Which factors contribute to Average Velocity of Gas Particles?

[axp107]

Roman Numeral Question

I. Molecular Mass
II. Temperature
III. Pressure

I said Temperature and Pressure. Don't have the answer

I reasoned... Temperature gives the particles more kinetic energy... and decreasing pressure.. makes them collide with each other more (faster.. maybe. not sure)

I thought Molecular mass did not play a role b/c aren't we told that all gases behave the same way? That's why we have the ideal gas law?

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[EECStoMed]

I would pick Molecular Mass and Temperature only. NO pressure in involved. TEMP obviously. As for molecular mass, think about the diffusion law between two molecules, the heavier it is, the slower it will move.

Roman Numeral Question

I. Molecular Mass
II. Temperature
III. Pressure

I said Temperature and Pressure. Don't have the answer

I reasoned... Temperature gives the particles more kinetic energy... and decreasing pressure.. makes them collide with each other more (faster.. maybe. not sure)

I thought Molecular mass did not play a role b/c aren't we told that all gases behave the same way? That's why we have the ideal gas law?

 

[axp107]

How would pressure affect a gas particle in general?

Lets say you increase Pressure. Would this Lower its Kinetic Energy? But not velocity? I don't see how that works?

 

[EECStoMed]

So what's your answer?

How would pressure affect a gas particle in general?

Lets say you increase Pressure. Would this Lower its Kinetic Energy? But not velocity? I don't see how that works?

 

[Funky]

I think it cant be pressure unless you are putting work into the system. I do believe there is a formula with molecular mass that deals with KE but I'm not sure.

 

[Tensyle]

I think only Temperature will affect the average velocity of the gas particles in a gas. I dont think it can be molecular mass. Molecular mass only affects KE, not the average velocity. Also, Molecular mass can affect individual velocity of gas particles in a gas, but the average velocity only depends on the temperature, right?

 

[w00tz]

How would pressure affect a gas particle in general?

Lets say you increase Pressure. Would this Lower its Kinetic Energy? But not velocity? I don't see how that works?

Pressure is only considered when it is extremely high. This would cause the intermolecular forces that were negligible at low pressures to become significant. I THINK these attractive forces decrease the velocities of the particles.

 

[apoptos]

I'm hoping someone can really articulate how to distinguish and unify the kinetic molecular theory of gases and graham's law... I've thoroughly confused myself at this point, so I thought I'd share the pain.

according to the kinetic molecular theory of gases, the average kinetic energy of a gas particle is proportional to the absolute temperature of the gas. The speeds of gases are defined in terms of their average molecular speed, which represents the mathematical average of all the gas particles in the sample:

Graham's law shows mathematically that under isothermal and isobaric conditions, the rates at which two gases diffuse are inversely proportional to the square root of their molar masses...

So temperature is a factor, and molar mass is a factor. What happens to PV=nRT? How can pressure have no effect whatsoever?

 

[EECStoMed]

Ok, I think I know how this works.

Temperatures tells you the average KE within a certain gas, so at let's say 373K or 100C all the gas molecules within say container A and container B have the same KE. Notice there's no mention of velocity here.

But if you're going to talk about average velocity, then I believe it's all relative, which is where the molar mass comes into play and you can compare two separate avg velocities of two gases using Graham's law.

And as far as pressure goes, unless the container's volume stays the same and you are doing work on the gas by increasing the pressure within the container, it says nothing about KE or avg velocity.

I'm hoping someone can really articulate how to distinguish and unify the kinetic molecular theory of gases and graham's law... I've thoroughly confused myself at this point, so I thought I'd share the pain.

according to the kinetic molecular theory of gases, the average kinetic energy of a gas particle is proportional to the absolute temperature of the gas. The speeds of gases are defined in terms of their average molecular speed, which represents the mathematical average of all the gas particles in the sample:

Graham's law shows mathematically that under isothermal and isobaric conditions, the rates at which two gases diffuse are inversely proportional to the square root of their molar masses...

So temperature is a factor, and molar mass is a factor. What happens to PV=nRT? How can pressure have no effect whatsoever?

 

[Foghorn]

KE (avg) = (1/2)m [v(avg)]^2. KE is proportional to the temperature.


Rearranging the equation: 2[KE (avg)/m] = [v(avg)]^2

 

[EECStoMed]

KE (avg) = (1/2)m [v(avg)]^2. KE is proportional to the temperature.


Rearranging the equation: 2[KE (avg)/m] = [v(avg)]^2

Temperature, definitely. The question is if mass contributes to avg velocity.

 

[Foghorn]

Temperature, definitely. The question is if mass contributes to avg velocity.

Yes. Mass is inversely proprotional to the square of the average velocity.

 

[ArchieMoses]

Roman Numeral Question

I. Molecular Mass
II. Temperature
III. Pressure

I said Temperature and Pressure. Don't have the answer

I reasoned... Temperature gives the particles more kinetic energy... and decreasing pressure.. makes them collide with each other more (faster.. maybe. not sure)

I thought Molecular mass did not play a role b/c aren't we told that all gases behave the same way? That's why we have the ideal gas law?

I think this is like the effusion of different gases down a tube. The heavier gases move slower (molecular mass), and the avg. kinetic energy (temperature) also influences how fast they move. so i would pick I and II.

the pressure would work too but I don't think you can assume that volume is constant here unless they tell you.


edit: foghorn stop typing so fast

 

[Foghorn]

I think this is like the effusion of different gases down a tube. The heavier gases move slower (molecular mass), and the avg. kinetic energy (temperature) also influences how fast they move. so i would pick I and II.

the pressure would work too but I don't think you can assume that volume is constant here unless they tell you.


edit: foghorn stop typing so fast

The effusion/diffusion argument, I think, is inappropriate in this case since that has to do with gas molecules going through a semi-permeable membrane. That's not what's being asked.

To simplify matters, pressure is dependent on velocity which is dependent on temperature. In other words, temperature is what causes gas molecules of a particular mass to have velocity i.e. KE. This velocity in turn causes the molecules to collide with the container walls. Without those collisions, there is no pressure. Therefore, pressure is a consequence of temperature, not the other way around.

 

[Schaden Freud]

Molecular Mass and Temperature.

The average kinetic energy of a gas is directly proportional to the temperature of that gas, and since KE=1/2mv^2, molecular mass and velocity are inversely related.

Roman Numeral Question

I. Molecular Mass
II. Temperature
III. Pressure

I said Temperature and Pressure. Don't have the answer

I reasoned... Temperature gives the particles more kinetic energy... and decreasing pressure.. makes them collide with each other more (faster.. maybe. not sure)

I thought Molecular mass did not play a role b/c aren't we told that all gases behave the same way? That's why we have the ideal gas law?

 

[Tensyle]

Roman Numeral Question

I. Molecular Mass
II. Temperature
III. Pressure

I said Temperature and Pressure. Don't have the answer

I reasoned... Temperature gives the particles more kinetic energy... and decreasing pressure.. makes them collide with each other more (faster.. maybe. not sure)

I thought Molecular mass did not play a role b/c aren't we told that all gases behave the same way? That's why we have the ideal gas law?

Where did you get this question from? If it is a practice test from somewhere, I can see if I can find an answer....

 

[AWhitehair]

The correct answer is molecular mass and temperature as some have eluded to above.

 

[AWhitehair]

Where did you get this question from? If it is a practice test from somewhere, I can see if I can find an answer....

It was on the Kaplan Diagnostic BTW.

 

[Tensyle]

okk its settled. haha

 

[pandulce321]

First off, the original post reasoned that a decrease in pressure would cause an increase in collisions. Correct me if I am wrong, but that is backwards... increasing the pressure increases the number of collisions with the walls of the container (the amount of force over area).

The part I am having a hard time with is that an increase in pressure, increases temperature (PV=nRT) and an increase in temperature increases the KE of a gas (and in turn its velocity).

 

[Foghorn]

The part I am having a hard time with is that an increase in pressure, increases temperature (PV=nRT) and an increase in temperature increases the KE of a gas (and in turn its velocity).

PV is just another way expressing average KE. Specifically, PV = (2/3)average translational KE. Since the the masses and velocities of individual gas molecules within a container can't be directly measured, other methods must be used. Namely, putting gas molecules within a container of particular volume then measuring pressure. That pressure represents the average collision frequency of all gas molecules. Collisions can only occur if the molecules have velocity i.e. KE. Molecules only have KE (translational) if energy, in the form of temperature, is transferred to them.