- 117 Views
- Uploaded on
- Presentation posted in: General

5.1 Pressure

Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author.While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server.

- - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - -

- Pressure
a. Devices used to measure pressure

1. Barometer: measures atmospheric pressure

*Atmospheric pressure results from the mass of air being pulled toward the center of the earth by gravity.

2. Manometer: measures the pressure of a gas in a container

b. Units of Pressure

1. mmHg = Torr (invented the first barometer)

2. 1 atm = 760 mmHg = 760 torr = 1.013 x 105 Pa

3. Pressure = force/ area

- Gas Laws
*pressure, temp, vol, amount

A. Boyle’s Law: volume and pressure are inversely related at constant temp

PV = k

P1V1 = P2V2

B. Charles’ Law: Kelvin temperature and volume vary directly with each other at constant temperature.

*if you graph a line of V vs. T all lines extrapolate to zero volume at 0K.

*V1/ T1 = V2/ T2

C. Avogadro’s Law: for a gas at constant temperature and pressure, the volume is directly proportional to the number of moles of gas present

* V1/ n1 = V2/ n2

- Ideal Gas Law
A. PV = nRT, R = 0.08206 Latm/ Kmol

B. Gas laws describe ideal gases, not real gases

C. Real gas behavior approaches ideal at high temps and low pressures

- Gas Stoichiometry
A. The moalr volume of a gas at STP is 22.4 L ideally

B. STP 0oC, 1atm (273 K, 1.013 x 105Pa)

- CH4 (g) 2.80L @25oC 1.65 atm
- O2(g) 35.0 L @31oC 1.25 atm
CH4(g) + 2O2(g) CO2(g) + 2H2O (g)

CO2(g) __L @ 125oC 2.50 atm

C. n = grams of gas/ molar mass

D = m/v

P = dRT/ molar mass

- Dalton: for a mixture of gases in a container, the total pressure exerted is the sum of the pressures that each gas would exert if it were alone.
A. Ptotal = P1 + P2 + P3 + …

B. Calculate the partial pressure using the ideal gas law

Ptotal = (n1 + n2…) RT/V

C. Total pressure depends on the total number of moles of gas, not the identity of the particles

D. Mole fraction is the ratio of the number of moles of a given component in a mixture to the total number of moles in the mixture (x = chi)

x = n2/ ntotal = P2/ Ptotal

n = PV/RT if V, R, T are constant then n = P

*the partial pressure of a particular component of a gaseous mixture is the mole fraction of that component times the total pressure

- Kinetic Theory
-based on observations used to explain and predict the behavior of ideal gases

A. Volume of particles can be assumed to be negligible

B. Particles are in constant motion, pressure is created by the collision of particles with the walls of the container

C. The particles are assumed to exert no forces on each other

D. The average KE of a collection of gas particles is assumed to be directly proportional to the Kelvin temp. of the gas.

- Effusion: passage of a gas through tiny holes in the container
A. Graham’s Law: the rate of effusion of a gas is inversely proportional to the square root of the mass of the particles of gas

B. Diffusion: mixing of gases due to random motion of particles

- Real Gases
A. No gas exactly follows ideal gas behavior (close under high temp, and low pressure)

B. Van der Waal’s equation: describes the behavior of a real gas by correcting for the actual volume of gas particles and forces between gas particles

[Pobs + a(n/v)2] x (V – nb) = nRT

Pobs = observed pressure

a(n/v)2 = pressure correction

V = volume of container

V – nb = volume correction

*a and b are constants obtained from observation of real gases

Read on your own!