Gases

Properties of ideal gases Gases are evenly distributed in a volume. They have very weak IMF, intermolecular forces. (van der Waal) They have frequent elastic collisions They have high kinetic energy due to their high velocities. Most of the volume they occupy is empty space; they are easily compressed.

Kinetic molecular theory of gases Kinetic energy (KE) is the energy a particle has as a result of its molecular movement. KE=1/2(molar mass)velocity2 Temperature is directly proportional to KE. For example: an increase in molecular movement causes an increase in temperature. This theory is best applied to gases since they are free moving particles and are subject to great changes in velocity. http://www.youtube.com/watch?v=EzWRvK0zhQk KE and Temp.

Applying the kinetic molecular theory to gases An example: Two gases, He and H2 are at the same temperature of 20C, have the same mole value, but are in separate but identical containers. Both gases have the same KE due to their same temperature. Both gases will exert the same pressure since there is the same number of molecules of each gas. (same mole value) and are in the same volume. The hydrogen will move faster than the He since it is a lighter gas H2= 2g/mol, He=4g/mol

Applying the kinetic molecular theory to gases continued Summary: Remember KE= 1/2mv2. So if gases are at the same temp. then they have the same KE. If gases have different masses, then their velocities must be different. The rates of diffusion will differ. Regardless of the mass of the gas, gases at the same temperature have the same force of collision/molecule. Heavy gases hit just as hard as light gases if they are at the same temp. Heavier gases have slower velocities, lighter gases have higher velocities. This equalizes the force of collision.

Kinetic energy videos. http://www.youtube.com/watch?v=UNn_trajMFo KE and size of molecules http://www.youtube.com/watch?v=EzWRvK0zhQk KE and Temp. (too long)

Applying the kinetic molecule theory to solids and liquids. This theory does not apply as easily to solids and liquids due to their restricted motion. It is safer to apply the KE theory to solids and liquids if they are at the same temp. H2O and alcohol at 20C have similar KE Or apply it to the same solid or liquid. H2O at 20C has less KE than H2O at 100C

Kelvin temperature scale/ absolute temperature scale. Properties of this temperature scale Based upon the molecular movement of gases At zero Kelvin there is no molecular movement. Since temp. is based upon molecular movement, O Kelvin is the coldest temp. possible anywhere. O Kelvin = -273°C Add 273 to C to get Kelvin

Non Ideal Gas- SF6 http://www.youtube.com/watch?v=xQo-v_F1P9U Mythbusters inhale SF6

Properties of gases

Pressure Pressure= force/area For gases: the force is created by the strength and the number of collisions. The strength of the collisions is determined by the kinetic energy of the molecules. Higher KE = higher velocity= greater force of collisions. The number of collisions is determined by: the mole value/volume of the gas and the velocity of the gas. For gases: the area is defined by the pressure unit. (walls of the container).

Units of pressure Millimeters of mercury (mmHg) = based upon the height of a mercury column in a barometer. Pressure at sea level=760mmHg Torr (named after Torricelli) is the same unit as the mmHg. You will see this unit in the texts. 760 torr= pressure at sea level. Atmospheres (atm)= 1atm is equal to the pressure at sea level.

Pascal (Pa)= SI unit, =1newton/meter2, 101,325 Pa = pressure at sea level. kPa is most often used. 101.3kPa Pounds per square inch (psi) =14.7 psi is the pressure at sea level.

http://www.youtube.com/watch?v=GgBE8_SyQCU A Torrecilli Barometer

Summary of pressure at sea level 1 atm 101,325 Pa or 101.3kPa 760mmHg 760 torr 14.7psi (will not be using this unit.)

Standard temperature and pressure: STP For research and convenience a standard temp and pressure are used when working with gases. Standard temperature: 273K or 0 celcius Standard pressure: sea level= 1atm, 760mmHg, 760torr, 101.3 kPa, 14.7psi MEMORIZE STP!!!!!!!!!!!

Avogadro’s Hypothesis Equal volumes of gases at the same temperature and pressure have the same number of molecules. At S.T.P. the volume of one mole of any ideal gas is 22.4L. At STP the volume of a gas can be easily determined using a proportion as long as the amount (moles) of gas is known.

Pressure and volume relationship An increase in pressure will decrease volume, or decreasing the volume will increase the pressure. When the volume is decreased, the number of collisions increase per time. The force of collisions stays the same. Pressure and volume are inversely related. The equation: P=k/V or P1V1=P2V2 when moving the gas into another container. Amount and temperature of gas remains constant.

Pressure and volume

Pressure and volume: Boyles law

Effects of temperature on pressure and volume Temperature and pressure in a fixed volume Increasing the temperature increases the pressure. Molecules have more KE which means their velocity increases. (mass stays constant) Higher velocity= more collisions and greater force in the collisions so the pressure increases

Effects of temperature on pressure and volume Temperature and volume (container can expand) Increasing temperature will increase the volume, the pressure remains constant. Temperature is directly proportional to volume. T=kV (charles law) Molecules have more KE which means their velocity increases. Higher velocity=more collisions and greater force in the collisions. This expands the container. The container will expand until the pressure inside=the pressure outside.

Gas Laws so far that you will be using on your homework. Boyle’s Law: P1V2 = P2V2 Charle’s Law: V1/T1 =V2/T2 Gay-Lussac law: P1/T1 =P2/T2 Combined Gas Law: uses the above laws: P1V1/T1 =P2V2/T2 Important note: you must use the Kelvin temperature scale for all “T’s”

IDEAL GAS LAW: uses only ideal gases and is the only law that allows for a change in moles. PV=nRT P=pressure V=volume(L) n=moles of gas R=constant T=temperature(K)

R Constants