Diesel

n7axw wrote:I am not a techy, so I tend to get in trouble when discussions get too technical...

But...could someone explain why diesel explodes under pressure? I kow there is heat involved, but where does it come from? How hot does it have to get in a cylinder with pressurized diesel before it explodes? I've looked it up, but so far none of the discussions have explained this...

Don

Diesel doesn't explode under pressure. In an Otto cycle or diesel engine, Air is drawn into a cylinder, The air intake valve closes and the piston rises compressing the air. When a gas such as air is compressed it heats up. When the piston reaches the top of its travel, diesel is injected into the compressed air. The air is hotter than the ignition temperature of the diesel at his point so the diesel burns spontaneously which is why no spark plugs are required in diesel engines.

n7axw wrote:But...could someone explain why diesel explodes under pressure? I kow there is heat involved, but where does it come from?

The heat comes from compression. When you compress something, it gets hotter -- someone else can get all sciency and technical to explain why, but Compressed=Hotter is a general principle that should be remembered.

I don't recall the required temperature, but it would vary based on the characteristics of the fuel anyway. IIRC, Diesels require a compression ratio of 10:1 or higher and Gasoline engines can't go above 8:1 without very high octane fuel and an anti-knock compound like tetraethyl lead.

Weird Harold wrote:

n7axw wrote:But...could someone explain why diesel explodes under pressure? I kow there is heat involved, but where does it come from?

The heat comes from compression. When you compress something, it gets hotter -- someone else can get all sciency and technical to explain why, but Compressed=Hotter is a general principle that should be remembered.

I don't recall the required temperature, but it would vary based on the characteristics of the fuel anyway. IIRC, Diesels require a compression ratio of 10:1 or higher and Gasoline engines can't go above 8:1 without very high octane fuel and an anti-knock compound like tetraethyl lead.

Here's the "sciency" description:

Gay-Lussac's law states that P1/T1 = P2/T2 (http://en.wikipedia.org/wiki/Gas_laws#Gay-Lussac.27s_law)

Where:
P = pressure
T = Temperature


The final temperature (T2) can be written as:
T2 = P2/P1 * T1

And so, as the pressure (P2) increases (from the decreasing volume - to approximately 4 MPa), the temperature (T2) increases (to approximately 550 degrees celcius).

n7axw wrote: But...could someone explain why diesel explodes under pressure? I kow there is heat involved, but where does it come from? How hot does it have to get in a cylinder with pressurized diesel before it explodes? I've looked it up, but so far none of the discussions have explained this...n7ax

MWadwell wrote an excellent sciency explanation. Here's my intuitive explanation.

Temperature is directly proportional to the average velocity of molecules. If you double the average velocity then you double the temperature (in an absolute sort of way). The molecules themselves are moving at different velocities, colliding off each other, speeding up and slowing down. The molecules also attract each other. That means that as they move toward each other they speed up, and they slow down as they move away from each other. (At very short distances they also repel each other, but I don't want to complicate this too much. I merely want to complicate it the right amount.)

Now, when you decrease the volume of a diesel/air mixture then the molecules don't have a chance to move apart from each other, before they collide and start moving in a different direction. Since they don't move as far then they don't slow down as much. Since they don't slow down as much their average velocity is higher. And since the average velocity is higher then the temperature is higher.

The temperature is a measure of the heat energy in the mixture. The work to increase the heat energy comes from the force needed to compress the diesel/air mixture, and from the distance the piston has to move in compressing it.

With higher average velocities that means that it's more likely that the oxygen molecules and the diesel molecules will collide in exactly the right way to overcome the repulsive force I mentioned parenthetically and to react to form water and carbon dioxide. It explodes. The auto-ignition temperature for diesel/air is 256 °C.

~Tonto

Tonto Silerheels wrote:n7axw wrote: But...could someone explain why diesel explodes under pressure? I kow there is heat involved, but where does it come from? How hot does it have to get in a cylinder with pressurized diesel before it explodes? I've looked it up, but so far none of the discussions have explained this...n7ax

MWadwell wrote an excellent sciency explanation. Here's my intuitive explanation.

Temperature is directly proportional to the average velocity of molecules. If you double the average velocity then you double the temperature (in an absolute sort of way). The molecules themselves are moving at different velocities, colliding off each other, speeding up and slowing down. The molecules also attract each other. That means that as they move toward each other they speed up, and they slow down as they move away from each other. (At very short distances they also repel each other, but I don't want to complicate this too much. I merely want to complicate it the right amount.)

Now, when you decrease the volume of a diesel/air mixture then the molecules don't have a chance to move apart from each other, before they collide and start moving in a different direction. Since they don't move as far then they don't slow down as much. Since they don't slow down as much their average velocity is higher. And since the average velocity is higher then the temperature is higher.

The temperature is a measure of the heat energy in the mixture. The work to increase the heat energy comes from the force needed to compress the diesel/air mixture, and from the distance the piston has to move in compressing it.

With higher average velocities that means that it's more likely that the oxygen molecules and the diesel molecules will collide in exactly the right way to overcome the repulsive force I mentioned parenthetically and to react to form water and carbon dioxide. It explodes. The auto-ignition temperature for diesel/air is 256 °C.

~Tonto

I'm afraid this is completely erroneous. The attraction between molecules does not significantly affect the velocity of the molecules in a gas.

SWM wrote: I'm afraid this is completely erroneous. The attraction between molecules does not significantly affect the velocity of the molecules in a gas.

I suspect you subscribe to the kinetic molecular theory. The kinetic molecular theory is very useful in that it can explain a wide variety of phenomena. Unfortunately, there are some few phenomena it cannot explain such as expansion into a vacuum. Suppose you have one mole of neon in a one-liter chamber at standard temperature, and suppose that this container is piped through a valve to an identical chamber containing vacuum. Under the kinetic molecular theory the temperature of the gas would remain unchanged after opening the valve since there is nothing to change the average speed of the molecules. (Assume the two chambers are at an equal height, and that there is no electrical or magnetic fields imposed on them.) If you were to perform the experiment, though, you would find that the temperature falls just like the ideal gas law predicts.

I will confess, though, to two errors. The first is that I discussed the average velocity of the molecules. The average velocity of the molecules is, of course, zero. I should have said the average speed of the molecules. The second is that I said that the temperature is directly proportional to the average velocity of the molecules. In fact, the temperature is proportional to the kinetic energy of the molecules, which, in turn, is proportional to the square of the average speed. I should have said that if you double the average speed then you quadruple the temperature.

There is no reason to dump the kinetic molecular theory for those times when it's applicable. I can't recommend confusing it with reality, though.

~Tonto

Tonto Silerheels is correct.

The velocity distribution for molecules is independent of their positions or current potential energy. The interactions between molecules have no effect on their average kinetic energy, only on their trajectories if you wait a while and watch. Thus, the velocity distribution (likelihood of finding a particular v for a randomly chosen molecule) for water vapor in air and for water in liquid water at the same temperature are the same.

See my textbook "Elementary Lectures in Statistical Mechanics" (Springer-Verlag) for the details.

Tonto Silerheels wrote:SWM wrote: I'm afraid this is completely erroneous. The attraction between molecules does not significantly affect the velocity of the molecules in a gas.

I suspect you subscribe to the kinetic molecular theory. The kinetic molecular theory is very useful in that it can explain a wide variety of phenomena. Unfortunately, there are some few phenomena it cannot explain such as expansion into a vacuum. Suppose you have one mole of neon in a one-liter chamber at standard temperature, and suppose that this container is piped through a valve to an identical chamber containing vacuum. Under the kinetic molecular theory the temperature of the gas would remain unchanged after opening the valve since there is nothing to change the average speed of the molecules. (Assume the two chambers are at an equal height, and that there is no electrical or magnetic fields imposed on them.) If you were to perform the experiment, though, you would find that the temperature falls just like the ideal gas law predicts.

I will confess, though, to two errors. The first is that I discussed the average velocity of the molecules. The average velocity of the molecules is, of course, zero. I should have said the average speed of the molecules. The second is that I said that the temperature is directly proportional to the average velocity of the molecules. In fact, the temperature is proportional to the kinetic energy of the molecules, which, in turn, is proportional to the square of the average speed. I should have said that if you double the average speed then you quadruple the temperature.

There is no reason to dump the kinetic molecular theory for those times when it's applicable. I can't recommend confusing it with reality, though.

~Tonto

MWadwell wrote:[
Here's the "sciency" description:

Gay-Lussac's law states that P1/T1 = P2/T2 (http://en.wikipedia.org/wiki/Gas_laws#Gay-Lussac.27s_law)

Where:
P = pressure
T = Temperature


The final temperature (T2) can be written as:
T2 = P2/P1 * T1

And so, as the pressure (P2) increases (from the decreasing volume - to approximately 4 MPa), the temperature (T2) increases (to approximately 550 degrees celcius).

You need to go further than Gay-Lussac's law, and on to the "Combined Gas Law" which gives that (For an ideal gas) and for a given mass of gas

p1v1/T1 = p2v2/T2 where T is the Absolute temperature of the gas.

I understand the Ideal Gas Law and was not trying to argue in favor of the Kinetic Molecular Theory. What I was objecting to were the statements:

The molecules also attract each other. That means that as they move toward each other they speed up, and they slow down as they move away from each other.

and

Now, when you decrease the volume of a diesel/air mixture then the molecules don't have a chance to move apart from each other, before they collide and start moving in a different direction. Since they don't move as far then they don't slow down as much.

I don't believe that correctly describes the situation.

SWM wrote:I understand the Ideal Gas Law and was not trying to argue in favor of the Kinetic Molecular Theory. What I was objecting to were the statements:

The molecules also attract each other. That means that as they move toward each other they speed up, and they slow down as they move away from each other.

and

Now, when you decrease the volume of a diesel/air mixture then the molecules don't have a chance to move apart from each other, before they collide and start moving in a different direction. Since they don't move as far then they don't slow down as much.

I don't believe that correctly describes the situation.

You are correct. To compress a gas, you need a moving wall, and collisions with the moving wall change atomic velocities. There is also an issue with whether the gas can exchange heat with the wall during the compression, whcih could be adiabatic or isothermal. or etc etc etc. "isobaric" is especially amusing, as students have trouble seeing how that is possible. How do I do an isobaric (constant pressure) expansion of a gas ina cylinder with piston?

Hint: Use a blowtorch.

Real gasses have a Joule Thompson coefficient, which describes the change of temperature on expansion. For hydrogen, the coefficient is negative...the gas gets hotter as it expands. Readers may correctly infer a safety issue, complexified by the technical issue that hydrogen flames are close to invisible. (The Hindenberg flame is the hull metal burning.)