Atmospheric Pressure:

• Atmospheric pressure is the force exerted by gas molecules in the air as they strike surfaces.

• Measured using a barometer; standard atmospheric pressure is 760 mm Hg (1 atm = 760 mm Hg = 101,325 Pa = 14.7 psi).

• Factors such as altitude and weather can affect barometric pressure.

Units of Pressure:

• Common units: atmospheres (atm), millimeters of mercury (mm Hg or torr), pascals (Pa), and pounds per square inch (psi).

• Conversion between units is essential for calculations.

Gas Laws:

• Boyle's Law: At constant temperature, pressure and volume are inversely proportional. PV=k or P1V1=P2V2

• Charles' Law: At constant pressure, volume and temperature (in Kelvin) are directly proportional. VT=k or V1T1=V2T2

• Avogadro's Law: At constant temperature and pressure, volume and moles of gas are directly proportional. Vn=k or V1n1=V2n2

• Combined Gas Law: Relates pressure, volume, and temperature for a fixed amount of gas. P1V1T1=P2V2T2

• Ideal Gas Law: Combines all the above laws into one equation. PV=nRT where R is the universal gas constant (0.0821 L·atm/(mol·K)).

Problem Solving with Gas Laws:

• Identify which law applies based on what variables are held constant.

• Convert all units to the appropriate SI units before solving (e.g., temperature to Kelvin, pressure to atm, volume to liters).

• Use the correct equation to solve for the unknown variable.

• For the ideal gas law, ensure all variables are in the correct units to match the value of R.

Conceptual Understanding:

• Pressure results from gas molecules colliding with surfaces.

• Volume, pressure, temperature, and amount of gas are interrelated; changing one affects the others according to the gas laws.

• Real gases may deviate from ideal behavior under high pressure or low temperature.