Why do gas exert pressure

Hello everyone, I have made number of topics recently I think they have all stem from my poor understanding in this matter. Ok in a gas, molecules can take up a volume and exert a pressure. Inside a fixed container when you increase the temperature pressure increases because more gas molecules are hitting the walls. Now in a liquid the molecules are close together and volume is determined by the shape of the container. So inside the container liquid molecules don't hit the walls and exert pressure some other way. What is this way? Is this called hydrostatic pressure? Also when you make the molecules move faster in a liquid would that increase the pressure like in a gas? According to what I have read it actually decrease hydrostatic pressure. However I have a picture in mind more kinetic more molecules hitting the walls higher the pressure. Maybe I'm just not understanding how gases and liquids behave in a molecular level when it comes to pressure. So please help me with this. I would greatly appreciate. Thanks In a gas the molecules move about freely and hit all the walls equally - it is this motion that causes pressure which is the same everywhere in the container. The pressure only depends on the number and speed of the impacts - so only on the amount of gas and it's temperature. In a liquid the molecules near the walls push against the wall and are pushed by the other molecules above/behind them. The deeper you go the more molecules are above you and the more force they...

Answer:- The random motion of particles in a gas causes the pressure it exerts. Gases have weak intermolecular interactions, and the particles are in a constant state of random motion, colliding with the container’s walls. These encounters with the container’s walls put pressure on the gas. How does a gas exert pressure. How does a gas exert pressure quizlet? Gas exerts pressure on the walls of its container because the particles collide with the wall. What happens to the temperature in gases? If the gas gets hotter, the pressure will be bigger. If the gas gets cooler, the pressure will be smaller. Table of Contents • • • • • • Does gas have exert pressure? Gases exert pressure, which is force per unit area. The pressure of a gas may be expressed in the SI unit of pascal or kilopascal, as well as in many other units including torr, atmosphere, and bar. Why does a gas exert pressure ?`? (a) Why does a gas exert pressure? Answer: (a) When gas is stored in a container, the fast-moving particles of the gas collide with each other and with the walls of the container; thus, exerting pressure on the walls of the container. How does the gas exert pressure on its container? The moving particles in a gas collide with each other and also with the walls of the container and due to these collisions, Gases exert pressure on the walls of the container. How does a gas exert pressure? Answer: (a) When gas is stored in a container, the fast-moving particles of the gas collide with each othe...

In the The molecules exert no intermolecular forces on each other and they have no potential energy associated with them in their motion. Thus the energy is wholly kinetic. The concept of pressure is explained in kinetic theory as a consequence of kinetic energy of gases. Let me explain it. Consider an ideal gas (the ones with which kinetic theory is concerned) in a closed container. Due to the troublesome motion of the gas molecules, they will collide with each other, some of them shall collide with the container walls and then bounce back and this process continues. Now if #mvecv# be the momentum of a single molecules before it bumps into the wall and it makes an angle #theta# with the wall and since the wall is assumed to be perfectly smooth it rebounds at an angle #theta# again (directed opposite to the direction from which the gas molecule initially came. Thus, it transfers a momentum, #mvCos theta# and undergoes a momentum change and rebounds with a final momentum #-mvCos theta#. Thus the momentum change turns out to be, #-mvCos theta + (-mvCos theta) = -2mvCos theta# This implies that there is a momentum transfer of #2mvCos theta# to the wall from the molecule. Thus considering a number of molecules striking the walls, they transfer momentum. Now force is defined to be, the time rate of change of momentum. Thus the molecules exert a force on the walls. Further since, pressure is defined to be force per unit area, then the molecules exert a pressure on the container ...

According to kinetic theory, molecules inside a volume (e.g. a balloon) are constantly moving around freely. During this molecular motion, they constantly collide with each other and with the walls of the container. In a small balloon, that would be many thousands of billions of collisions each second. The force of impact of a single collision is too small to measure. However, taken all together, this large number of impacts exerts a considerable force on the surface of the container. If they hit the surface of the balloon straight on (at a 90 ° angle), they exert their maximum force. If they hit the surface at an angle less than 90 °, they exert a smaller force. The sum of all these forces causes the pressure, #p#, that is exerted by the gas. The diagram above represents a balloon containing molecules of a gas (the red dots). The yellow arrows indicate that the gas pressure, #"p"#, in the balloon is exerted outward against the walls of the balloon. The larger the number of collisions per area of the container, the larger the pressure: Pressure = #"Force"/"Area"# or #p = F/A#. The direction of this force is always perpendicular to the surface of the container at every point. The video below gives a good explanation of gas pressure.

\( \newcommand\) • • • • • • Another important property of liquids (and solids) that is governed by intermolecular forces is vapor pressure. Vapor pressure is defined as the partial pressure of a substance in the gas phase (vapor) that exists above a sample of the liquidin a closed container. The phenomenon of vapor pressure is explained by the kinetic molecular theory again, which shows that a liquid always exists in equilibrium with its vapor. Evaporation: Liquid/Vapor Equilibrium The average energy of the particles in a liquid is governed by the temperature. The higher the temperature, the higher the average energy. But within that average, some particles have energies higher than the average, and others have energies lower than the average. Some of the more energetic particles on the surface of the liquid can be moving fast enough to escape from the attractive forces holding the liquid together. They evaporate. The diagram shows a small region of a liquid near its surface. Notice that evaporation only takes place on the surface of the liquid. That's quite different from If you look at water which is just evaporating in the sun, you don't see any bubbles. Water molecules are simply breaking away from the surface layer. Eventually, the water will all evaporate in this way. The energy which is lost as the particles evaporate is replaced from the surroundings. As the molecules in the water jostle with each other, new molecules will gain enough energy to escape from the sur...

Gases are complicated. They're full of billions and billions of energetic gas molecules that can collide and possibly interact with each other. Since it's hard to exactly describe a real gas, people created the concept of an Ideal gas as an approximation that helps us model and predict the behavior of real gases. The term ideal gas refers to a hypothetical gas composed of molecules which follow a few rules: If this sounds too ideal to be true, you're right. There are no gases that are exactly ideal, but there are plenty of gases that are close enough that the concept of an ideal gas is an extremely useful approximation for many situations. In fact, for temperatures near room temperature and pressures near atmospheric pressure, many of the gases we care about are very nearly ideal. If the pressure of the gas is too large (e.g. hundreds of times larger than atmospheric pressure), or the temperature is too low (e.g. − 200 C -200 \text − 2 0 0 C minus, 200, start text, space, C, end text ) there can be significant deviations from the ideal gas law. For more on non-ideal gases read Perhaps the most confusing thing about using the ideal gas law is making sure we use the right units when plugging in numbers. If you use the gas constant R = 8.31 J K ⋅ m o l R=8.31 \dfrac K kelvin K start text, k, e, l, v, i, n, space, end text, K . If you use the gas constant R = 0.082 L ⋅ a t m K ⋅ m o l R=0.082 \dfrac K kelvin K start text, k, e, l, v, i, n, space, end text, K . Units to use fo...

I always hear pressure defined as the force exerted by particles on the walls of the container they're being held in. This makes sense since the mathematical definition of pressure is $ p = \frac $. So, can pressure exist without walls to exert force on? My understanding of pressure motivates me to think that even without a container, particles of gas in a vacuum could create pressure since their collisions with each other result in forces being exerted on areas (the areas being the surfaces of the particles being collided with). On the other hand, a liquid in a vacuum could not exert pressure since on a microscopic level, the particles of liquid can't collide due to the cumulative strength of the bonds holding them together. $\begingroup$ "an example would be a self gravitating ball of gas, such as a star." Another example of how an (enormous) pressure can exist even without containing walls is a 'dynamic' situation, such as the laser-driven implosion of a NIF capsule. In this case very large pressures are supported by the inertia of the mass of the capsule itself. Of course such a situation can only exist for a short time before the capsule flies apart. $\endgroup$ An addition to Daniel's answer: It regularly happens that a certain definition isn't applicable when talking about liquids and gases. For example, almost all definitions in thermodynamics are defined for closed systems, which is rarely applicable in the real world. To overcome this, there is the notion of a Re...

where P is pressure, V is volume, n is the number of moles (related to mass), R is a constant, and T is temperature. The volume is not infinite because the Earth's gravity has enough "pull" on the molecules to hold them close to the planet. Some gases escape, like helium, but heavier gases like nitrogen, oxygen, water vapor, and carbon dioxide are bound more tightly. Yes, some of these larger molecules still bleed off into space, but terrestrial processes both absorb gases (like Because there is a measurable temperature, the molecules of the atmosphere have energy. They vibrate and move around, bumping into other gas molecules. These collisions are mostly elastic, meaning the molecules bounce away more than they stick together. The "bounce" is a force. When it is applied over an area, like your skin or the Earth's surface, it becomes pressure. How Much Is Atmospheric Pressure? Pressure depends on altitude, temperature, and weather (largely the amount of water vapor), so it's not a constant. However, the average pressure of air under ordinary conditions at sea level is 14.7 lbs per square inch, 29.92 inches of mercury, or 1.01 × 10 5 pascals. Atmospheric pressure is only about half as much at 5 km altitude (about 3.1 miles). Why is pressure so much higher close to the Earth's surface? It's because it's really a measure of the weight of all the air pressing down at that point. If you are high in the atmosphere, there isn't much air above you to press down. At the Earth's sur...

I have studied on the internet that gases exert equal pressure in all directions in a container but liquids do not. In liquids pressure exerted on the wall of a container increases with depth. Why is that so? any logical or intuitive if not conceptual answer is also appreciated. Also can there be a situation in which liquid can exert equal pressure on the walls of their container, independent of depth? Note: I found a pretty similar question here but is a bit complicated and so I couldn't understand it. Also please use liquid/gas terms instead of fluid as it mixes up things for me. $\begingroup$ You say that gases exert equal pressure in all directions in a container. Air in the atmosphere is a gas. If you put a large container (impossibly large I admit) around it, it would not exert equal pressure on all the walls of the container. The reason liquids exert higher pressures at larger depths is because of gravity. On any reasonable scale for a gas, I assume that gravity is ignored when considering pressure differences. (To add to this, fluid pressure is $\rho g h$, where $\rho$ is the density of the fluid, and gases are very dilute fluids) $\endgroup$ I have studied on the internet that gases exert equal pressure in all directions in a container but liquids do not. [...] Why is that so? Actually they both work in the same manner. The cause is the presence of gravity. Pressure increases with depth in a liquid, because the heavy (dense) liquid has to carry the whole column of...