal number of particles.

i.e., V = const ´ n
At Standard Temperature and Pressure (STP), i.e., 273 K and 1.00 atm, 1 mol (i.e., 6.02 ´ 10²³ particles) of any gas (actually ideal gas) occupies 22.4 L, which is the molar volume of gases.
For example (Table 5.1): H₂, 22.43 L; He, 22.42 L; N₂, 22.38 L

"Combined gas law" (of the above relationships) at constant amount of gas:
P₁V₁/T₁=P₂V₂/T₂

P and V are  proportional to the mount of gas

i.e., P₁V₁/nT₁ = P₂V₂/nT₂

The ideal gas law

1 mol of an ideal gas occupies 22.4 L at STP

i.e., PV/nT = (1 atm ´ 22.4 L)/(1 mol  273 K) =  0.0821 (atm L/mol K)

The constant is called the ideal gas constant (R).

The ideal gas law is P x V = n x R x T (PV = nRT)

(You need to know how to obtain n, the number of moles, e.g., 32 g oxygen gas is 1.0 moles, 28 g nitrogen gas is 1.0 mole, etc.)

Dalton's law of partial pressures: at constant volume and temperature, the total pressure exerted by a mixture of gases is equal to the sum of the partial pressures by the different gases,

i.e., P_total = P₁ + P₂ + P₃ +...