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pressure in the case of force applied to an area
P = F/A
1 Pa is equal to
1 N / m² is equal to
1 bar is equal to
10^5 Pa
1.00 atm is equal to
760 torr is equal to
values for STP
T - 273 K or 0 degrees C
P - 1.00 atm
760 torr or 1 atm is equal to
101.325 kPa is equal to
Boyles’s Law
P = c/v to P1 V1 = P2 V2
Charles’ Law
c = V/T to V1 / T1 = V2 / T2
Avogadro’s Law
v= cN to V1 / n1 = V2 / n2
ideal gas equation
PV=nRT
ideal gas assumptions
- no attractive forces
- neglect particle volume
- there is identical molar volume at STP no matter the gas ID
- the gas consists of individual point particles
density of ideal gas formula
d= PM / RT
Dalton’s Law of Partial Pressures
P1 = n1 / nT times Pt
Dalton’s Law of Partial Pressures important notice
the total pressure of a mixture of gases equals the sum of the partial pressures that each would exert if present alone
kinetic molecular theory
1) gases consists of large numbers of molecules that are in continuous, random motion
2) the combined volume of all the molecules of the gas is negligible relative to the total volume in which the gas is contained
3) no attractive forces
4) average kinetic energy depends only on absolute temp
effusion
the escape of gas molecules through a tiny hole into an evacuated space
diffusion
the spread of one substance throughout a space or throughout a second substance
deviations from ideal gas behavior
assumptions made in KE molecular theory don’t hold true if there is high pressure and / or low temperature
real gases behavior
most ideal behavior is a higher temperature (1000K) and the molecules have lower average KE so the attractive forces can take over
van der Waals equation
(P+n²a/v²) (V-nb) = nRT
