chapter 12 the behavior of gases

Boyles Law

at constant temp, has pressure and volume are inversely related P1V2=P2V2

pressire

can be affected by changing the number of moles of a gas--directly proportional because more gas, more collisions volume of a gas--indirectly proportional because increasing size of container means less collisions temperature of a gas-directly proportional because higher temp means higher KE means more collisions

van der waals equation

correct ideal-gas equation known as VDW equation. its takes these deviations of real gas into account

combined gas law

deals with the situation where only the number of molecules stays constant (P1*V1)/T1=(P2*V2)/T2

Real Gas

no real gas behaves like the ideal gas equation predicts intermolecular attractions exist real gas molecules do occupy a finite volume The effect of IMA reduces the actual pressure exerted by the gas particle with the wall of a container. Less forceful collisions due to its attraction to its neighboring thus pressure of real gas is less than pressure of ideal gas

Dalton's Law of Partial Pressures

for a mixture of gases, the tool pressure is the sum of the pressures each gas would exert if it were alone Ptot=P1 + P2 + P3 + .... using the ideal gas law, can change to the sum of the moles (RT/V)

characteristics of gases

in gas state, molecules are separated by large distances. IMF are very small Unlike liquids and solids, they expand to fill their containers, are highly compressible, have extremely low densities

standard pressure

normal atmospheric pressure at sea level 1.00 atm, 760 torr, 760 mm Hg (u can use torr and mm Hg interchangeably) and 101.3 kPa

density

mass divided by volume D= m/V D=MP/RT M=molar mass

Pressure and Temperature of a Gas

raising the temp of a gas increases the pressure, if the volume is held constant. the molecules hit the walls hard and more frequently. temp and pressure are directly prop heating aerosol can cause it to explode or rupture because of the increased pressure

real gases behave like ideal gases when...

when the molecules are far apart when they do not take up as big a percentage of the space we can ignore their volume This is at LOW PRESSURE when the molecules are moving fast (high temperature). collisions are harder and faster. molecules are not next to each other very long. attractive forces can't play a role

Kinetic Theory Revisited

1. Gases consist of hard, spherical particles (usually molecules or atoms) 2. small-so the individual volume is considered to be insignificant 3. large empty space between them 4. easily compressed and expanded 5. no attractive or repulsive forces 6. collisions between gas particles are perfectly elastic collisions 7. move rapidly in constant motion the AVG Kinetic Energy of a collection of gas particles is directly proportional to the Kelvin temperature of the gas

Variable that describe a gas

4 variables and their common units 1. pressure P in kilopascals or atmospheres 2. volume (v) in liters 3. temperature in kelvin 4. number of moles (n)

pressure

amount of force applied to an area P=F/A atmospheric pressure is the weight of air per unit of area

Gas Laws

describe HOW gases behave. temperature must always be in KELVIN when working w Gas Law problems

Ideal gas law

describes the relationship between the amount of the gas and its temperature pressure and volume . STP= temperature is 0 degrees Celsius, 1atm, 22.4 L, and 1 mole PV=nRT R=Ideal Gas Law Constant R=o.o8206 L * atm/ K* moles To use ideal gas laws we must assume the gases behave like ideal gases which 1. have no volume 2. no attractive forces ideal gas does not really exist because gases have both volume and attractive forces but it gives us a very close approximation

avogadro's hypothesis

equal volumes of gases at the same temperature and pressure contain equal numbers of particles saying that two rooms of the same size could be filled with the same number of objects, whether they were marbles or baseballs 22.4 L of a gas= 1 mole of a gas at STP

molecular diffusion

molecules moving from areas of high concentration to low concentration high pressure to low pressure askk about the question If a container of perfume is opened at one end of the room and you're at the other end, why can't u smell the perfume? (remember the speeds of gas molecules is large at r.t.!) although the gas molecules travel in straight lines, there are many collisions with other gas molecules that will cause trajectory changes

graham's law

rate of effusion and diffusions depends on the speed of the gas molecules heavier molecules move slower at the same temperature o a lighter molecule so they effuse and diffuse slower rate of effusion of a gas is INVERSELY proportional to the square root of its molar mass can be used to compare the relative rate of effusion between two gases gas 1 should be the lighter of the two gases in order to get a whole number r1/r2= square root of M2/M1 tells us how much faster one is moving relative to the other

Mole fraction

ratio of number of moles of a certain component of a mixture to number of moles total in mixture mole fraciton: x1= n1/ntotal

effusion

the rate of escape of a gas thru a pinhole into an evacuation space

Gay-Lussac's Law

the temp and pressure of a gas are directly relation at constant volume P1/T1=P2/T2

Charles Law

the volume of a gas is directly proportional to the kelvin temperature, if the pressure is held constant V1/T1=V2/T2