Microscopic Gas

1 microscopic

1.1 kinetic theory
1.2 brownian motion
1.3 intermolecular forces

microscopic

if 1 observe gas under powerful microscope, 1 see collection of particles (molecules, atoms, ions, electrons, etc.) without definite shape or volume in more or less random motion. these neutral gas particles change direction when collide particle or sides of container. in ideal gas, these collisions elastic. particle or microscopic view of gas described kinetic-molecular theory. assumptions behind theory can found in postulates section of kinetic theory.

kinetic theory

kinetic theory provides insight macroscopic properties of gases considering molecular composition , motion. starting definitions of momentum , kinetic energy, 1 can use conservation of momentum , geometric relationships of cube relate macroscopic system properties of temperature , pressure microscopic property of kinetic energy per molecule. theory provides averaged values these 2 properties.

the theory explains how gas system responds change. example, gas heated absolute zero, when (in theory) still, internal energy (temperature) increased. gas heated, particles speed , temperature rises. results in greater numbers of collisions container per unit time due higher particle speeds associated elevated temperatures. pressure increases in proportion number of collisions per unit time.

brownian motion

random motion of gas particles results in diffusion.

brownian motion mathematical model used describe random movement of particles suspended in fluid. gas particle animation, using pink , green particles, illustrates how behavior results in spreading out of gases (entropy). these events described particle theory.

since @ limit of (or beyond) current technology observe individual gas particles (atoms or molecules), theoretical calculations give suggestions how move, motion different brownian motion because brownian motion involves smooth drag due frictional force of many gas molecules, punctuated violent collisions of individual (or several) gas molecule(s) particle. particle (generally consisting of millions or billions of atoms) moves in jagged course, yet not jagged expected if individual gas molecule examined.

intermolecular forces

when gases compressed, intermolecular forces shown here start play more active role.

as discussed earlier, momentary attractions (or repulsions) between particles have effect on gas dynamics. in physical chemistry, name given these intermolecular forces van der waals force. these forces play key role in determining physical properties of gas such viscosity , flow rate (see physical characteristics section). ignoring these forces in conditions (see kinetic-molecular theory) allows real gas treated ideal gas. assumption allows use of ideal gas laws simplifies calculations.

proper use of these gas relationships requires kinetic-molecular theory (kmt). when gas particles possess magnetic charge or intermolecular force gradually influence 1 spacing between them reduced (the hydrogen bond model illustrates 1 example). in absence of charge, @ point when spacing between gas particles reduced can no longer avoid collisions between @ normal gas temperatures. case increased collisions among gas particles include fixed volume of gas, upon heating contain fast particles. means these ideal equations provide reasonable results except extremely high pressure (compressible) or high temperature (ionized) conditions. notice of these excepted conditions allow energy transfer take place within gas system. absence of these internal transfers referred ideal conditions in energy exchange occurs @ boundaries of system. real gases experience of these collisions , intermolecular forces. when these collisions statistically negligible (incompressible), results these ideal equations still meaningful. if gas particles compressed close proximity behave more liquid (see fluid dynamics).